Xenon

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## Contents

* 1 Overview
* 2 History
* 3 Occurrence
* 4 Characteristics
* 5 Isotopes
* 6 Compounds
* 7 Applications
* 7.1 Illumination and optics
* 7.1.1 Gas-discharge lamps
* 7.1.2 Lasers
* 7.2 Anesthesia
* 7.3 Medical imaging
* 7.4 Other
* 8 Precautions
* 9 References
* 10 External links

## Overview

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Xenon (Template:PronEng in the UK, Template:IPA in the US) is a chemical element that has the symbol Xe and atomic number 54\. A colorless, heavy, odorless noble gas, xenon occurs in the earth's atmosphere in trace amounts.[1] Although generally unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized.[2][3][4]

Naturally occurring xenon is made of nine stable isotopes, but there are also over 40 unstable isotopes that undergo radioactive decay. The isotope ratios of xenon are an important tool for studying the early history of the Solar System.[5] Xenon-135 is produced as a result of nuclear fission and acts as a neutron absorber in nuclear reactors.[6]

Xenon is used in flash lamps[7] and arc lamps,[8] and as a general anesthetic.[9] The first excimer laser design used a xenon dimer molecule (Xe2) as its lasing medium,[10] and the earliest laser designs used xenon flash lamps as pumps.[11] Xenon is also being used to search for hypothetical weakly interactive massive particles[12] and as the propellant for ion thrusters in spacecraft.[13]

## History

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Xenon was discovered in England by William Ramsay and Morris Travers on July 12, 1898, shortly after their discovery of the elements krypton and neon. They found it in the residue left over from evaporating components of liquid air.[14][15] Ramsay suggested the name xenon for this gas from the Greek word ξένον [xenon], neuter singular form of ξένος [xenos], meaning foreign, strange, or host.[16][17] In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere as one part in 20 million.[18]

During the 1930s, the engineer Harold Edgerton began exploring strobe light technology for high-speed photography. This led him to the invention of the xenon flash lamp, in which light is generated by sending a brief electrical current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method.[7][19][20]

Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers in 1939. 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 Lazharev, in Russia, apparently studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by J. H. Lawrence, who experimented on mice. Xenon was first used as a surgical anesthetic in 1951 by Stuart C. Cullen, who successfully operated on two patients.[21]

In 1960, the physicist John H. Reynolds discovered that certain meteorites contained an isotopic anomaly in the form of an overabundance of xenon-129. He inferred that this was a decay product of radioactive iodine-129. As the half-life of 129I is 16 million years, this demonstrated that the meteorites were formed during the early history of the Solar System, as the 129I isotope was likely generated before the Solar System was formed.[22][23]

Xenon and the other noble gases were for a long time considered to be completely chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride (PtF6) was a powerful oxidizing agent that could oxidize oxygen (O2) to form dioxygenyl hexafluoroplatinate (O2+[PtF6]−).[24] Since O2 and xenon have almost the same first ionization potential, Bartlett realized that platinum hexafluoride might also be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate.[25][4] Bartlett thought its composition to be Xe+[PtF6]−, although later work has revealed that it was probably a mixture of various xenon-containing salts.[26][27][28] Since then, many other xenon compounds have been discovered,[29] and some compounds of the noble gases argon, krypton, and radon have been identified, including argon fluorohydride (HArF),[30] krypton difluoride (KrF2),[31][32] and radon fluoride.[33]

## Occurrence

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Xenon is a trace gas in Earth's atmosphere, occurring at 0.087±0.001 parts per million (μL/L).[34] It is also found in gases emitted from some mineral springs. Some radioactive species of xenon, for example, 133Xe and 135Xe, are produced by neutron irradiation of fissionable material within nuclear reactors.[2]

Xenon is obtained commercially as a byproduct 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 steps, the liquid oxygen may be enriched to contain 0.1-0.2% of a krypton/xenon mixture, which is extracted either via adsorption onto silica gel or by distillation. Finally, the krypton/xenon mixture may be separated into krypton and xenon via distillation.[35][36] Extraction of a liter of xenon from the atmosphere requires 220 watt-hours of energy.[37] Worldwide production of xenon in 1998 was estimated at 5,000–7,000 m3.[38] Due to its low abundance, 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.[38]

