The idle ramblings of a Jack of some trades, Master of none

Jul 3, 2011


Chemistry, one can argue, is physics applied to electron transfer. How do compounds form? In a most basic way, this happens when one element hands over some of its electrons to another (creating what is known as a bond). Electrons, as is well known, are negatively charged particles that orbit the positively charged nucleus of every atom in an element. Opposite charges attract, and so the electrons keep revolving around the nucleus. In fact, they do so at various 'distances' from the nucleus, called 'shells'. The reason why some elements are happy to hand over electrons and others are pleased to accept them is that all elements yearn to achieve 'perfection', and this occurs when there are (usually) 8 electrons in their outermost shell.

As it happens, some elements are perfect in that they already have the full complement of electrons in their outermost shells. These elements therefore are non-reactive. They are called 'noble', and there are several of them - all gases - helium, neon, argon, krypton, xenon, radon. They resolutely eschew the company of any other elements. There are no naturally occurring compounds involving the noble ones. And for much of modern chemistry, it was widely held that there was no way the noble gases would ever react.

In 1962, though, a particularly enthusiastic man called Neil Bartlett set out to create a compound that contained a noble gas.

Neil Bartlett (1932-2008) (MSU Gallery of Chemists' Mini-portraits)
Now here's an important point. Some elements have atoms that are larger than other elements' atoms. The larger the atoms, the farther away the outer shells are from the hold of the nucleus. So they are easier to dislodge and share with other elements. Bartlett realised that the technology available to him was insufficient to force the smaller noble elements into compounds, so he chose one of the larger ones - xenon (chemical symbol Xe).

He realised that the amount of energy required to remove an electron from xenon was about the same as that to ionize O2. Oxygen, usually, was an electron borrower, but in the presence of a highly reactive substance known as platinum hexafluoride, it could be forced to release an electron. If oxygen, then why not xenon, which had roughly the same ionization energy? Bartlett managed to produce an orange crystal: xenon hexafluoroplatinate (Xe+[PtF6]) Interestingly, the reaction could take place at room temperature.

Encouraged by this finding, other chemists attempted to create compounds with the lighter noble gases. The next gas to succumb was krypton (Kr), and did so barely a year after Bartlett's epochal work. Krypton was so resilient, however, that it needed to be cooled down to -151°C, when its electrons slowed enough that fluorine could rip them away. Here's how the creators of krypton tetrafluoride (Kr+F4) reported their methodology:
The experimental setup (a reaction vessel of volume approximately 650 mm3, with copper electrodes 2.0 cm in diameter and 7 cm apart) and the experimental conditions (current of 24 to 37 ma, 700 to 2200 volts) were the same as in the earlier investigations (5). The mixtures of Kr and F2 (1 and 2 volumes respectively, to within 0.1 percent) was admitted, at a pressure of 7 to 12 mm-Hg, into the discharge vessel, which had been cooled to 84°K to 86°K by mixtures of liquid O2 and N2. In a successful experiment, 500 cm3 of the mixture of Kr and F2 (at normal temperature and pressure) was completely converted to 1.15g of KrF4 in 4.0 hours. 2
It took another 37 years for another noble gas to form a compound. In 2000, Finnish scientists achieved the feat with argon (Ar). Here's how:
It was an experiment of Fabergé delicacy, requiring solid argon; hydrogen gas; fluorine gas; a highly reactive starter compound, cesium iodide, to get the reaction going; and well-timed bursts of ultraviolet light, all set to bake at a frigid -445°F. When things got a little warmer, the argon compound collapsed.

Nevertheless, below that temperature argon fluorohydride was a durable crystal. The Finnish scientists announced the feat in a paper with a refreshingly accessible title for a scientific work, "A Stable Argon Compound." 3

  1. Mark Sampson, 'Neil Bartlett and Reactive Noble Gases', American Chemical Society.
  2. Grosse, A.V. et al, 'Krypton Tetrafluoride: Preparation and Some Properties', Science, 15 Mar 1963.
  3. Sam Kean, The Disappearing Spoon: and Other True Tales of Madness, Love and the History of the World from the Periodic Table of the Elements, Little, Brown, 2010.


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