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452 Noble-Gas Chemistry

reacting with krypton over a period of several hours to of tetrameric and hexameric rings of virtually undistorted
+
form KrF 2 . An excellent and experimentally simpler pho- square pyramidal XeF cations (see Section III.E) linked
5
tochemical method for the preparation of KrF 2 has been together by fluoride ion bridges.
developed which involves the photolysis of solid kryp- Xenon dichloride (XeCl 2 ) and XeClF have been pre-
◦
ton/liquid fluorine mixtures at −196 C. Krypton difluo- pared by photochemical and electric discharge meth-
ride can routinely be obtained in yields of several grams to ods and have been examined at low temperatures by
several tens of grams using low-temperature photochem- matrix-isolation techniques. The dichloride has a lin-
ical means and generally requires no further purification. ear structure like that of XeF 2 . Evidence for the exis-
It can be stored indefinitely at temperatures of −78 Cor tence of XeCl 2 , XeBr 2 , and XeCl 4 has been obtained
◦
below. However, great care must be exercised in avoid- from 129 Xe M¨ossbauer studies. The compounds were
ing inadvertent contact with organic substances and the obtained by γ -decay of the corresponding 129 I com-
introduction of moisture into solid KrF 2 samples, as the pounds. Owing to their thermal chemical instabilities,
potent oxidant properties of KrF 2 result in the rapid oxi- no dihalide other than the binary fluorides have been
dation of organic compounds and water. For example, vi- prepared in macroscopic amounts. Other examples of
+
olent explosions have been known to occur upon warming Xe(II) Cl bonds are C 6 F 5 XeCl and (C 6 F 5 Xe) 2 Cl (see
of moisture-contaminated bulk samples to or near room Section III.F.3).
temperature. Unstable monohalides of xenon (XeF, XeCl, XeBr, and
Radon fluoride is most conveniently prepared by reac- XeI) have been produced in the gas phase by electron
tion of radon gas with a liquid halogen fluoride (ClF, ClF 3 , bombardment methods and in solid matrices by gamma
ClF 5 , BrF 3 ,orIF 7 ) at room temperature. Millicurie or and ultraviolet irradiation methods. Although short lived
larger amounts of radon react spontaneously with gaseous in the gas phase, these halides are of considerable impor-
or liquid fluorine in a small volume (25 or 50 mL flask) tance as light-emitting species in gas lasers.
within about 30 min. The radon behaves as both activator
and reactant. The product is nonvolatile and hence remains
B. Oxides
in the reaction vessel when the excess reagent is removed
by vacuum distillation. Two oxides of xenon are known: xenon trioxide (XeO 3 )
and xenon tetroxide (XeO 4 ). Xenon trioxide is most ef-
ficiently prepared by the hydrolysis of XeF 6 or by the
III. XENON COMPOUNDS reaction of XeF 6 with HOPOF 2 . The XeO 3 molecule has
a trigonal pyramidal shape [C 3v point symmetry; Xe O,
˚
The principal neutral, cationic, and anionic fluorides; ox- 1.76(3) A], and XeO 4 has a tetrahedral shape in the gas
˚
ide fluorides; and oxides of xenon are listed in Table I phase [T d point symmetry; Xe O, 1.736(2) A]. Xenon
along with their geometries. tetroxide has also been studied in solution by 129 Xe and
131 Xe NMR spectroscopy. Xenon tetroxide is prepared by
the interaction of concentrated sulfuric acid with sodium
A. Halides
or barium perxenate (Na 4 XeO 6 ,Ba 2 XeO 6 ). Both oxides
Xenon fluorides (and XeOF 4 ) and their complexes are are thermodynamically unstable, explosive solids which
the only thermodynamically stable xenon compounds. must be handled with the greatest care. On decomposing
Xenon difluoride (XeF 2 ), xenon tetrafluoride (XeF 4 ), and to the elements, solid XeO 3 and gaseous XeO 4 release
−1
xenon hexafluoride (XeF 6 ) are stable, colorless, crys- 402 and 642 kJ mol , respectively. Xenon trioxide has a
talline solids which can be sublimed under vacuum negligible vapor pressure at room temperature and read-
at 25 C. The mean thermochemical bond energies are ily dissolves in water to give stable solutions containing
◦
−1
−1
−
XeF 2 , 132.3 ± 0.7 kJ mol ; XeF 4 , 130.3 ± 0.5 kJ mol ; mainly molecular XeO 3 and xenic acid anion, HXeO ,
4
−1
and XeF 6 , 125.3 ± 0.7kJmol . Xenon hexafluoride is whichis vanishingly small (K ≈ 3 × 10 −11 for XeO 3aq + 2
+
yellow-green as a liquid or gas. Reports of xenon octaflu- H 2 O l − HXeO − + H 3 O ), except in basic solution
4aq aq
3
oride, XeF 8 , in the early noble-gas chemistry literature (K ≈ 1.5 × 10 for XeO 3aq + OH − HXeO − ). Xenon
−
aq 4aq
remain unsubstantiated. Xenon difluoride is a linear sym- tetroxide is volatile at 25 C, but frequently decomposes
◦
˚
metrical molecule [Xe F, 1.9773(15) A], and XeF 4 is a explosively well before this temperature is reached.
˚
square planar molecule [Xe F, 1.953(4) A]. Experimental
evidence is consistent with a distorted octahedral structure
C. Oxide Fluorides
for gaseous XeF 6 arising from the presence of an extra
pair of nonbonding electrons in the xenon valence shell. The oxide fluoride O(XeF) 2 is isoelectronic with the
+
Solid XeF 6 exists in at least four phases which consist Xe 2 F cation and is also pale yellow in color and has
3