Xenon is relatively rare in the Sun's atmosphere, on Earth, and in the asteroids and comets. The atmosphere of Mars shows a similar xenon abundance to that of Earth: 0.08 parts per million.[39] However, Mars shows a higher proportion of 129Xe than the Earth or the Sun. As 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.[40][41] By contrast, the planet Jupiter has an unusually high abundance of xenon in its atmosphere; about 2.6 times as much as the Sun.[42] This high abundance remains unexplained, but may have been caused by an early and rapid buildup of planetesimals--small, subplanetary bodies--before the presolar disk began to heat up.[43] (Otherwise, xenon would not have been trapped in the planetesimal ices.) Within the Solar System, the nucleon fraction for all isotopes of xenon is 1.56 × 10-8, or one part in 64 million of the total mass.[44] The problem of the low terrestrial xenon may potentially be explained by covalent bonding of xenon to oxygen within quartz, hence reducing the outgassing of xenon into the atmosphere.[45]

Unlike the lower mass noble gases, the normal stellar nucleosynthesis process inside a star does not form xenon. Elements more massive than iron-56 have a net energy cost to produce through fusion, so there is no energy gain for a star to create xenon.[46] Instead, many isotopes of xenon are formed during supernova explosions.[47]

## Characteristics

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Error creating thumbnail: File missingAn electron shell diagram for xenon. Note the eight electrons in the outer shell.

An atom of xenon is defined as having a nucleus with 54 protons. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the surface density of the Earth's atmosphere, 1.217 kg/m3.[48] As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point.[49] Under the same conditions, the density of solid xenon, 3.640 g/cm3, is larger than the average density of granite, 2.75 g/cm3.[49] Using gigapascals of pressure, xenon has been forced into a metallic phase.[50]

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 is completely filled with eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.[51] However, xenon can be oxidized by powerful oxidizing agents, and many xenon compounds have been synthesized.

In a gas-filled tube, xenon emits a blue or lavenderish glow when the gas is excited by electrical discharge. Xenon emits a band of emission lines that span the visual spectrum,[52] but the most intense lines occur in the region of blue light, which produces the coloration.[53]

## Isotopes

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Main article: Isotopes of xenon

Naturally occurring xenon is made of nine stable isotopes. The isotopes 124Xe, 134Xe and 136Xe are predicted to undergo double beta decay, but this has never been observed so they are considered to be stable.[54][55] Besides these stable forms, there are over 40 unstable isotopes that have been studied. 129Xe is produced by beta decay of 129I, which has a half-life of 16 million years, while 131mXe, 133Xe, 133mXe, and 135Xe are some of the fission products of both 235U and 239Pu,[56] and therefore used as indicators of nuclear explosions. The various isotopes of xenon are produced from supernova explosions,[47] red giant stars that have exhausted the hydrogen at their cores and entered the asymptotic giant branch, classical novae explosions[57] and the radioactive decay of elements such as iodine, uranium and plutonium.[56]

The artificial isotope 135Xe is of considerable significance in the operation of nuclear fission reactors. 135Xe has a huge cross section for thermal neutrons, 2.6×106 barns,[6] so it acts 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. Fortunately 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).[58]

Under adverse conditions, relatively high concentrations of radioactive xenon isotopes may be found emanating from nuclear reactors due to the release of fission products from cracked fuel rods,[59] or fissioning of uranium in cooling water.[60]

Because xenon is a tracer for two parent isotopes, xenon isotope ratios in meteorites are a powerful tool for studying the formation of the solar system. The iodine-xenon method of dating gives the time elapsed between nucleosynthesis and the condensation of a solid object from the solar nebula. Xenon isotopic ratios such as 129Xe/130Xe and 136Xe/130Xe are also a powerful tool for understanding terrestrial differentiation and early outgassing.[5] Excess 129Xe found in carbon dioxide well gases from New Mexico was believed to be from the decay of mantle-derived gases soon after Earth's formation.[61][56]

## Compounds

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Xenon tetrafluoride

Xenon hexafluoroplatinate was the first chemical compound of xenon, synthesized in 1962.[25] Following this, many additional compounds of xenon have been discovered. These include xenon difluoride (XeF2), xenon tetrafluoride (XeF4), xenon hexafluoride (XeF6), xenon tetroxide (XeO4), and sodium perxenate (Na4XeO6). A highly explosive compound, xenon trioxide (XeO3), has also been made. Most of the more than 80[62][63] xenon compounds found to date contain electro-negative fluorine or oxygen. When other atoms are bound (such as hydrogen or carbon), they are often part of a molecule containing fluorine or oxygen.[64] Some compounds of xenon are colored but most are colorless.[62]

In 1995, a group of scientists at the University of Helsinki in Finland (M. Räsänen and co-workers) announced the preparation of xenon dihydride (HXeH), and later xenon hydride-hydroxide (HXeOH), hydroxenoacetylene (HXeCCH), and other Xe-containing molecules. [65][66] Deuterated molecules, HXeOD and DXeOH, have also been produced.[67]

XeF4 crystals. 1962.

