Xenon compounds are compounds containing the element xenon (Xe). After Neil Bartlett's discovery in 1962 that xenon can form chemical compounds, a large number of xenon compounds have been discovered and described. Almost all known xenon compounds contain atoms of electronegative fluorine or oxygen. The chemistry of xenon in each oxidation state is analogous to that of the neighboring element iodine in the immediately lower oxidation state.
Halides
Three fluorides are known: XeF2, XeF4, and XeF6. These are the starting points for the synthesis of almost all xenon compounds. Neutral XeF is theorized to be unstable; however, fluoroxenonium pentafluoroplatinate (XeF+PtF−5), more commonly known as xenon hexafluoroplatinate, has been described. The solid, crystalline difluoride XeF2 is formed when a mixture of fluorine and xenon gases is exposed to ultraviolet light. The ultraviolet component of ordinary daylight is sufficient. Long-term heating of XeF2 at high temperatures under an NiF2 catalyst yields XeF6. Pyrolysis of XeF6 in the presence of NaF yields high-purity XeF4. The xenon fluorides behave as both fluoride acceptors and fluoride donors, forming salts that contain such cations as XeF+ and Xe2F+3, and anions such as XeF−5, XeF−7, and XeF2−8. The green, paramagnetic Xe+2 is formed by the reduction of XeF2 by xenon gas. XeF2 also forms coordination complexes with transition metal ions. More than 30 such complexes have been synthesized and characterized. Whereas the xenon fluorides are well characterized, the other halides are not. Xenon dichloride, formed by the high-frequency irradiation of a mixture of xenon, fluorine, and silicon or carbon tetrachloride, is reported to be an endothermic, colorless, crystalline compound that decomposes into the elements at 80 °C. However, XeCl2 may be merely a van der Waals molecule of weakly bound Xe atoms and Cl2 molecules and not a real compound. Theoretical calculations indicate that the linear molecule XeCl2 is less stable than the van der Waals complex. Xenon tetrachloride and xenon dibromide are sufficiently more unstable that they cannot be synthesized by chemical reactions. They were created by radioactive decay of 129ICl−4 and 129IBr−2, respectively.
Oxides and oxohalides Three oxides of xenon are known: xenon trioxide (XeO3) and xenon tetroxide (XeO4), both of which are dangerously explosive and powerful oxidizing agents, and xenon dioxide (XeO2), which was reported in 2011 with a coordination number of four. XeO2 forms when xenon tetrafluoride is poured over ice. Its crystal structure may allow it to replace silicon in silicate minerals. The XeOO+ cation has been identified by infrared spectroscopy in solid argon. Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF6:
XeF6 + 3 H2O → XeO3 + 6 HF XeO3 is weakly acidic, dissolving in alkali to form unstable xenate salts containing the HXeO−4 anion. These unstable salts easily disproportionate into xenon gas and perxenate salts, containing the XeO4−6 anion. Barium perxenate, when treated with concentrated sulfuric acid, yields gaseous xenon tetroxide:
Ba2XeO6 + 2 H2SO4 → 2 BaSO4 + 2 H2O + XeO4 To prevent decomposition, the xenon tetroxide thus formed is quickly cooled into a pale-yellow solid. It explodes above −35.9 °C into xenon and oxygen gas, but is otherwise stable. A number of xenon oxyfluorides are known, including XeOF2, XeOF4, XeO2F2, and XeO3F2. XeOF2 is formed by reacting OF2 with xenon gas at low temperatures. It may also be obtained by partial hydrolysis of XeF4. It disproportionates at −20 °C into XeF2 and XeO2F2. XeOF4 is formed by the partial hydrolysis of XeF6...
XeF6 + H2O → XeOF4 + 2 HF ...or the reaction of XeF6 with sodium perxenate, Na4XeO6. The latter reaction also produces a small amount of XeO3F2. XeO2F2 is also formed by partial hydrolysis of XeF6.
XeF6 + 2 H2O → XeO2F2 + 4 HF XeOF4 reacts with CsF to form the XeOF−5 anion, while XeOF3 reacts with the alkali metal fluorides KF, RbF and CsF to form the XeOF−4 anion.
Other compounds Xenon directly bonds with a less-electronegative element than fluorine or oxygen, particularly carbon. Electron-withdrawing groups, such as groups with fluorine substitution, are necessary to stabilize these compounds. Numerous such compounds have been characterized, including:
C6F5–Xe+–N≡C–CH3, where C6F5 is the pentafluorophenyl group. (C6F5)2Xe C6F5–Xe–C≡N C6F5–Xe–F C6F5–Xe–Cl C2F5–C≡C–Xe+ (CH3)3C–C≡C–Xe+ C6F5–XeF+2 (C6F5Xe)2Cl+ Other compounds containing xenon bonded with a less-electronegative element include F–Xe–N(SO2F)2 and F–Xe–BF2. The latter is synthesized from dioxygenyl tetrafluoroborate, O2BF4, at −100 °C. An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe2+4, which contains Xe–Au bonds. This ion occurs in the compound AuXe4(Sb2F11)2, and is remarkable in having direct chemical bonds between two notoriously unreactive atoms, xenon and gold, with xenon acting as a transition metal ligand. A similar mercury complex (HgXe)(Sb3F17) (formulated as [HgXe2+][Sb2F11–][SbF6–]) is also known. Xenon reversibly complexes gaseous M(CO)5, where M=Cr, Mo, or W. p-block metals also bind noble gases: XeBeO has been observed spectroscopically and both XeBeS and FXeBO are predicted stable. The compound Xe2Sb2F11 contains a Xe–Xe bond, the longest element-element bond known (308.71 pm = 3.0871 Å). In 1995, M. Räsänen and co-workers, scientists at the University of Helsinki in Finland, announced the preparation of xenon dihydride (HXeH), and later xenon hydride-hydroxide (HXeOH), hydroxenoacetylene (HXeCCH), and other Xe-containing molecules. In 2008, Khriachtchev et al. reported the preparation of HXeOXeH by the photolysis of water within a cryogenic xenon matrix. Deuterated molecules, HXeOD and DXeOH, have also been produced.
Clathrates and excimers
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