Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers in September 1898, shortly after their discovery of the elements krypton and neon. They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον xénon, neuter singular form of ξένος xénos, meaning 'foreign(er)', 'strange(r)', or 'guest'. In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere to be one part in 20 million.
In 1939, American physician Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, and discovered that this caused the divers to perceive a change in depth. From his results, he deduced that xenon gas could serve as an anesthetic. Although Russian toxicologist Nikolay V. Lazarev apparently studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice. Xenon was first used as a surgical anesthetic in 1951 by American anesthesiologist Stuart C. Cullen, who successfully used it with two patients.
An acrylic cube specially prepared for element collectors containing liquefied xenon
In November 1989, IBM scientists demonstrated a technology capable of manipulating individual atoms. The program, called IBM in atoms, used a scanning tunneling microscope to arrange 35 individual xenon atoms on a substrate of chilled crystal of nickel to spell out the three letter company initialism. It was the first time atoms had been precisely positioned on a flat surface.
A layer of solid xenon floating on top of liquid xenon inside a high voltage apparatus.
Liquid (featureless) and crystalline solid Xe nanoparticles produced by implanting Xe+ ions into aluminium at room temperature.
Xenon has atomic number 54; that is, its nucleus contains 54 protons. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Liquid xenon has a high polarizability due to its large atomic volume, and thus is an excellent solvent. It can dissolve hydrocarbons, biological molecules, and even water. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is greater than the average density of granite, 2.75 g/cm3. Under gigapascals of pressure, xenon forms a metallic phase.
Solid xenon changes from face-centered cubic (fcc) to hexagonal close packed (hcp) crystal phase under pressure and begins to turn metallic at about 140 GPa, with no noticeable volume change in the hcp phase. It is completely metallic at 155 GPa. When metallized, xenon appears sky blue because it absorbs red light and transmits other visible frequencies. Such behavior is unusual for a metal and is explained by the relatively small width of the electron bands in that state.
Liquid or solid xenon nanoparticles can be formed at room temperature by implanting Xe+ ions into a solid matrix. Many solids have lattice constants smaller than solid Xe. This results in compression of the implanted Xe to pressures that may be sufficient for its liquefaction or solidification.
Xenon is a member of the zero-valence elements that are called noble or inert gases. It is inert to most common chemical reactions (such as combustion, for example) because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.
Xenon is obtained commercially as a by-product of the separation of air into oxygen and nitrogen. After this separation, generally performed by fractional distillation in a double-column plant, the liquid oxygen produced will contain small quantities of krypton and xenon. By additional fractional distillation, the liquid oxygen may be enriched to contain 0.1–0.2% of a krypton/xenon mixture, which is extracted either by absorption onto silica gel or by distillation. Finally, the krypton/xenon mixture may be separated into krypton and xenon by further distillation.
Worldwide production of xenon in 1998 was estimated at 5,000–7,000 m3. Because of its scarcity, xenon is much more expensive than the lighter noble gases—approximate prices for the purchase of small quantities in Europe in 1999 were 10 €/L for xenon, 1 €/L for krypton, and 0.20 €/L for neon, while the much more plentiful argon costs less than a cent per liter.
Within the Solar System, the nucleon fraction of xenon is 1.56 × 10−8, for an abundance of approximately one part in 630 thousand of the total mass. Xenon is relatively rare in the Sun's atmosphere, on Earth, and in asteroids and comets. The abundance of xenon in the atmosphere of planet Jupiter is unusually high, about 2.6 times that of the Sun. This abundance remains unexplained, but may have been caused by an early and rapid buildup of planetesimals—small, subplanetary bodies—before the heating of the presolar disk. (Otherwise, xenon would not have been trapped in the planetesimal ices.) The problem of the low terrestrial xenon may be explained by covalent bonding of xenon to oxygen within quartz, reducing the outgassing of xenon into the atmosphere.