As well as compounds where xenon forms a chemical bond, xenon can form clathrates--substances where xenon atoms are trapped by the crystalline lattice of another compound. An example is xenon hydrate (Xe*5.75 H2O), where xenon atoms occupy vacancies in a lattice of water molecules.[68] The deuterated version of this hydrate has also been produced.[69] Such clathrate hydrates can occur naturally under conditions of high pressure, such as in Lake Vostok underneath the Antarctic ice sheet.[70] Clathrate formation can be used to fractionally distill xenon, argon and krypton.[71] 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 monitored via 129Xe nuclear magnetic resonance spectroscopy. Using this technique, chemical reactions on the fullerene molecule can be analyzed, due to the sensitivity of the chemical shift of the xenon atom to its environment. However, the xenon atom also has an electronic influence on the reactivity of the fullerene.[72]

## Applications

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Although xenon is rare and relatively expensive to extract from the Earth's atmosphere, it still has a number of applications.

### Illumination and optics

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#### Gas-discharge lamps

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Xenon is used in light-emitting devices called xenon flash lamps, which are used in photographic flashes and stroboscopic lamps;[7] to excite the active medium in lasers which then generate coherent light;[73] and, occasionally, in bactericidal lamps.[74] The first solid-state laser, invented in 1960, was pumped by a xenon flash lamp,[11] and lasers used to power inertial confinement fusion are also pumped by xenon flash lamps.[75]

Continuous, short-arc, high pressure xenon arc lamps have a color temperature closely approximating noon sunlight and are used in solar simulators. That is, the chromaticity of these lamps closely approximates a heated black-body radiator that has a temperature close to that observed from the Sun. After they were first introduced during the 1940s, these lamps began replacing the shorter-lived carbon arc lamps in movie projectors.[8] They are employed in typical 35mm and IMAX film projection systems, automotive HID headlights and other specialized uses. These arc lamps are an excellent source of short wavelength ultraviolet radiation and they have intense emissions in the near infrared, which is used in some night vision systems.

The individual cells in a plasma display use a mixture of xenon and neon that is converted into a plasma using electrodes. The interaction of this plasma with the electrodes generates ultraviolet photons, which then excite the phosphor coating on the front of the display.[76][77]

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.[78]

#### Lasers

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In 1962, a group of researchers at Bell Laboratories discovered laser action in xenon,[79] and later found that the laser gain was improved by adding helium to the lasing medium.[80][81] The first excimer laser used a xenon dimer (Xe2) energized by a beam of electrons to produce stimulated emission at an ultraviolet wavelength of 176 nm.[10] Xenon chloride and xenon fluoride have also been used in excimer (or, more accurately, exciplex) lasers.[82] The xenon chloride excimer laser has been employed, for example, in certain dermatological uses.[83]

### Anesthesia

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Xenon has been used as a general anesthetic, although it is expensive. Even so, anesthesia machines that can deliver xenon are about to appear on the European market.[84] Two mechanisms for xenon anesthesia have been proposed. The first one involves the inhibition of the calcium ATPase pump--the mechanism cells use to remove calcium (Ca2+)--in the cell membrane of synapses.[85] This results from a conformational change when xenon binds to nonpolar sites inside the protein.[86] The second mechanism focuses on the non-specific interactions between the anesthetic and the lipid membrane.[87]

Xenon has a minimum alveolar concentration (MAC) of 0.63, making it 50% more potent than N2O as an anesthetic. Thus it can be used in concentrations with oxygen that have a lower risk of hypoxia. Unlike nitrous oxide (N2O), xenon is not a greenhouse gas and so it is also viewed as environmentally friendly. Because of the high cost of xenon, however, economic application will require a closed system so that the gas can be recycled, with the gas being appropriately filtered for contaminants between uses.[37]

