Chemical compounds containing at least one xenon atom
Xenon compounds are compounds containing the elementxenon (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 theelectronegative atoms fluorine or oxygen. The chemistry of xenon in eachoxidation state is analogous to that of the neighboring elementiodine in the immediately lower oxidation state.[1]
Threefluorides are known:XeF 2,XeF 4, andXeF 6. XeF is theorized to be unstable.[2] These are the starting points for the synthesis of almost all xenon compounds.
The solid, crystalline difluorideXeF 2 is formed when a mixture offluorine and xenon gases is exposed to ultraviolet light.[3] The ultraviolet component of ordinary daylight is sufficient.[4] Long-term heating ofXeF 2 at high temperatures under anNiF 2 catalyst yieldsXeF 6.[5]Pyrolysis ofXeF 6 in the presence ofNaF yields high-purityXeF 4.[6]
The xenon fluorides behave as both fluoride acceptors and fluoride donors, forming salts that contain such cations asXeF+ andXe 2F+ 3, and anions such asXeF− 5,XeF− 7, andXeF2− 8. The green, paramagneticXe+ 2 is formed by the reduction ofXeF 2 by xenon gas.[1]
XeF 2 also formscoordination complexes with transition metal ions. More than 30 such complexes have been synthesized and characterized.[5]
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, andsilicon orcarbon tetrachloride,[7] is reported to be an endothermic, colorless, crystalline compound that decomposes into the elements at 80 °C. However,XeCl 2 may be merely avan der Waals molecule of weakly bound Xe atoms andCl 2 molecules and not a real compound.[8] Theoretical calculations indicate that the linear moleculeXeCl 2 is less stable than the van der Waals complex.[9]Xenon tetrachloride andxenon dibromide are more unstable that they cannot be synthesized by chemical reactions. They werecreated by radioactive decay of129 ICl− 4 and129 IBr− 2, respectively.[10][11]
Three oxides of xenon are known:xenon trioxide (XeO 3) andxenon tetroxide (XeO 4), both of which are dangerously explosive and powerful oxidizing agents, andxenon dioxide (XeO2), which was reported in 2011 with acoordination number of four.[12] XeO2 forms when xenon tetrafluoride is poured over ice. Its crystal structure may allow it to replace silicon in silicate minerals.[13] The XeOO+ cation has been identified byinfrared spectroscopy in solidargon.[14]
Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis ofXeF 6:[15]
XeF 6 + 3H 2O →XeO 3 + 6 HF
XeO 3 is weakly acidic, dissolving in alkali to form unstablexenate salts containing theHXeO− 4 anion. These unstable salts easilydisproportionate into xenon gas andperxenate salts, containing theXeO4− 6 anion.[16]
Barium perxenate, when treated with concentratedsulfuric acid, yields gaseous xenon tetroxide:[7]
Ba 2XeO 6 + 2H 2SO 4 → 2BaSO 4 + 2H 2O +XeO 4
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, includingXeOF 2,XeOF 4,XeO 2F 2, andXeO 3F 2.XeOF 2 is formed by reactingOF 2 with xenon gas at low temperatures. It may also be obtained by partial hydrolysis ofXeF 4. It disproportionates at −20 °C intoXeF 2 andXeO 2F 2.[17]XeOF 4 is formed by the partial hydrolysis ofXeF 6...[18]
XeF 6 +H 2O →XeOF 4 + 2HF
...or the reaction ofXeF 6 with sodium perxenate,Na 4XeO 6. The latter reaction also produces a small amount ofXeO 3F 2.
XeO 2F 2 is also formed by partial hydrolysis ofXeF 6.[19]
XeF 6 + 2H 2O →XeO 2F 2 + 4HF
XeOF 4 reacts withCsF to form theXeOF− 5 anion,[17][20] while XeOF3 reacts with thealkali metal fluoridesKF,RbF and CsF to form theXeOF− 4 anion.[21]
Xenon can be directly bonded to a less electronegative element than fluorine or oxygen, particularlycarbon.[22] Electron-withdrawing groups, such as groups with fluorine substitution, are necessary to stabilize these compounds.[16] Numerous such compounds have been characterized, including:[17][23]
C 6F 5–Xe+ –N≡C–CH 3, where C6F5 is the pentafluorophenyl group.