Naturally occurring xenon is composed of seven stableisotopes: 126Xe, 128–132Xe, and 134Xe. The isotopes 126Xe and 134Xe are predicted by theory to undergo double beta decay, but this has never been observed so they are considered stable. In addition, more than 40 unstable isotopes that have been studied. The longest lived of these isotopes are the primordial124Xe, which undergoes double electron capture with a half-life of 1.8 × 1022 yr, and 136Xe, which undergoes double beta decay with a half-life of 2.11 × 1021 yr.129Xe is produced by beta decay of 129I, which has a half-life of 16 million years. 131mXe, 133Xe, 133mXe, and 135Xe are some of the fission products of 235U and 239Pu, and are used to detect and monitor nuclear explosions.
Because a 129Xe nucleus has a spin of 1/2, and therefore a zero electricquadrupole moment, the 129Xe nucleus does not experience any quadrupolar interactions during collisions with other atoms, and the hyperpolarization persists for long periods even after the engendering light and vapor have been removed. Spin polarization of 129Xe can persist from several seconds for xenon atoms dissolved in blood to several hours in the gas phase and several days in deeply frozen solid xenon. In contrast, 131Xe has a nuclear spin value of 3⁄2 and a nonzero quadrupole moment, and has t1 relaxation times in the millisecond and second ranges.
Some radioactive isotopes of xenon (for example, 133Xe and 135Xe) are produced by neutron irradiation of fissionable material within nuclear reactors.135Xe is of considerable significance in the operation of nuclear fission reactors. 135Xe has a huge cross section for thermal neutrons, 2.6×106barns, and operates as a neutron absorber or "poison" that can slow or stop the chain reaction after a period of operation. This was discovered in the earliest nuclear reactors built by the American Manhattan Project for plutonium production. However, the designers had made provisions in the design to increase the reactor's reactivity (the number of neutrons per fission that go on to fission other atoms of nuclear fuel).135Xe reactor poisoning was a major factor in the Chernobyl disaster. A shutdown or decrease of power of a reactor can result in buildup of 135Xe, with reactor operation going into a condition known as the iodine pit.
Under adverse conditions, relatively high concentrations of radioactive xenon isotopes may emanate from cracked fuel rods, or fissioning of uranium in cooling water.
Because the half-life of 129I is comparatively short on a cosmological time scale (16 million years), this demonstrated that only a short time had passed between the supernova and the time the meteorites had solidified and trapped the 129I. These two events (supernova and solidification of gas cloud) were inferred to have happened during the early history of the Solar System, because the 129I isotope was likely generated shortly before the Solar System was formed, seeding the solar gas cloud with isotopes from a second source. This supernova source may also have caused collapse of the solar gas cloud.
In a similar way, xenon isotopic ratios such as 129Xe/130Xe and 136Xe/130Xe are a powerful tool for understanding planetary differentiation and early outgassing. For example, the atmosphere of Mars shows a xenon abundance similar to that of Earth (0.08 parts per million) but Mars shows a greater abundance of 129Xe than the Earth or the Sun. Since this isotope is generated by radioactive decay, the result may indicate that Mars lost most of its primordial atmosphere, possibly within the first 100 million years after the planet was formed. In another example, excess 129Xe found in carbon dioxide well gases from New Mexico is believed to be from the decay of mantle-derived gases from soon after Earth's formation.
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.
Three fluorides are known: XeF 2, XeF 4, and XeF 6. XeF is theorized to be unstable. 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. The ultraviolet component of ordinary daylight is sufficient. Long-term heating of XeF 2 at high temperatures under an NiF 2 catalyst yields XeF 6. Pyrolysis of XeF 6 in the presence of NaF yields high-purity XeF 4.
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.
XeF 2 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, with the exception of dichloride XeCl2 and XeCl4, the other halides are not known. 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, XeCl 2 may be merely a van der Waals molecule of weakly bound Xe atoms and Cl 2 molecules and not a real compound. Theoretical calculations indicate that the linear molecule XeCl 2 is less stable than the van der Waals complex. Xenon tetrachloride is more unstable that can’t synthesized by chemical reaction.It was created by radioactive129 ICl− 4 decay.
Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF 6:
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.
Barium perxenate, when treated with concentrated sulfuric acid, yields gaseous xenon tetroxide:
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.XeOF 4 is formed by the partial hydrolysis of XeF 6, or the reaction of XeF 6 with sodium perxenate, Na 4XeO 6. The latter reaction also produces a small amount of XeO 3F 2. XeOF 4 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.