### Medical imaging

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Gamma emission from the radioisotope 133Xe of xenon can be used to image the heart, lungs, and brain, for example, by means of single photon emission computed tomography. 133Xe has also been used to measure blood flow.[88][89][90]

Nuclei of only two of the stable isotopes of xenon, 129Xe and 131Xe, have non-zero intrinsic angular momenta (nuclear spins). When mixed with alkali vapor and nitrogen, their nuclear spins can be aligned along the laser beam of circularly-polarized light that is tuned to an absorption line of the alkali atoms. Typically, pure rubidium metal, heated above 100 °C, is used to produce the alkali vapor. This alignment (spin polarization) of xenon nuclei can surpass 50% of its maximum possible value, greatly exceeding the equilibrium value dictated by the Boltzmann distribution (typically 0.001% of the maximum value at room temperature, even in the strongest magnets). Such non-equilibrium alignment of spins is a temporary condition, and is called hyperpolarization. Because the 129Xe isotope has a nuclear spin value of 1/2 (and therefore the electric quadrupole moment of 129Xe nucleus must be zero), 129Xe nucleus does not experience any quadrupolar interactions during collisions with other atoms, and thus its hyperpolarization can be maintained for long periods of time even after the laser beam has been turned off and the alcali vapor removed by condensation on a room-temperature surface. The time it takes for a collection of spins to return to their equilibrium (Boltzmann) polarization is called the T1 relaxation time. For 129Xe isotope it can range from several seconds for xenon atoms dissolved in blood[91] to several hours in the gas phase[92] and to several days in the deeply-frozen solid xenon.[93] In contrast, the 131Xe isotope has a nuclear spin value of 3/2, does possess a non-zero quadrupole moment, and has T1 relaxation times in the millisecond and second ranges.[94] The hyperpolarization process (such as Spin-Exchange optical pumping described above) renders the 129Xe isotope much more detectable via magnetic resonance imaging and has been used for studies of the lungs and other tissues. It can be used, for example, to trace the flow of gases within the lungs.[95][96]

### Other

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In nuclear energy applications, xenon is used in bubble chambers,[97] probes, and in other areas where a high molecular weight and inert nature is desirable. Liquid xenon is being used as a medium for detecting hypothetical weakly interactive massive particles, or WIMPs. When a WIMP collides with a xenon nucleus, it should, theoretically, strip an electron and create a primary scintillation. By using xenon, this burst of energy could then be readily distinguished from similar events caused by particles such as cosmic rays.[12] However, the XENON experiment at the Gran Sasso National Laboratory in Italy has thus far failed to find any confirmed WIMPs. Even if no WIMPs are detected though, the experiment will serve to constrain the properties of dark matter and some physics models.[98] The current detector at this facility is five times as sensitive as other instruments world-wide, and the sensitivity will be increased by an order of magnitude in 2008.[99]

Xenon is the preferred fuel for ion propulsion of spacecraft because of its low ionization potential per atomic weight, the ability to store it as a liquid at near room temperature (but at high pressure) yet easily converts back into a gas to fuel the engine. The inert nature of xenon makes it environmentally friendly and less corrosive to an ion engine than other fuels such as mercury or caesium. Xenon was first used for satellite ion engines during the 1970s.[100] It was later employed as a propellant for Europe's SMART-1 spacecraft[13] and for the three ion propulsion engines on NASA's Dawn Spacecraft.[101]

Chemically, the perxenate compounds are used as oxidizing agents in analytical chemistry. Xenon difluoride is used as an etchant for silicon, particularly in the production of microelectromechanical systems (MEMS).[102] The anticancer drug 5-fluorouracil can be produced by reacting Xenon difluoride with Uracil.[103] Xenon is also used in protein crystallography. Applied at pressures from 0.5 to 5 MPa (5 to 50 atm) to a protein crystal, xenon atoms bind in predominantly hydrophobic cavities, often creating a high quality, isomorphous, heavy-atom derivative, which can be used for solving the phase problem.[104][105]

## Precautions

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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 anaesthesia when inhaled in very high concentrations (see anesthesia subsection above). Many of xenon compounds are explosive and toxic due to their strong oxidative properties.[106]