[C 6F 5] 2Xe
C 6F 5–Xe–C≡N
C 6F 5–Xe–F
C 6F 5–Xe–Cl
C 2F 5–C≡C–Xe+
[CH 3] 3C–C≡C–Xe+
C 6F 5–XeF+ 2
(C 6F 5Xe) 2Cl+
Other compounds containing xenon bonded to a less electronegative element includeF–Xe–N(SO 2F) 2 andF–Xe–BF 2. The latter is synthesized fromdioxygenyl tetrafluoroborate,O 2BF 4, at −100 °C.[17][24]
An unusual ion containing xenon is thetetraxenonogold(II) cation,AuXe2+ 4, which contains Xe–Au bonds.[25] This ion occurs in the compoundAuXe 4(Sb 2F 11) 2, and is remarkable in having direct chemical bonds between two notoriously unreactive atoms, xenon andgold, with xenon acting as a transition metal ligand. A similar mercury complex (HgXe)(Sb3F17) (formulated as [HgXe2+][Sb2F11–][SbF6–]) is also known.[26] Xenon reversibly complexes gaseousM(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.[27]
The compoundXe 2Sb 2F 11 contains a Xe–Xe bond, the longest element-element bond known (308.71 pm = 3.0871 Å).[28]
In 1995, M. Räsänen and co-workers, scientists at theUniversity of Helsinki inFinland, announced the preparation of xenon dihydride (HXeH), and later xenon hydride-hydroxide (HXeOH), hydroxenoacetylene (HXeCCH), and other Xe-containing molecules.[29] In 2008, Khriachtchevet al. reported the preparation of HXeOXeH by thephotolysis of water within acryogenic xenon matrix.[30]Deuterated molecules, HXeOD and DXeOH, have also been produced.[31]
In addition to compounds where xenon forms achemical bond, xenon can formclathrates—substances where xenon atoms or pairs are trapped by thecrystalline lattice of another compound. One example isxenon hydrate (Xe·5+3⁄4H2O), where xenon atoms occupy vacancies in a lattice of water molecules.[32] This clathrate has amelting point of 24 °C.[33] Thedeuterated version of this hydrate has also been produced.[34] Another example is xenonhydride (Xe(H2)8), in which xenon pairs (dimers) are trapped insidesolid hydrogen.[35] Suchclathrate hydrates can occur naturally under conditions of high pressure, such as inLake Vostok underneath theAntarctic ice sheet.[36] Clathrate formation can be used to fractionally distill xenon, argon and krypton.[37]
Xenon can also formendohedral fullerene compounds, where a xenon atom is trapped inside afullerene molecule. The xenon atom trapped in the fullerene can be observed by129Xenuclear magnetic resonance (NMR) spectroscopy. Through the sensitivechemical shift of the xenon atom to its environment, chemical reactions on the fullerene molecule can be analyzed. These observations are not without caveat, however, because the xenon atom has an electronic influence on the reactivity of the fullerene.[38]
When xenon atoms are in theground energy state, they repel each other and will not form a bond. When xenon atoms becomes energized, however, they can form anexcimer (excited dimer) until the electrons return to theground state. This entity is formed because the xenon atom tends to complete the outermostelectronic shell by adding an electron from a neighboring xenon atom. The typical lifetime of a xenon excimer is 1–5 nanoseconds, and the decay releasesphotons withwavelengths of about 150 and 173 nm.[39][40] Xenon can also form excimers with other elements, such as thehalogensbromine,chlorine, andfluorine.[41]
^Dean H Liskow; Henry F Schaefer III; Paul S Bagus; Bowen Liu (1973). "Probable nonexistence of xenon monofluoride as a chemically bound species in the gas phase".J Am Chem Soc.95 (12):4056–57.doi:10.1021/ja00793a042.
^Weeks, James L.; Chernick, Cedric; Matheson, Max S. (1962). "Photochemical Preparation of Xenon Difluoride".Journal of the American Chemical Society.84 (23):4612–13.doi:10.1021/ja00882a063.
^Streng, L. V.; Streng, A. G. (1965). "Formation of Xenon Difluoride from Xenon and Oxygen Difluoride or Fluorine in Pyrex Glass at Room Temperature".Inorganic Chemistry.4 (9):1370–71.doi:10.1021/ic50031a035.