Xenon can be directly bonded to 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:
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.
An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe2+ 4, which contains Xe–Au bonds. 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.
The compound Xe 2Sb 2F 11 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.
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.
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 1–5 nanoseconds, and the decay releases photons with wavelengths of about 150 and 173 nm. Xenon can also form excimers with other elements, such as the halogensbromine, chlorine, and fluorine.
Although xenon is rare and relatively expensive to extract from the Earth's atmosphere, it has a number of applications.
The individual cells in a plasma display contain a mixture of xenon and neon ionized with electrodes. The interaction of this plasma with the electrodes generates ultraviolet photons, which then excite the phosphor coating on the front of the display.
Xenon is used as a "starter gas" in high pressure sodium lamps. It has the lowest thermal conductivity and lowest ionization potential of all the non-radioactive noble gases. As a noble gas, it does not interfere with the chemical reactions occurring in the operating lamp. The low thermal conductivity minimizes thermal losses in the lamp while in the operating state, and the low ionization potential causes the breakdown voltage of the gas to be relatively low in the cold state, which allows the lamp to be more easily started.
Xenon interacts with many different receptors and ion channels, and like many theoretically multi-modal inhalation anesthetics, these interactions are likely complementary. Xenon is a high-affinity glycine-site NMDA receptor antagonist. However, xenon is different from certain other NMDA receptor antagonists in that it is not neurotoxic and it inhibits the neurotoxicity of ketamine and nitrous oxide, while actually producing neuroprotective effects. Unlike ketamine and nitrous oxide, xenon does not stimulate a dopamine efflux in the nucleus accumbens.
Like nitrous oxide and cyclopropane, xenon activates the two-pore domain potassium channel TREK-1. A related channel TASK-3 also implicated in the actions of inhalation anesthetics is insensitive to xenon. Xenon inhibits nicotinic acetylcholine α4β2 receptors which contribute to spinally mediated analgesia. Xenon is an effective inhibitor of plasma membrane Ca2+ ATPase. Xenon inhibits Ca2+ ATPase by binding to a hydrophobic pore within the enzyme and preventing the enzyme from assuming active conformations.
Xenon is a competitive inhibitor of the serotonin 5-HT3 receptor. While neither anesthetic nor antinociceptive, this reduces anesthesia-emergent nausea and vomiting.
Xenon induces robust cardioprotection and neuroprotection through a variety of mechanisms. Through its influence on Ca2+, K+, KATP\HIF, and NMDA antagonism, xenon is neuroprotective when administered before, during and after ischemic insults. Xenon is a high affinity antagonist at the NMDA receptor glycine site. Xenon is cardioprotective in ischemia-reperfusion conditions by inducing pharmacologic non-ischemic preconditioning. Xenon is cardioprotective by activating PKC-epsilon and downstream p38-MAPK. Xenon mimics neuronal ischemic preconditioning by activating ATP sensitive potassium channels. Xenon allosterically reduces ATP mediated channel activation inhibition independently of the sulfonylurea receptor1 subunit, increasing KATP open-channel time and frequency.
Inhaling a xenon/oxygen mixture activates production of the transcription factorHIF-1-alpha, which may lead to increased production of erythropoietin. The latter hormone is known to increase red blood cell production and athletic performance. Reportedly, doping with xenon inhalation has been used in Russia since 2004 and perhaps earlier. On August 31, 2014, the World Anti Doping Agency (WADA) added xenon (and argon) to the list of prohibited substances and methods, although no reliable doping tests for these gases have yet been developed. In addition, effects of xenon on erythropoietin production in humans have not been demonstrated, so far.
Xenon, particularly hyperpolarized 129Xe, is a useful contrast agent for magnetic resonance imaging (MRI). In the gas phase, it can image cavities in a porous sample, alveoli in lungs, or the flow of gases within the lungs. Because xenon is soluble both in water and in hydrophobic solvents, it can image various soft living tissues.