At 169 m/s, the speed of sound in xenon gas is slower than that in air[107] (due to the slower average speed of the heavy xenon atoms compared to nitrogen and oxygen molecules), so xenon lowers the resonant frequencies of the vocal tract when inhaled. This produces a characteristic lowered voice pitch, opposite the high-pitched voice caused by inhalation of helium. Like helium, xenon does not satisfy the body's need for oxygen and is a simple asphyxiant; consequently, many universities no longer allow the voice stunt as a general chemistry demonstration. As xenon is expensive, the gas sulfur hexafluoride, which is similar to xenon in molecular weight (146 versus 131), is generally used in this stunt, although it too is an asphyxiant.[108]

It is possible to safely breathe heavy gases such as xenon or sulfur hexafluoride when they include a 20% mixture of oxygen. The lungs mix the gases very effectively and rapidly, so that the heavy gases are purged along with the oxygen and do not accumulate at the bottom of the lungs.[109] There is, however, a danger associated with any heavy gas in large quantities: it may sit invisibly in a container, and if a person enters a container filled with an odorless, colorless gas, they may find themselves breathing it unknowingly. 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.[110]

## References

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1. ↑ "xenon", Columbia Electronic Encyclopedia, 6th ed., Columbia University Press, 2007. Accessed on line October 23, 2007.
2. ↑ 2.0 2.1 Husted, Robert; Boorman, Mollie (December 15, 2003). "Xenon". Los Alamos National Laboratory, Chemical Division. Retrieved 2007-09-26.CS1 maint: Multiple names: authors list (link)
3. ↑ Rabinovich, Viktor Abramovich (1988). Thermophysical properties of neon, argon, krypton, and xenon (English-language edition ed.). Washington, DC: Hemisphere Publishing Corp. ISBN 0195218337. Unknown parameter `|coauthors=` ignored (help)CS1 maint: Extra text (link) --National Standard Reference Data Service of the USSR. Volume 10.
4. ↑ 4.0 4.1 Freemantel, Michael (August 25, 2003). "Chemistry at its Most Beautiful" (PDF). Chemical & Engineering News. Retrieved 2007-09-13.
5. ↑ 5.0 5.1 Kaneoka, Ichiro (1998). "Xenon's Inside Story". Science. 280 (5365): 851–852. Retrieved 2007-10-10.
6. ↑ 6.0 6.1 Stacey, Weston M. (2007). Nuclear Reactor Physics. Wiley-VCH. pp. p. 213. ISBN 3527406794.CS1 maint: Extra text (link)
7. ↑ 7.0 7.1 7.2 Burke, James (2003). Twin Tracks: The Unexpected Origins of the Modern World. Oxford University Press. p. 33\. ISBN 0743226194.
8. ↑ 8.0 8.1 Mellor, David (2000). Sound Person's Guide to Video. Focal Press. pp. p. 186. ISBN 0240515951.CS1 maint: Extra text (link)
9. ↑ Sanders, Robert D.; Ma, Daqing; Maze, Mervyn (2005). "Xenon: elemental anaesthesia in clinical practice". British Medical Bulletin. 71 (1): 115–135. Retrieved 2007-10-02.CS1 maint: Multiple names: authors list (link)
10. ↑ 10.0 10.1 Basov, N. G.; Danilychev, V. A.; Popov, Yu. M. (1971). "Stimulated Emission in the Vacuum Ultraviolet Region". Soviet Journal of Quantum Electronics. 1 (1): 18–22. doi:10.1070/QE1971v001n01ABEH003011.CS1 maint: Multiple names: authors list (link)
11. ↑ 11.0 11.1 Toyserkani, E.; Khajepour, A.; Corbin, S. (2004). Laser Cladding. CRC Press. p. 48\. ISBN 0849321727.CS1 maint: Multiple names: authors list (link)
12. ↑ 12.0 12.1 Ball, Philip (May 1, 2002). "Xenon outs WIMPs". Nature. Retrieved 2007-10-08.
13. ↑ 13.0 13.1 Saccoccia, G.; del Amo, J. G.; Estublier, D. (August 31, 2006). "Ion engine gets SMART-1 to the Moon". ESA. Retrieved 2007-10-01.CS1 maint: Multiple names: authors list (link)
