The story of the noble gases started in 1785 in the studies into the composition of air by H Cavendish. During his studies into the separation of air he was left with a residual gas which he was unable to remove chemically which would be identified as being mostly argon a century later. Another noble gas element was discovered from observation of a new absorption line in the spectrum of light from the Sun during the solar eclipse of 1868. JN Lockyer and E Frankland suggested the existence of a new element which they named helium after the Greek word for sun. This spectral line was also observed by L Palmieri in the volcanic gas from Mt Vesuvius in 1881. In work to test the hypothesis that the atomic weights of elements should be multiples of that of hydrogen, Lord Rayleigh consistently observed a slight difference in the densities of nitrogen gas obtained from air and that obtained from ammonia. To try and account for this discrepancy he reacted the nitrogen from air with heated magnesium forming magnesium nitride, Mg₃N₂, a small amount of a dense monatomic gas. This gas, observed to be a different element due to it’s inert nature was called argon from the Greek word for lazy or idle. W Ramsay then suggested that a new group should be added to the periodic table. He then went on, with MW Travers, to isolate and purify the nobel gas elements krypton, neon and xenon by low temperature liquification and distillation of air. He also identified the presence of helium in uranium ores. The remaining element radon was later isolated and studied by Rutherford and Soddy in 1902.

Physical Properties and Uses

The noble gas elements are all monatomic gases and are colourless, odourless and non-polar. Selected properties of the elements are shown in table 1. Weak van der Waal’s forces are the only interactive forces acting between noble gas atoms. As the size of the atoms increase going from He to Rn, the electron clouds around the nuclei become more polarisable leading to a greater degree of interatomic attraction and an increase in the boiling and melting points.

Out of all the nobel gas elements, helium possesses most bizarre properties. Firstly, it can’t be frozen as the other elements can without the application of pressure. It also shows very strange behaviour below a transition temperature called the λ-point at 2.2 K. Above this temperature it appears to be a normal liquid and is called HeI. Below this temperature, HeII, it become a superfluid where it’s viscosity becomes zero. It will spread out to cover any surface connected to it that is also below the λ-point temperature with a layer a few hundred atoms thick. If an empty container is immersed into HeII, the liquid helium will climb it’s sides and fill it until the levels are the same inside and out. Another property of these element, especially helium, is their ability to diffuse through many material such as rubber and PVC. Helium will also diffuse through most forms of glass.

Table 1: Selected properties of the noble gases.

Radon, the heaviest member of the group, is formed from the radioactive decay of other elements in rocks and is radioactive itself. This presents a health hazard in certain areas where there are large amounts of granite rock. Houses in these areas must have thorough ventilation to prevent build up of this carcinogen.

One of the most common uses of helium is a cryogen and also has a more familiar use in balloons. It is also used in deep sea dive breathing gas as a replacement for nitrogen to minimize the chances of suffering the “bends”. Neon’s main use is in illuminated signs. Argon is most frequently used as an inert gas to provide an innocuous atmosphere inside laboratory equipment for air sensitive chemical reactions and is also used in light bulbs. Krypton and xenon have applications in gas lasers.

Inert Gases, Not So Inert

Chemists’ scientific curiosity often leads them to try and do things that are thought to be inherently difficult if not impossible. Early attempts to initiate chemical reactions with noble gases proved unsuccessful seemingly confirming the opinion that they were totally inert. It was not until 1962 that the very first noble gas containing compound was made. In studying the chemistry of the extremely reactive and volatile gas PtF₆, N Bartlett found that accidental exposure to atmospheric oxygen lead to a colour change and the immediate formation of O₂⁺[PtF₆]^-. He noted that the ionisation energy of O₂ was similar to that of Xe and proposed that a similar reaction might take place between Xe and the highly oxidising PtF₆. This was indeed observed with the deep red vapour of PF₆ forming a orange-yellow solid on contact with Xe.

Xe + PtF₆→ Xe⁺[PtF₆]^-

Within a few months the fluoride compounds XeF₄ and XeF₂ had also been made and there is now an extensive body of xenon chemistry known (see table 2). Only the heavier noble gases form isolable compounds and most of these are exclusively flourides and oxides though a few highly unstable compounds with bonds to other elements are known. Radon would be thought to have a richer chemistry than that of xenon, however it’s radioactive nature not only makes is hazardous to work with but radioactive decay products also decompose the compounds.

Table 2: Examples of some xenon oxides and fluorides.

Xenon Fluorides

Xenon difluoride, XeF₂ can be prepared either by reaction of excess Xe with F₂ at 400 °C in a sealed nickel vessel or by the action of sunlight on mixtures of the two gases. XeF₂ will react with further F₂ to give XeF₄ which will react further to give XeF₆. The are the only neutral pure fluorides and an equilibrium exists between the them above 250 °C. XeF₂ is soluble in water but readily undergoes hydrolysis in the presence of base

2XeF₂ + H₂O → 2Xe + 4HF + O₂

It also acts as a mild fluorinating agent and is able to add F₂ across the C=C double bond in alkenes to give difluoroalkanes.

Xenon tetrafluoride, XeF₄, is the easiest of the fluorides to make by heating a 1:1.5 mixture of Xe and F₂ at 400 °C under about 6 atmospheres pressure. It instantly hydrolyses in water to give a mixture of products including XeO₃ which is highly explosive therefore any reaction to make XeF₄ requires rigorous exclusion of moisture. It is also a much stronger fluorinating agent than XeF₂