Because of the xenon atom's large, flexible outer electron shell, the NMR spectrum changes in response to surrounding conditions and can be used to monitor the surrounding chemical circumstances. For instance, xenon dissolved in water, xenon dissolved in hydrophobic solvent, and xenon associated with certain proteins can be distinguished by NMR.
Hyperpolarized xenon can be used by surface chemists. Normally, it is difficult to characterize surfaces with NMR because signals from a surface are overwhelmed by signals from the atomic nuclei in the bulk of the sample, which are much more numerous than surface nuclei. However, nuclear spins on solid surfaces can be selectively polarized by transferring spin polarization to them from hyperpolarized xenon gas. This makes the surface signals strong enough to measure and distinguish from bulk signals.
Liquid xenon is used in calorimeters to measure gamma rays, and as a detector of hypothetical weakly interacting massive particles, or WIMPs. When a WIMP collides with a xenon nucleus, theory predicts it will impart enough energy to cause ionization and scintillation. Liquid xenon is useful for these experiments because its density makes dark matter interaction more likely and it permits a quiet detector through self-shielding.
Because they are strongly oxidative, many oxygen–xenon compounds are toxic; they are also explosive (highly exothermic), breaking down to elemental xenon and diatomic oxygen (O2) with much stronger chemical bonds than the xenon compounds.
Xenon gas can be safely kept in normal sealed glass or metal containers at standard temperature and pressure. However, it readily dissolves in most plastics and rubber, and will gradually escape from a container sealed with such materials. Xenon is non-toxic, although it does dissolve in blood and belongs to a select group of substances that penetrate the blood–brain barrier, causing mild to full surgical anesthesia when inhaled in high concentrations with oxygen.
The speed of sound in xenon gas (169 m/s) is less than that in air because the average velocity of the heavy xenon atoms is less than that of nitrogen and oxygen molecules in air. Hence, xenon vibrates more slowly in the vocal cords when exhaled and produces lowered voice tones, an effect opposite to the high-toned voice produced in helium. Like helium, xenon does not satisfy the body's need for oxygen, and it is both a simple asphyxiant and an anesthetic more powerful than nitrous oxide; consequently, and because xenon is expensive, many universities have prohibited the voice stunt as a general chemistry demonstration. The gas sulfur hexafluoride is similar to xenon in molecular weight (146 versus 131), less expensive, and though an asphyxiant, not toxic or anesthetic; it is often substituted in these demonstrations.
Dense gases such as xenon and sulfur hexafluoride can be breathed safely when mixed with at least 20% oxygen. Xenon at 80% concentration along with 20% oxygen rapidly produces the unconsciousness of general anesthesia (and has been used for this, as discussed above). Breathing mixes gases of different densities very effectively and rapidly so that heavier gases are purged along with the oxygen, and do not accumulate at the bottom of the lungs. There is, however, a danger associated with any heavy gas in large quantities: it may sit invisibly in a container, and a person who enters an area filled with an odorless, colorless gas may be asphyxiated without warning. Xenon is rarely used in large enough quantities for this to be a concern, though the potential for danger exists any time a tank or container of xenon is kept in an unventilated space.
^Albert, J. B.; Auger, M.; Auty, D. J.; Barbeau, P. S.; Beauchamp, E.; Beck, D.; Belov, V.; Benitez-Medina, C.; Bonatt, J.; Breidenbach, M.; Brunner, T.; Burenkov, A.; Cao, G. F.; Chambers, C.; Chaves, J.; Cleveland, B.; Cook, S.; Craycraft, A.; Daniels, T.; Danilov, M.; Daugherty, S. J.; Davis, C. G.; Davis, J.; Devoe, R.; Delaquis, S.; Dobi, A.; Dolgolenko, A.; Dolinski, M. J.; Dunford, M.; et al. (2014). "Improved measurement of the 2νββ half-life of 136Xe with the EXO-200 detector". Physical Review C. 89. arXiv:1306.6106. Bibcode:2014PhRvC..89a5502A. doi:10.1103/PhysRevC.89.015502.
^Staff (2007). "Xenon". Columbia Electronic Encyclopedia (6th ed.). Columbia University Press. Retrieved 2007-10-23.
