Noble gas compound

In chemistry, noble gas compounds are chemical compounds that include an element from the noble gases, group 8 or 18 of the periodic table. Although the noble gases are generally unreactive elements, many such compounds have been observed, particularly involving the element xenon. From the standpoint of chemistry, the noble gases may be divided into two groups: the relatively reactive krypton (ionisation energy 14.0 eV), xenon (12.1 eV), and radon (10.7 eV) on one side, and the very unreactive argon (15.8 eV), neon (21.6 eV), and helium (24.6 eV) on the other. Consistent with this classification, Kr, Xe, and Rn form compounds that can be isolated in bulk at or near standard temperature and pressure, whereas He, Ne, Ar have been observed to form true chemical bonds using spectroscopic techniques, but only when frozen into a noble gas matrix at temperatures of 40 K (−233 °C; −388 °F) or lower, in supersonic jets of noble gas, or under extremely high pressures with metals. The heavier noble gases have more electron shells than the lighter ones. Hence, the outermost electrons are subject to a shielding effect from the inner electrons that makes them more easily ionized, since they are less strongly attracted to the positively-charged nucleus. This results in an ionization energy low enough to form stable compounds with the most electronegative elements, fluorine and oxygen, and even with less electronegative elements such as nitrogen and carbon under certain circumstances.

History and background When the family of noble gases was first identified at the end of the nineteenth century, none of them were observed to form any compounds and so it was initially believed that they were all inert gases (as they were then called) which could not form compounds. With the development of atomic theory in the early twentieth century, their inertness was ascribed to a full valence shell of electrons which render them very chemically stable and nonreactive. All noble gases have full s and p outer electron shells (except helium, which has no p sublevel), and so do not form chemical compounds easily. Their high ionization energy and negative electron affinity explain their non-reactivity. Immediately after the discovery of noble gases, chemists attempted to produce their compounds. The early attempts were failures. Moissan tried to react helium and fluorine at room temperature with electric spark. Berthelot announced a compound of helium with benzene and carbon disulfide using electric discharge. Boomer announced WHe2 formed under intense electric discharge from a tungsten filament in low-pressure helium. Morrison thought that helium, if irradiated so that one electron is pushed to a higher orbit, would behave similar to hydrogen. Consequently, he predicted that radioactive elements might form helides, and claimed to succeed in forming a compound of helium with lead-214, and a compound of helium with bismuth-214. See for an extensive listing of failed attempts before Bartlett's 1962 success. See Helium_compounds#Discredited_or_unlikely_observations for a list of false discoveries of helium compounds. Walther Kossel in 1916 predicted theoretically, on the basis of ionization energy, that xenon fluoride and krypton fluoride can be made. Partly based on Kossel, Andreas von Antropoff argued in 1924 theoretically that the noble gases should be placed in group 8b, rather than group 18, and thus potentially reactive. Until 1933, he also attempted to experimentally produce nobel gas compounds, unsuccessfully. In 1933, Linus Pauling predicted that the heavier noble gases would be able to form compounds with fluorine and oxygen. Specifically, he predicted the existence of krypton hexafluoride (KrF6) and xenon hexafluoride (XeF6), speculated that XeF8 might exist as an unstable compound, and suggested that xenic acid would form perxenate salts. Quantum-chemical calculations subsequently supported Pauling's speculation, suggesting that bonding in a (then-hypothetical) noble gas compound would resemble bonding in the well-known trihalogenide ions, although these were ignored by the broader chemistry community. Pauling later retracted his hypothesis, claming in 1961 that "Xenon is completely unreactive chemically". On Pauling's suggestion, Yost and Kaye in 1933 attempted and failed to make xenon react with fluorine, though Rudolf Hoppe in 1962 would use a modified version of their approach to make XeF2 for the first time. In June 1962, Neil Bartlett gave the first creditable report of a noble gas compound. Bartlett had noticed that the highly oxidising compound platinum hexafluoride ionised O2 to O+2. As the ionisation energy of O2 to O+2 (1165 kJ mol−1) is nearly equal to the ionisation energy of Xe to Xe+ (1170 kJ mol−1), he tried the reaction of Xe with PtF6. This yielded a crystalline product, xenon hexafluoroplatinate, whose formula was proposed to be Xe+[PtF6]−. It was later shown that the compound is actually more complex, containing both [XeF]+[PtF5]− and [XeF]+[Pt2F11]−. Nonetheless, this was the first real compound of any noble gas. The psychological barrier broken, the first binary noble gas compounds appeared later that year. Bartlett subjected a mixture of xenon and fluorine to high temperature, obtaining xenon tetrafluoride (XeF4). Meanwhile, Rudolf Hoppe, among other groups, synthesized xenon difluoride (XeF2) from the elements. Pauling et al's predictions thus proved quite accurate, although XeF8 appears not only thermodynamically, but kinetically unstable. As of 2022, XeF8 has not been made, and only the octafluoroxenate(VI) anion ([XeF8]2−) observed. Following the first successful synthesis of xenon compounds, synthesis of krypton difluoride (KrF2) was reported in 1963.

True noble gas compounds In this section, the non-radioactive noble gases are considered in decreasing order of atomic weight, which generally reflects the priority of their discovery, and the breadth of available information for these compounds. The radioactive elements radon and oganesson are harder to study and are considered at the end of the section.

Xenon compounds

… excerpt ends here. Continue reading the full article.