Chemical compounds containing at least one xenon atom
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 the electronegative atoms 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.[1]
Three fluorides are known: XeF 2, XeF 4, and XeF 6. XeF is theorized to be unstable.[2] These are the starting points for the synthesis of almost all xenon compounds.
The solid, crystalline difluoride XeF 2 is formed when a mixture of fluorine and xenon gases is exposed to ultraviolet light.[3] The ultraviolet component of ordinary daylight is sufficient.[4] Long-term heating of XeF 2 at high temperatures under an NiF 2 catalyst yields XeF 6.[5] Pyrolysis of XeF 6 in the presence of NaF yields high-purity XeF 4.[6]
The xenon fluorides behave as both fluoride acceptors and fluoride donors, forming salts that contain such cations as XeF+ and Xe 2F+ 3, and anions such as XeF− 5, XeF− 7, and XeF2− 8. The green, paramagnetic Xe+ 2 is formed by the reduction of XeF 2 by xenon gas.[1]
XeF 2 also forms coordination 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, and silicon or carbon 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 a van der Waals molecule of weakly bound Xe atoms and Cl 2 molecules and not a real compound.[8] Theoretical calculations indicate that the linear molecule XeCl 2 is less stable than the van der Waals complex.[9]Xenon tetrachloride and xenon dibromide are more unstable that they cannot be synthesized by chemical reactions. They were created by radioactive decay of 129 ICl− 4 and 129 IBr− 2, respectively.[10][11]
Oxides and oxohalides
Three oxides of xenon are known: xenon trioxide (XeO 3) and xenon tetroxide (XeO 4), 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.[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 by infrared spectroscopy in solid argon.[14]
Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF 6:[15]
XeF 6 + 3 H 2O → XeO 3 + 6 HF
XeO 3 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.[16]
Barium perxenate, when treated with concentrated sulfuric acid, yields gaseous xenon tetroxide:[7]
Ba 2XeO 6 + 2 H 2SO 4 → 2 BaSO 4 + 2 H 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, including XeOF 2, XeOF 4, XeO 2F 2, and XeO 3F 2. XeOF 2 is formed by reacting OF 2 with xenon gas at low temperatures. It may also be obtained by partial hydrolysis of XeF 4. It disproportionates at −20°C into XeF 2 and XeO 2F 2.[17]XeOF 4 is formed by the partial hydrolysis of XeF 6...[18]
XeF 6 + H 2O → XeOF 4 + 2 HF
...or the reaction of XeF 6 with sodium perxenate, Na 4XeO 6. The latter reaction also produces a small amount of XeO 3F 2.
XeO 2F 2 is also formed by partial hydrolysis of XeF 6.[19]
XeF 6 + 2 H 2O → XeO 2F 2 + 4 HF
XeOF 4 reacts with CsF to form the XeOF− 5 anion,[17][20] while XeOF3 reacts with the alkali metal fluorides KF, RbF and CsF to form the XeOF− 4 anion.[21]
Other compounds
Xenon can be directly bonded to a less electronegative element than fluorine or oxygen, particularly carbon.[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 include F–Xe–N(SO 2F) 2 and F–Xe–BF 2. The latter is synthesized from dioxygenyl tetrafluoroborate, O 2BF 4, at −100°C.[17][24]
An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe2+ 4, which contains Xe–Au bonds.[25] This ion occurs in the compound AuXe 4(Sb 2F 11) 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.[26]
The compound Xe 2Sb 2F 11 contains a Xe–Xe bond, the longest element-element bond known (308.71pm = 3.0871Å).[27]
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.[28] In 2008, Khriachtchev et al. reported the preparation of HXeOXeH by the photolysis of water within a cryogenic xenon matrix.[29]Deuterated molecules, HXeOD and DXeOH, have also been produced.[30]
In addition to compounds where xenon forms a chemical bond, xenon can form clathrates—substances where xenon atoms or pairs are trapped by the crystalline lattice of another compound. One example is xenon hydrate (Xe·5+3⁄4H2O), where xenon atoms occupy vacancies in a lattice of water molecules.[31] This clathrate has a melting point of 24°C.[32] The deuterated version of this hydrate has also been produced.[33] Another example is xenon hydride (Xe(H2)8), in which xenon pairs (dimers) are trapped inside solid hydrogen.[34] Such clathrate hydrates can occur naturally under conditions of high pressure, such as in Lake Vostok underneath the Antarctic ice sheet.[35] Clathrate formation can be used to fractionally distill xenon, argon and krypton.[36]
Xenon can also form endohedral fullerene compounds, where a xenon atom is trapped inside a fullerene molecule. The xenon atom trapped in the fullerene can be observed by 129Xe nuclear magnetic resonance (NMR) spectroscopy. Through the sensitive chemical 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.[37]
When xenon atoms are in the ground energy state, they repel each other and will not form a bond. When xenon atoms becomes energized, however, they can form an excimer (excited dimer) until the electrons return to the ground state. This entity is formed because the xenon atom tends to complete the outermost electronic shell by adding an electron from a neighboring xenon atom. The typical lifetime of a xenon excimer is 1–5nanoseconds, and the decay releases photons with wavelengths of about 150 and 173nm.[38][39] Xenon can also form excimers with other elements, such as the halogensbromine, chlorine, and fluorine.[40]
Related Research Articles
The noble gases are the naturally occurring members of group 18 of the periodic table: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Under standard conditions, these elements are odorless, colorless, monatomic gases with very low chemical reactivity and cryogenic boiling points.