^Rabinovich, Viktor Abramovich; Vasserman, A. A.; Nedostup, V. I.; Veksler, L. S. (1988). Thermophysical properties of neon, argon, krypton, and xenon. Washington. 10. Washington, DC: Hemisphere Publishing Corp. Bibcode:1988wdch...10.....R. ISBN0-89116-675-0.—National Standard Reference Data Service of the USSR. Volume 10.
^ abFreemantle, Michael (August 25, 2003). "Chemistry at its Most Beautiful". Chemical & Engineering News. Vol. 81 no. 34. pp. 27–30. doi:10.1021/cen-v081n034.p027.
^Graham, L.; Graudejus, O.; Jha N.K.; Bartlett, N. (2000). "Concerning the nature of XePtF6". Coordination Chemistry Reviews. 197 (1): 321–334. doi:10.1016/S0010-8545(99)00190-3.
^Holleman, A. F.; Wiberg, Egon (2001). Bernhard J. Aylett (ed.). Inorganic Chemistry. translated by Mary Eagleson and William Brewer. San Diego: Academic Press. ISBN0-12-352651-5.; translation of Lehrbuch der Anorganischen Chemie, founded by A. F. Holleman, continued by Egon Wiberg, edited by Nils Wiberg, Berlin: de Gruyter, 1995, 34th edition, ISBN3-11-012641-9.
^Mahaffy, P. R.; Niemann, H. B.; Alpert, A.; Atreya, S. K.; Demick, J.; Donahue, T. M.; Harpold, D. N.; Owen, T. C. (2000). "Noble gas abundance and isotope ratios in the atmosphere of Jupiter from the Galileo Probe Mass Spectrometer". Journal of Geophysical Research. 105 (E6): 15061–15072. Bibcode:2000JGR...10515061M. doi:10.1029/1999JE001224.
^Mass fraction calculated from the average mass of an atom in the solar system of about 1.29 atomic mass units
^Heymann, D.; Dziczkaniec, M. (March 19–23, 1979). Xenon from intermediate zones of supernovae. Proceedings 10th Lunar and Planetary Science Conference. Houston, Texas: Pergamon Press, Inc. pp. 1943–1959. Bibcode:1979LPSC...10.1943H.
^Pignatari, M.; Gallino, R.; Straniero, O.; Davis, A. (2004). "The origin of xenon trapped in presolar mainstream SiC grains". Memorie della Societa Astronomica Italiana. 75: 729–734. Bibcode:2004MmSAI..75..729P.
^Beer, H.; Kaeppeler, F.; Reffo, G.; Venturini, G. (November 1983). "Neutron capture cross-sections of stable xenon isotopes and their application in stellar nucleosynthesis". Astrophysics and Space Science. 97 (1): 95–119. Bibcode:1983Ap&SS..97...95B. doi:10.1007/BF00684613.
^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): 4056–4057. 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): 4612–4613. 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): 1370–1371. doi:10.1021/ic50031a035.
^ abTramšek, Melita; Žemva, Boris (December 5, 2006). "Synthesis, Properties and Chemistry of Xenon(II) Fluoride". Acta Chimica Slovenica. 53 (2): 105–116. 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): 7184–7189. 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): 10952–10954. doi:10.1021/j100144a009.
^Bell, C.F. (2013). Syntheses and Physical Studies of Inorganic Compounds. Elsevier Science. p. 143. ISBN9781483280608.
^Cockett, A.H.; Smith, K.C.; Bartlett, N. (2013). The Chemistry of the Monatomic Gases: Pergamon Texts in Inorganic Chemistry. Elsevier Science. p. 292. ISBN9781483157368.
^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): 11904–11905. 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): 10623–10630. Bibcode:2000JPCA..10410623I. doi:10.1021/jp001313j.
^Kleppe, Annette K.; Amboage, Mónica; Jephcoat, Andrew P. (2014). "New high-pressure van der Waals compound Kr(H2)4 discovered in the krypton-hydrogen binary system". Scientific Reports. 4: 4989. Bibcode:2014NatSR...4E4989K. doi:10.1038/srep04989.
