The noble gases — helium, neon, argon, krypton and xenon — all occur as minor constituents of air. Although this seems to make them very accessible, they remained unknown until the end of the nineteenth century. The most abundant, argon, was actually isolated in 1785 by Henry Cavendish, but he did not recognize this unknown component of air as a new element. It was only in 1894 that Sir William Ramsay and Lord Rayleigh jointly announced its discovery. This marked the beginning of an extraordinary scientific adventure, which led Ramsay and his co-workers to isolate, in the space of few years, an entire group of new elements.

Among the group 18 elements, neon (from the greek νέον, 'new', a name suggested by Ramsay's 13-year-old son), krypton and xenon were obtained by fractionation of liquid air using an apparatus that had just been invented by the engineers William Hampson and Carl von Linde, which efficiently produced large amounts of liquid gases — a wonderful example of pure and applied science working in concert. The fraction containing neon was distilled in June 1898. Element 10, once isolated, presented peculiar spectroscopic lines, including the bright reddish–orange lights that now brighten up our city tours by night. The same red emission is also behind the helium–neon lasers employed, for example, in barcode scanners, CD players and medical applications such as laser eye surgery and the analysis of blood cells.

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In 1912, J. J. Thomson observed that the canal rays (beams of positive ions) obtained from ionized neon followed two distinct trajectories when passed through a magnetic and an electric field. He deduced the presence of neon atoms with two different atomic masses, 20Ne and 22Ne, thus discovering isotopes of a stable element. The separation of ions by their mass was soon improved by Arthur Dempster and Francis Aston, and developed into the modern technique of mass spectrometry.

Naturally, chemists attempted reactions with the noble gases, but early attempts were unsuccessful. No negative results, however, proved more informative: this reluctance to react became a founding principle of the modern theories of chemical bonding, which considers the elements' valence electron shells. In the case of the noble gases, their complete shells lead them to being inert.

Yet, unfazed, chemists didn't give up on reacting the noble gases. If chemical bonds form by electron sharing or donation, it was reasonable to expect that the inertness of these elements would progressively decrease from helium to xenon. Moving down the periodic table in this order, polarizability increases and ionization potential progressively decreases, to reach values comparable to those of commonly oxidizable molecules. Guided by these arguments, on a Friday afternoon of March 1962, all alone in his lab, Neil Bartlett succeeded in oxidizing xenon using platinum hexafluoride. Xenon chemistry soon greatly increased, and is now well established. Several krypton compounds also went on to be prepared, as well as one argon compound1 (the triatomic HArF); no helium and neon compounds have yet been reported.

By the same token, neon should be more reactive than helium. Yet according to theoretical investigations, neutral and even anionic species such as HHeF, H3CHeF, (LiF)2(HeO) and FHeX− (X = O, S, Se) are metastable structures featuring covalent helium bonds, whereas the neon analogues of these species are predicted to be unbound. These calculations are in agreement with the fact that complexes of neon with neutral metal acceptors, including the compounds NeAuF and NeBeS recently detected in cold matrices2,3, are also in general less stable than their helium counterparts. There are also examples of helium and neon cations that feature this reversed order of stability4.

Neon is bigger than helium, and possesses occupied p orbitals. This is thought to produce less effective electrostatic interactions and higher orbital repulsions, which typically make the neon compounds either unstable or only marginally stable, although the contributions of these factors are still to be further investigated. Chemists studying neon thus face two challenges: the experimental preparation, but also the accurate theoretical prediction of its compounds.

There have been suggestions to shift helium to group 2 of the periodic table, next to hydrogen and just above beryllium. Supporting arguments are the isoelectronic analogy (it has two electrons in its outside shell), and the anchoring of otherwise concealed periodic regularities. The lower stability of neon compounds compared with helium ones is in line with this proposal; moving helium would make neon occupy the top position in that column of the periodic table, which would suit well its situation of the most inert noble gas.