G. J. Moody UWlST Cardiff. United Kingdom
A Decade of Xenon Chemistry
Until the end of the nineteenth century air was held to comprise nitrogen, oxygen, carbon dioxide, and water vapor. The isolation ( I ) of argon (Ar) from the air by Rayleigh and Ramsay in 1895 was startling enough; even more so was its complete lack of activity towards a variety of chemicals. Moissan, (2) for example, failed to achieve reaction even on sparking with fluorine. Indeed, its very isolation depended on this total chemical inactivity. Within a few years five other gaseous elements (He, Ne, Kr, Xe, and Rn) were discovered. All apparently lacked normal chemical behavior and these six remarkable elements were justifiably classified as Inert Gases. Moreover, they fitted perfectly between the electronegative halogens and electropositive alkali metals in Mendeleev's Periodic Table. The early electronic valency theories strengthened the concept of inertness hy emphasizing the significance of the stable electronic octet. In fact, it is the electronic duplet, not the octet, which is so fundamentally important. Periodically the few serious attempts to react xenon or krypton with chlorine or fluorine also failed. Thus Yost and Kaye (3) could not confirm any definite reaction between xenon and fluorine subjected to an electrical discharge. Their attempt to combine xenon and chlorine nhotochemicallv fared no better. Ironicallv this latter experiment was not performed for xenon and fluorine. It is a tantalizine thoueht that had thev done so then xenon difluoride would Lave certainly L e n made 40 years ago, since Strenn and Strene (4) obtained XeFz crystals within hours of exposing the constituent elements to sunlight in a dry Pyrex flask a t room temperature. Had Ramsay been able to give Moissan as much xenon as he had argon, it is conceivable that xenon fluorides would have been made 70 years ago, and surely Ruff (who first made IF,) would have succeeded even if Moissan had failed. Xenon chemistry would then have evolved at a leisurely pace. hen in 1962 the well-entrenched ine~t-gasconcept, at least for xenon, was shattered. Bartlett and Lohmann (5) had found oxygen ( I = 12.2 eV) to react a t room temperature with PtFe to form Oz+PtFe-. Bartlett (6) reasoned that xenon ( I = 12.13 eV) would react likewise. It did in a visually rapid dramatic fashion according to Xe(g)
+ PtFdred vapor)
-
XetPtFB-(yellowarangevapor)
That this reaction is more complicated is unimportant for it triggered tremendous interest and research efforts. This has been neatly put by Anderson (7) thus: "If Ramsay enjoyed exclusive prospecting rights in virgin territory, Bartlett started a eold rush." Todav more is known about xenon chemistry t h k some far more"common elements. Naturallv. other hexafluorides (RhFa and RuFd were tried and i&nd to rapidly react with ;enon, hut-the simultaneous reduction of the hexafluorides implied them to be acting as fluorinating agents, and that perhaps fluorine itself would also react with xenon. The now famous experiment (August 2nd, 1962) was conducted by Claassen, Malm, and Selig (8) a t the Argonne National Laboratories in which xenon and fluorine (1:5 ratio) were heated in a nickel can a t 400'C and 6 atm for 1 hr. Xenon consumption was essentially complete! The can contents 828
/
Journal of Chemlcal Educatlon
Noble Gas Compounds Establlrhed 51nca 1962
XeCI1; XeFa; X e h ; xaF3.XeF1
readily sublimed into a glass vessel as brilliant colorless XeFl crystals. Xenon di and hexafluorides were also quickly synthesized from the elements, but xenon-oxygen compounds cannot yet be directly synthesized. These binary xenon fluorides constitute the only starting materials for preparing the compounds listed in the table by complexation or hydrolysis 2XeF6
+ 4Naf + 160H-
-+ + -- +
+
+
Na,Xe08 0, Xe 12F 8HP ZXeF, + SiO, ZXeOF, + SiF, BalXeOs + 2H80, XeO, 2BaS0, + 2H,O XeO,F, XeOF, XeF, + XeO,
-
+
Other interesting products are realized by substitution with highly electronegative ligands in acetonitrile media
-
+
2HF XeF, + 2HOS0,F -+ Xe(OSO,F), FXeOTeF, HF XeF2 + HOTeF, Xe(OTeFSX WOS0,F -+ XdOSO,F), + 2HOTeF,
+
+