Xenon is a chemical element; it has symbol Xe and atomic number 54. It is a dense, colorless, odorless noble gas found in Earth's atmosphere in trace amounts. Although generally unreactive, it can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized.
In chemistry, noble gas compounds are chemical compounds that include an element from the noble gases, group 18 of the periodic table. Although the noble gases are generally unreactive elements, many such compounds have been observed, particularly involving the element xenon.
Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:
Dioxygen difluoride is a compound of fluorine and oxygen with the molecular formula O2F2. It can exist as an orange-colored solid which melts into a red liquid at −163 °C (110 K). It is an extremely strong oxidant and decomposes into oxygen and fluorine even at −160 °C (113 K) at a rate of 4% per day—its lifetime at room temperature is thus extremely short. Dioxygen difluoride reacts vigorously with nearly every chemical it encounters (including ordinary ice) leading to its onomatopoeic nickname FOOF (a play on its chemical structure and its explosive tendencies).
Xenon tetrafluoride is a chemical compound with chemical formula XeF 4. It was the first discovered binary compound of a noble gas. It is produced by the chemical reaction of xenon with fluorine:
Xenon hexafluoride is a noble gas compound with the formula XeF6. It is one of the three binary fluorides of xenon that have been studied experimentally, the other two being XeF2 and XeF4. All known are exergonic and stable at normal temperatures. XeF6 is the strongest fluorinating agent of the series. It is a colorless solid that readily sublimes into intensely yellow vapors.
Platinum hexafluoride is the chemical compound with the formula PtF6, and is one of seventeen known binary hexafluorides. It is a dark-red volatile solid that forms a red gas. The compound is a unique example of platinum in the +6 oxidation state. With only four d-electrons, it is paramagnetic with a triplet ground state. PtF6 is a strong fluorinating agent and one of the strongest oxidants, capable of oxidising xenon and O2. PtF6 is octahedral in both the solid state and in the gaseous state. The Pt-F bond lengths are 185 picometers.
Silver(II) fluoride is a chemical compound with the formula AgF2. It is a rare example of a silver(II) compound - silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.
Xenon trioxide is an unstable compound of xenon in its +6 oxidation state. It is a very powerful oxidizing agent, and liberates oxygen from water slowly, accelerated by exposure to sunlight. It is dangerously explosive upon contact with organic materials. When it detonates, it releases xenon and oxygen gas.
Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF 2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.
Selenium tetrafluoride (SeF4) is an inorganic compound. It is a colourless liquid that reacts readily with water. It can be used as a fluorinating reagent in organic syntheses (fluorination of alcohols, carboxylic acids or carbonyl compounds) and has advantages over sulfur tetrafluoride in that milder conditions can be employed and it is a liquid rather than a gas.
Xenon oxytetrafluoride is an inorganic chemical compound. It is an unstable colorless liquid with a melting point of −46.2 °C that can be synthesized by partial hydrolysis of XeF 6, or the reaction of XeF 6 with silica or NaNO 3:
Krypton difluoride, KrF2 is a chemical compound of krypton and fluorine. It was the first compound of krypton discovered. It is a volatile, colourless solid at room temperature. The structure of the KrF2 molecule is linear, with Kr−F distances of 188.9 pm. It reacts with strong Lewis acids to form salts of the KrF+ and Kr 2F+ 3 cations.
The dioxygenyl ion, O+ 2, is a rarely-encountered oxycation in which both oxygen atoms have a formal oxidation state of +1/2. It is formally derived from oxygen by the removal of an electron:
A hexafluoride is a chemical compound with the general formula QXnF6, QXnF6m−, or QXnF6m+. Many molecules fit this formula. An important hexafluoride is hexafluorosilicic acid (H2SiF6), which is a byproduct of the mining of phosphate rock. In the nuclear industry, uranium hexafluoride (UF6) is an important intermediate in the purification of this element.
The tetrafluoroammonium cation is a positively charged polyatomic ion with chemical formula NF+ 4. It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine. Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF 4, trifluoramine oxide ONF 3 and the tetrafluoroborate BF− 4 anion.
Organoxenon chemistry is the study of the properties of organoxenon compounds, which contain carbon to xenon chemical bonds. The first organoxenon compounds were divalent, such as (C6F5)2Xe. The first tetravalent organoxenon compound, [C6F5XeF2][BF4], was synthesized in 2004. So far, more than one hundred organoxenon compounds have been researched.
Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.
Radon compounds are chemical compounds formed by the element radon (Rn). Radon is a noble gas, i.e. a zero-valence element, and is chemically not very reactive. The 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer. Because radon is a gas under normal circumstances, and its decay-chain parents are not, it can readily be extracted from them for research.
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↑ 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.
1 2 Tramšek, 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.
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↑ 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.
↑ 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.
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