^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.
^ abBanks, P.; Franks, N. P.; Dickinson, R. (2010). "Competitive inhibition at the glycine site of the N-methyl-D-aspartate receptor mediates xenon neuroprotection against hypoxia-ischemia". Anesthesiology. 112 (3): 614–22. doi:10.1097/ALN.0b013e3181cea398. PMID20124979.
^Yamakura, T.; Harris, R. A. (2000). "Effects of gaseous anesthetics nitrous oxide and xenon on ligand-gated ion channels. Comparison with isoflurane and ethanol". Anesthesiology. 93 (4): 1095–101. doi:10.1097/00000542-200010000-00034. PMID11020766.
^Rashid, M. H.; Furue, H.; Yoshimura, M.; Ueda, H. (2006). "Tonic inhibitory role of α4β2 subtype of nicotinic acetylcholine receptors on nociceptive transmission in the spinal cord in mice". Pain. 125 (1–2): 125–35. doi:10.1016/j.pain.2006.05.011. PMID16781069.
^Nickalls, R.W.D.; Mapleson, W.W. (August 2003). "Age-related iso-MAC charts for isoflurane, sevoflurane and desflurane in man". British Journal of Anaesthesia. 91 (2): 170–174. doi:10.1093/bja/aeg132. PMID12878613.
^Wolber, J.; Rowland, I. J.; Leach, M. O.; Bifone, A. (1998). "Intravascular delivery of hyperpolarized 129Xenon for in vivo MRI". Applied Magnetic Resonance. 15 (3–4): 343–352. doi:10.1007/BF03162020.
^Baltás, E.; Csoma, Z.; Bodai, L.; Ignácz, F.; Dobozy, A.; Kemény, L. (2006). "Treatment of atopic dermatitis with the xenon chloride excimer laser". Journal of the European Academy of Dermatology and Venereology. 20 (6): 657–60. doi:10.1111/j.1468-3083.2006.01495.x. PMID16836491.
^Luhmer, M.; Dejaegere, A.; Reisse, J. (1989). "Interpretation of the solvent effect on the screening constant of Xe-129". Magnetic Resonance in Chemistry. 27 (10): 950–952. doi:10.1002/mrc.1260271009.
^Raftery, Daniel; MacNamara, Ernesto; Fisher, Gregory; Rice, Charles V.; Smith, Jay (1997). "Optical Pumping and Magic Angle Spinning: Sensitivity and Resolution Enhancement for Surface NMR Obtained with Laser-Polarized Xenon". Journal of the American Chemical Society. 119 (37): 8746–8747. doi:10.1021/ja972035d.
^Gaede, H. C.; Song, Y. -Q.; Taylor, R. E.; Munson, E. J.; Reimer, J. A.; Pines, A. (1995). "High-field cross polarization NMR from laser-polarized xenon to surface nuclei". Applied Magnetic Resonance. 8 (3–4): 373–384. doi:10.1007/BF03162652.
^Brazzle, J. D.; Dokmeci, M. R.; Mastrangelo, C. H. (July 28 – August 1, 1975). Modeling and Characterization of Sacrificial Polysilicon Etching Using Vapor-Phase Xenon Difluoride. Proceedings 17th IEEE International Conference on Micro Electro Mechanical Systems (MEMS). Maastricht, Netherlands: IEEE. pp. 737–740. ISBN978-0-7803-8265-7.
^169.44 m/s in xenon (at 0 °C and 107 kPa), compared to 344 m/s in air. See: Vacek, V.; Hallewell, G.; Lindsay, S. (2001). "Velocity of sound measurements in gaseous per-fluorocarbons and their mixtures". Fluid Phase Equilibria. 185 (1–2): 305–314. doi:10.1016/S0378-3812(01)00479-4.
^Yamaguchi, K.; Soejima, K.; Koda, E.; Sugiyama, N (2001). "Inhaling Gas With Different CT Densities Allows Detection of Abnormalities in the Lung Periphery of Patients With Smoking-Induced COPD". Chest. 120 (6): 1907–16. doi:10.1378/chest.120.6.1907. PMID11742921.