Xenon is one of the six elements that make up Group 18 of the periodic table, known collectively as the noble gases. These elements share a characteristic electron configuration, possessing a completely filled outermost valence shell. This unique atomic structure historically led scientists to classify them as “inert gases,” reflecting the belief that they were incapable of forming stable chemical bonds. For decades, this perceived non-reactivity was a fundamental principle in chemistry.
The Myth of Inertness
The belief in the chemical inertness of the noble gases stemmed directly from the stability afforded by their full valence shells. According to the foundational chemical concept known as the octet rule, atoms strive to achieve an outer shell of eight electrons, a configuration that mimics the noble gases. Since these elements already possessed this stable arrangement, it was theorized they had no energetic motivation to gain, lose, or share electrons with other atoms.
This chemical dogma guided research for the first half of the 20th century. The idea suggested that any attempt to force a reaction would be energetically unfavorable, yielding only fleeting or unstable products. Although some theoretical predictions hinted at the possibility of noble gas compounds under extreme conditions, the consensus remained that stable, isolable compounds were impossible to achieve in a laboratory setting.
The 1962 Synthesis
The long-held myth was definitively shattered in 1962 by the chemist Neil Bartlett at the University of British Columbia. His experiment was based on his prior work with the powerful oxidizing agent, platinum hexafluoride (\(\text{PtF}_6\)). Bartlett had successfully used \(\text{PtF}_6\) to oxidize molecular oxygen (\(\text{O}_2\)), forming the red solid dioxygenyl hexafluoroplatinate (\(\text{O}_2\text{PtF}_6\)). The key insight came from comparing the ionization energy of molecular oxygen (approximately 1175 kilojoules per mole (\(\text{kJ/mol}\))) to that of xenon.
Xenon has a closely comparable ionization energy of about 1170 \(\text{kJ/mol}\), suggesting that \(\text{PtF}_6\) might be potent enough to strip an electron from the xenon atom. Bartlett tested his hypothesis by mixing colorless xenon gas with the deep red vapor of platinum hexafluoride in a glass apparatus. An immediate reaction occurred as the gases combined to precipitate an orange-yellow solid on the sides of the vessel. The spontaneous formation of this solid compound, which was stable at room temperature, provided the first undeniable proof that noble gases could participate in chemical reactions.
The Nature of Xenon Hexafluoroplatinate
The orange-yellow substance synthesized by Bartlett was initially formulated as xenon hexafluoroplatinate (\(\text{XePtF}_6\)), a salt composed of the xenon cation (\(\text{Xe}^+\)) and the hexafluoroplatinate anion (\(\text{PtF}_6^-\)). The reaction succeeded because platinum hexafluoride is one of the most highly oxidizing substances known. This strong oxidizing ability allowed it to overcome the high ionization energy of xenon, effectively removing an electron from the atom.
Xenon, being the heaviest of the noble gases, has the lowest ionization energy, making it the most susceptible to chemical reaction among its group members. Later analysis revealed that the initial product was likely more complex than the simple \(\text{Xe}^+[\text{PtF}_6]^-\) salt. The resulting solid proved to be a mixture of salts, but the fundamental significance remained: a stable compound had been formed from a noble gas. The successful oxidation of xenon confirmed that chemical reactivity was a matter of energy balance between the noble gas and a sufficiently powerful reactant.
The Dawn of Noble Gas Chemistry
The discovery of xenon hexafluoroplatinate immediately led to a burst of new research. Scientists quickly abandoned the outdated “inert gas” label and began an intense search for other noble gas compounds. Within months of Bartlett’s initial report, other chemists successfully synthesized simpler xenon compounds.
These included xenon difluoride (\(\text{XeF}_2\)), xenon tetrafluoride (\(\text{XeF}_4\)), and xenon hexafluoride (\(\text{XeF}_6\)), which were easier to study and isolate than the initial complex salt. The rapid creation of these compounds confirmed that the first synthesis was the opening of an entirely new branch of chemistry. The research soon expanded to include other heavy noble gases, resulting in the successful preparation of compounds involving Krypton and Radon.