how the noble gases turned out not to be inert

The Mechanism

On the evening of *23 March 1962*, in a small inorganic-chemistry research laboratory in the chemistry building at the *University of British Columbia* in *Vancouver*, *Canada*, *Neil Bartlett* — a 29-year-old British-born chemistry lecturer — synthesised the first known *chemical compound of a noble-gas element*: an orange-yellow solid he initially formulated as *Xe⁺[PtF₆]⁻*, *xenon hexafluoroplatinate*. The result, published as a *242-word short communication* in the *Proceedings of the Chemical Society* in June 1962 (Bartlett, N., "Xenon hexafluoroplatinate(V) Xe⁺[PtF₆]⁻," *Proceedings of the Chemical Society*, p. 218, June 1962), demolished in a single afternoon the textbook position that had been held without serious challenge since *William Ramsay* and *Lord Rayleigh* isolated the first noble gas, *argon*, in 1894 and *Ramsay* had been awarded the *1904 Nobel Prize in Chemistry* for the discovery of the entire *Group 0* (now Group 18) of the periodic table — the so-called *inert gases* — between 1894 (argon) and 1898 (neon, krypton, xenon). The textbook position was that the noble gases, on account of their *closed-shell electronic configurations* (a full outer shell of 2 or 8 electrons), are chemically *unreactive*: they form no compounds, accept no electrons, donate no electrons, and exist always in the monatomic gas phase. *Linus Pauling*, the most prominent inorganic chemist of the 20th century, had stated in the first edition of *The Nature of the Chemical Bond* in 1939 that "argon, krypton, and xenon do not form chemical compounds" — and in 1933 had specifically predicted, on theoretical grounds, that *xenon hexafluoride (XeF₆)* might be preparable but had been talked out of pursuing the synthesis by his colleagues' insistence that the noble-gas inertness was a settled matter. The intellectual path to Bartlett's discovery began with an unrelated piece of work in his own laboratory in 1961. Bartlett had been studying the properties of *platinum hexafluoride (PtF₆)* — a deep red, highly volatile gas with the chemical formula of an octahedral molecular fluoride and the practical character of one of the most powerful oxidising agents then known. While preparing PtF₆ in a glass-and-nickel vacuum line, Bartlett had been routinely working with it in air, where small leaks were inevitable. In a paper published in *Proceedings of the Chemical Society* in *June 1961*, Bartlett and his graduate student *Derek H. Lohmann* reported that PtF₆ readily oxidised *molecular oxygen (O₂)*, producing the salt *dioxygenyl hexafluoroplatinate*, *O₂⁺[PtF₆]⁻*, as an orange crystalline solid. This was, in itself, a remarkable result: O₂ has an *ionisation energy* of *1175 kJ/mol* — the energy required to remove a single electron — and was generally considered too tightly-bound for any chemical oxidiser to ionise. PtF₆ had done so easily. Sitting in his office in *late 1961*, in front of a borrowed copy of a recent periodic-table data compilation, Bartlett happened to notice that the *ionisation energy of xenon* — at *1170 kJ/mol* — was *almost exactly equal* to that of *molecular oxygen*. The numerical coincidence was striking: if PtF₆ could ionise O₂, then a straightforward thermodynamic argument suggested it should also be able to ionise *Xe*. Bartlett went down the hall to the storeroom, signed out a small lecture-bottle of xenon, and set up an apparatus consisting of two glass containers — one holding the deep red PtF₆ gas, the other holding the colourless xenon gas — separated by a thin glass *break-seal*. On the evening of *23 March 1962*, alone in the laboratory, he tapped the apparatus to break the seal. The two gases mixed. *An immediate reaction occurred*, depositing an *orange-yellow crystalline solid* on the inside walls of the glass. Bartlett sealed the apparatus, examined the solid under a microscope, and verified by elemental analysis that the compound contained xenon, platinum, and fluorine in a 1:1:6 ratio — consistent with the formula *Xe⁺[PtF₆]⁻*, a salt in which xenon had been ionised by PtF₆ in exactly the same way O₂ had been the year before. He wrote up the result as a one-page short note in *Proceedings of the Chemical Society* and submitted it the following week; the paper was published, almost without revision, in the *June 1962 issue*. The reception was *swift and dramatic*. Within four months, three independent groups — *Howard Claassen* at *Argonne National Laboratory* (working with *Henry Selig* and *John Malm*), *Cedric Hoppe* at the *University of Münster*, and *Bartlett's own group* — had synthesised *xenon difluoride (XeF₂)*, *xenon tetrafluoride (XeF₄)*, and *xenon hexafluoride (XeF₆)* — the direct combination compounds of xenon and fluorine — by passing electric discharges through Xe-F₂ mixtures or by heating the elements together in nickel vessels. By the end of 1963 the chemistry of the noble gases was an established and rapidly-expanding field; an entire issue of *Science* (*16 August 1963*) was devoted to it. The class of noble-gas compounds now extends to compounds of *krypton* (*KrF₂*, the only known thermodynamically stable krypton compound, prepared at low temperature) and *radon* (*RnF₂*); compounds of *helium* and *neon* and *argon* exist only under exotic conditions (the helium "Van der Waals" complex *HeNe*, the argon compound *HArF* in solid argon matrices at 17 K, etc.). Of practical importance: *xenon difluoride* is now a standard reagent in semiconductor manufacturing for *selective etching of silicon* in MEMS device fabrication, with a global market of approximately *$80 million per year*. *Xenon tetrafluoride* and *xenon hexafluoride* are used as *fluorine-transfer reagents* in organic synthesis. *Xenon* itself, as an inhalation anaesthetic, is used in clinical practice for *neuroprotection during cardiac surgery* (its anaesthetic mechanism is distinct from that of every other clinical anaesthetic, involving inhibition of *NMDA receptors* at the post-synaptic membrane — a discovery itself made possible by the experimental access to xenon chemistry that Bartlett's 1962 result opened). The intellectual significance of the 1962 result extends beyond practical chemistry. The argument that the noble gases must be chemically unreactive *because* of their closed-shell electronic structure was an argument from a single principle to an empirical conclusion — and it was wrong. The lesson, drawn out at length in *George Olah's* 1994 Nobel Lecture, is that *every theoretical argument in chemistry that proves an empirical impossibility on the basis of a single elegant principle should be tested experimentally by someone who hasn't read the argument*. The noble gases are not inert. They were merely *under-studied*, for sixty-seven years, on the basis of an argument that no one had ever seriously checked, until a 29-year-old chemist in Vancouver in March of 1962 noticed that two numbers in a periodic-table data compilation were almost equal and decided to put two gases in two glass containers next to one another and break the seal.