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California Association of Chemistry Teachers
Herberl H. Hyman Argonne Notional Laboratory Argonne, Ninois
The Chemistry of the Noble Gases
This discussion of compounds of the noble gases will be divided into three parts. First, I would like to discuss the chemistry of this family in the periodic table, as if the first experiments on small quantities of xenon isolated from the atmosphere had succeeded in establishing the straightforward reaction of xenon and fluorine (1). The hundred or so publications on the chemistly of the noble gases which have appeared in the last year would have appeared gradually over the 30 succeeding years. Then the unusual xenon platinum hexafluoride adduct noted by Bartlett in May of 1962 (2) would be a rather peculiar and not well understood compound of xenon rather than the first true compound synthesized. I n this portion of this paper I will stress the experimental observations and call attention to the broad spectrum of techniques that have been applied. Then I would like to discuss how theoretical approaches currently used to explain chemical bond formation have met the challenges posed by the specific properties of noble gas compounds as well as the challenge of their very existence. Finally, I would like to review the course of developments in this last year, and see what one might learn from them to tell students. The Chemistry of Noble Gases
The chemical behavior (5) that has been observed for xenon and the other noble gases is easily reconciled with their position in the periodic table. Xenon appears to form compounds only when attached to the most electronegative of elements, primarily fluorine and oxygen. We must therefore assume that whatever binding takes place involves some partial or even substantial transfer of electrons from the noble gas to the combining element. The tendency to lose electrons and therefore to form compounds increases as we go to the heavier elements. Radon exists only as a species This paper ia based on a tak presented at the Aailomar Summer Conference of the California Association of Chemistry Teachem, August, 1963. I t is based on work performed under the auspices of the US.Atomic Energy Commission.
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with rather short half life, and therefore we know very little about the properties of radon compounds. We know that krypton tends to form compounds with fluorine and oxygen with much greater d 8 c u l t y than xenon does, and we believe that argon and the light noble gases are for purposes of conventional chemical manipulations, completely inert. The bulk of our experimental observations are confined to the chemistry of xenon. Xenon exhibits every even valence from two to eight. The difluoride, tetrafluoride, and hexafluoride are well characterized and available in substantial amounts; a good deal is known about their properties. I n the hexavalent states two atoms of fluorine may be substituted by one of oxygen to form successively XeOF4, Xe02F3,and XeOa. The first is apparently quite stable and easily isolated (4); the second has been observed with the aid of a mass spectrometer only as an impurity in mixtures (6); the last is a thermodynamically unstable compound, the end product of hydrolysis of xenon tetra- and hexafluoride, often inadvertently produced in laboratory systems. Since the trioxide explodes on the slightest provocation, its ready production adds considerable spice to research with xenon compounds (6). The corresponding tetravalent oxyfluoride XeOF* has been reported but not adequately studied (7). XeOn is undoubtedly unstable and has not yet been prepared or identified. The octavalent state is best represented in compounds in which the xenon is present as an oxygenated anion (8). As might be expected, the chemistry of the inert gases bears some resemblance to that of the halogens when behaving as positive elements, that is, in forming halogen fluorides and oxides and halogenates and perhalogenates. I will return t o this particular analogy in the last portion of the paper. Though it may be possible to prepare xenon oxygen compounds without going through fluorides, this has not yet been done. The fluorides can be prepared directly from the elements and, of course, with the aid of reactive fluorides. The simplest generalization seems to be that any process which exposes xenon to fluorine atoms,
~rhetlierthe atomic fluorine is formed by heating, electric discharge, ionizing radiation, or some equivalent process in either gaseous fluorine or a fluorine componnd wlrich can be dissociated, will result in the production of a xenon fluoride (9). If the concentration of fluorine is low, the difluoride will he the sole or principal product. Continuing exposure of the difluoride to fluorinating conditions yields the t,etrafluoride; in systems in drich flnorine is in snhstantial excess, the hexafluoride results. The evidrnce for formation of an oct,afluoride under the most severe fluorinahg conditions is not adequate to establish the existence of this compound, hut on the other hand, there does not sccm to he enough evidence t,o rule out rhe possihility of its eveittual isolation in pouderahle amnnnts (10). Observations on the exchange of xenon hexafluoridc and fluorine point strongly to a transitory existence (11). Xenon fluorides are strong fluorinating agents and potentially quite nsefol ones in the important field of production of fluoro chemicals. As yet, very little research has been recorded in this field; but I have no douht t,hat once the basic chemistry of xenon componnds is worked out, there \\-ill he many studies on the use of thesr chemicals ae reagents for introducing fluorine into organic molecules. Each of the xenon fluorides is a white solid a t room temperature (Figs. 1, 2, 3). A very broad range of experimental techniques has been employed in studying the xenon fluorides. I d l oot take the time to describe in detail the infrared spectrum, the Raman spectrum, the NRlR spectrum, or the accumulation of X-ray diffraction, neutron diffraction, and similar studies that have led t o the structural information which I will summarize. I have selected a few illustrations which suggest the diversity of data a t our disposal. The vibration spectra (12) (Figs. 4 and 5) revealed the shape of thc XeFl molecule very soon after its preparation as the first single noble gas fluoride. Similar data for the more complex molec~~le XeOFl (Figs. 6 and 7) showed the striking similarity of the arrangement of fluorines in the two molecules. The table of frequencies and associated vihrations cmphasiaes this similarity (Fig. 8).
Figure 1. Xenon hemfluoride crystdr in a p o l y c h l o r o t r i R ~ ~ r ~ e ~ ytube. lene
Figure 2.
I n additiou to structural informat,ion the ii~frarrd spectrum offers a convenient \yay to study vapor pressures as is shown in Fignre 9 (13). E'~ .k systematic N X R study of the xenon fluorides has heen very helpful in understanding the molecnl~s. Figure 10 summarizes mnch of the information (14) arid compares these conipounds with some related flnorides.
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SPECTRUM Xli4
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Figure 4.
Infrared rpectrvm of XeF, vapor.
Figure 5.
Roman rpeclrum of solid XeF,.
Xenon tetrafluoridesryrtalr.
Figure 3. tube.
Xenon diflvoride crystal in a ~ilico
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The vibrational spectra (Raman and infrared) coupled with electron diffraction data usually serve to define the size and shape of the gaseous molecule. X-ray and neutron diffraction data reveal the arrangement of atoms in the solid. For molecular crystals such as the fluorides of xenon, the arrangement found in the solid is preserved in the vapor phase.
Figure 6.
Infrared spectrum of XeOF4 vapor.
Figure 7.
Romcrn rpestrum of liquid XeOF