^abTramšek, Melita; Žemva, Boris (December 5, 2006). "Synthesis, Properties and Chemistry of Xenon(II) Fluoride".Acta Chimica Slovenica.53 (2):105–16.doi:10.1002/chin.200721209.
^Proserpio, Davide M.; Hoffmann, Roald; Janda, Kenneth C. (1991). "The xenon-chlorine conundrum: van der Waals complex or linear molecule?".Journal of the American Chemical Society.113 (19):7184–89.doi:10.1021/ja00019a014.
^Richardson, Nancy A.; Hall, Michael B. (1993). "The potential energy surface of xenon dichloride".The Journal of Physical Chemistry.97 (42):10952–54.doi:10.1021/j100144a009.
^Bell, C.F. (2013).Syntheses and Physical Studies of Inorganic Compounds. Elsevier Science. p. 143.ISBN978-1-48328060-8.
^Cockett, A.H.; Smith, K.C.; Bartlett, N. (2013).The Chemistry of the Monatomic Gases: Pergamon Texts in Inorganic Chemistry. Elsevier Science. p. 292.ISBN978-1-48315736-8.
^Christe, K. O.; Dixon, D. A.; Sanders, J. C. P.; Schrobilgen, G. J.; Tsai, S. S.; Wilson, W. W. (1995). "On the Structure of the [XeOF5]− Anion and of Heptacoordinated Complex Fluorides Containing One or Two Highly Repulsive Ligands or Sterically Active Free Valence Electron Pairs".Inorg. Chem.34 (7):1868–1874.doi:10.1021/ic00111a039.
^Christe, K. O.; Schack, C. J.; Pilipovich, D. (1972). "Chlorine trifluoride oxide. V. Complex formation with Lewis acids and bases".Inorg. Chem.11 (9):2205–2208.doi:10.1021/ic50115a044.
^Frohn, H.; Theißen, Michael (2004). "C6F5XeF, a versatile starting material in xenon–carbon chemistry".Journal of Fluorine Chemistry.125 (6):981–988.doi:10.1016/j.jfluchem.2004.01.019.
^Goetschel, Charles T.; Loos, Karl R. (1972). "Reaction of xenon with dioxygenyl tetrafluoroborate. Preparation of FXe-BF2".Journal of the American Chemical Society.94 (9):3018–3021.doi:10.1021/ja00764a022.
^Grochala, Wojciech (Oct 2007) [12 April 2007]. "Atypical compounds of gases, which have been called 'noble'".Chemical Society Reviews.36 (10). Royal Society of Chemistry. p. 1638 and fn. 53–55.doi:10.1039/b702109g – viaCiteSeerX.
^Khriachtchev, Leonid; Isokoski, Karoliina; Cohen, Arik; Räsänen, Markku; Gerber, R. Benny (2008). "A Small Neutral Molecule with Two Noble-Gas Atoms: HXeOXeH".Journal of the American Chemical Society.130 (19):6114–8.doi:10.1021/ja077835v.PMID18407641.
^Pettersson, Mika; Khriachtchev, Leonid; Lundell, Jan; Räsänen, Markku (1999). "A Chemical Compound Formed from Water and Xenon: HXeOH".Journal of the American Chemical Society.121 (50):11904–905.doi:10.1021/ja9932784.
^Ikeda, Tomoko; Mae, Shinji; Yamamuro, Osamu; Matsuo, Takasuke; Ikeda, Susumu; Ibberson, Richard M. (November 23, 2000). "Distortion of Host Lattice in Clathrate Hydrate as a Function of Guest Molecule and Temperature".Journal of Physical Chemistry A.104 (46):10623–30.Bibcode:2000JPCA..10410623I.doi:10.1021/jp001313j.
^McKay, C. P.; Hand, K. P.; Doran, P. T.; Andersen, D. T.; Priscu, J. C. (2003). "Clathrate formation and the fate of noble and biologically useful gases in Lake Vostok, Antarctica".Geophysical Research Letters.30 (13): 35.Bibcode:2003GeoRL..30.1702M.doi:10.1029/2003GL017490.S2CID20136021.