Attempts to effect the reaction produced O(TeFs)n, S20eFz, HF, 02, and Xe instead. Xenon compounds vary considerably in their stability. Thus xenon dichloride can only be prepared by the ingenious Nelson-Pimentel matrix isolation technique (9): xenon trioxide (AHorrai= 400 kJ mole-=) is an explosive comparable with TNT, while dirubidium and dicesium octafluoroxenate(V1) are stable to around 400°C. Perxe-
nates, which can he handled as conveniently as most lahoratory chemicals, and usually stable to 200°C, constitute very powerful oxidizing agents which for example instantly oxidize manganese(Il) to permanganate. Xenon chemistry which is based on the even II, VI, IV, and VIII oxidation states closely resembles that of iodine and, to some extent, osmium and tellurium. Other Noble Gas Compounds
Radon ( I = 10.75 eV) is more easily oxidized than xenon ( I = 12.13 eV); for example, it reacts spontaneously at r w m temperature with fluorine or chlorine trifluoride, whereas xenon requires thermal or photochemical initiation. Thus the element should have an extensive chemistry hut the 3.84 day half-life of the most stable 222-isotope presents formidable experimental problems. However, the existence, hut not stoichiometries, of both volatile and nonvolatile radon fluorides has heen demonstrated on a tracer level (10). Krypton ( I = 14.00 eV) is more difficult to oxidize than xenon or radon, and to date only two compounds are authenticated. The difluoride can only he kept a t Dry Ice temperature; whereas the KrFz.2SbFs complex is more stable, slowly decomposing at 25°C. Xenon and krypton, but not argon, react with fluorine using the Pimentel matrix isolation technique. This strongly implies that no sensible stable compounds of the lighter elements argon (I = 15.76 eV), or of course neon (I = 21.56 eV), or helium (I = 24.59 eV) will he prepared by proven present day methods. Therefore, the description of these three lighter elements as inert gases can still be justified. This inertness is particularly useful since argon is employed on a large scale for industrial blanketing as in titanium ~ ~ ~nroduction. ~ ~ ~ ~ ~ ~ Two important features emerge from the composition of the established noble eas comoounds (see the table). First ionization potentials, which &ate the comparative ease of removing outer orbital electrons, decrease steadily from helium through radon. This trend accords with the chemical reactivity of the three heaviest gases, and also underlines the inequality of the electron configuration among the whole family. Secondly the atoms, or groups, associated with say xenon in these compounds are among the most electronegative known, and considerable hond polarity might be exaccords pected. Thus a polarity of -o.5F-Xe+1-F-o.5 with many physical properties of xenon difluoride.
.~~~ ~~
rine) are arrayed in five pairs-two honding and three lone pairs as shown (see the figure). These three lone pairs in the trigonal plane exert equivalent repulsions on both honding pairs to give a linear molecule. Similarly, no distortion is expected or found in xenon tetrafluoride. The slight distortion observed in xenon oxide tetrafluoride ( lone-oairbonding pair > honding pair-honding pair: In bromine trifluoride the ten valence electrons are arrayed in three bonding and two nonbonding pairs. The lone pairs should, and do, cause some distortion (3') in the axial plane. Gillespie (11) has used the theory with considerable success to predict the symmetries for a number of xenon compounds(see the figure). In xenon and krypton difluorides the ten valence electrons (eight from the noble gas and one from each fluo-
The geometries of selected bromine end xenon compounds. Lone pairs are shaded dumbbells.
Volume 51. Number 10. October 1074
/
629
in x , y, and z coordinates predicts octahedral symmetry for xenon hexafluoride. By any standards xenon hexafluoride is an unusual molecule since few of its properties can he reconciled with the Oh symmetry established for many known hexafluorides; few small molecules have been studied in such detail. Gillespie predicted a distorted model on the basis of VSEPR which seems to he unrealistic in view of the reported (12) dipole moment (