14. ↑ W. Ramsay and M. W. Travers (1898). "On the extraction from air of the companions of argon, and neon". Report of the Meeting of the British Association for the Advancement of Science: 828.
15. ↑ Gagnon, Steve. "It's Elemental - Xenon". Thomas Jefferson National Accelerator Facility. Retrieved 2007-06-16.
16. ↑ Anonymous (1904). Daniel Coit Gilman, Harry Thurston Peck, Frank Moore Colby, ed. The New International Encyclopædia. Dodd, Mead and Company. pp. p. 906.CS1 maint: Multiple names: editors list (link) CS1 maint: Extra text (link)
17. ↑ Staff (1991). The Merriam-Webster New Book of Word Histories. Merriam-Webster, Inc. pp. p. 513. ISBN 0877796033.CS1 maint: Extra text (link)
18. ↑ Ramsay, William (1902). "An Attempt to Estimate the Relative Amounts of Krypton and of Xenon in Atmospheric Air". Proceedings of the Royal Society of London. 71: 421–426. Retrieved 2007-10-02.
19. ↑ Anonymous. "History". Millisecond Cinematography. Retrieved 2007-11-07.
20. ↑ Paschotta, Rüdiger (November 1, 2007). "Lamp-pumped lasers". Encyclopedia of Laser Physics and Technology. RP Photonics. Retrieved 2007-11-07.
21. ↑ Marx, Thomas; Schmidt, Michael; Schirmer, Uwe; Reinelt, Helmut (2000). "Xenon anesthesia" (PDF). Journal of the Royal Society of Medicine. 93: 513–517. Retrieved 2007-10-02.CS1 maint: Multiple names: authors list (link)
22. ↑ Clayton, Donald D. (1983). Principles of Stellar Evolution and Nucleosynthesis (2nd edition ed.). University of Chicago Press. pp. p. 75. ISBN 0226109534.CS1 maint: Extra text (link)
23. ↑ Bolt, B. A.; Packard, R. E.; Price, P. B. (2007). "John H. Reynolds, Physics: Berkeley". The University of California, Berkeley. Retrieved 2007-10-01.CS1 maint: Multiple names: authors list (link)
24. ↑ Neil Bartlett and D. H. Lohmann (1962). "Dioxygenyl hexafluoroplatinate (V), O2+[PtF6]−". Proceedings of the Chemical Society. London: Chemical Society (3): 115. doi:10.1039/PS9620000097. Unknown parameter `|month=` ignored (help)
25. ↑ 25.0 25.1 Bartlett, N. (1962). "Xenon hexafluoroplatinate (V) Xe+[PtF6]−". Proceedings of the Chemical Society. London: Chemical Society (6): 218. doi:10.1039/PS9620000197. Unknown parameter `|month=` ignored (help)
26. ↑ Graham, L. (2000). "Concerning the nature of XePtF6". Coordination Chemistry Reviews. 197: 321–334. doi:10.1016/S0010-8545(99)00190-3. Unknown parameter `|coauthors=` ignored (help)
27. ↑ p. 392, §11.4, Inorganic Chemistry, translated by Mary Eagleson and William Brewer, edited by Bernhard J. Aylett, San Diego: Academic Press, 2001, ISBN 0-12-352651-5; translation of Lehrbuch der Anorganischen Chemie, originally founded by A. F. Holleman, continued by Egon Wiberg, edited by Nils Wiberg, Berlin: de Gruyter, 1995, 34th edition, ISBN 3-11-012641-9.
28. ↑ Steel, Joanna (2007). "Biography of Neil Bartlett". College of Chemistry, University of California, Berkeley. Retrieved 2007-10-25.
29. ↑ Bartlett, Neil (September 8, 2003). "The Noble Gases". Chemical & Engineering News. American Chemical Society. 81 (36). Retrieved 2007-10-01.
30. ↑ Leonid Khriachtchev, Mika Pettersson, Nino Runeberg, Jan Lundell, and Markku Räsänen (August 24, 2000). "A stable argon compound". Nature. 406: 874&ndash, 876. doi:10.1038/35022551.CS1 maint: Multiple names: authors list (link)
31. ↑ Lynch, C. T.; Summitt, R.; Sliker, A. (1980). CRC Handbook of Materials Science. CRC Press. ISBN 087819231X.CS1 maint: Multiple names: authors list (link)
32. ↑ D. R. MacKenzie (September 20, 1963). "Krypton Difluoride: Preparation and Handling". Science. 141 (3586): 1171. doi:10.1126/science.141.3586.1171.
33. ↑ Paul R. Fields, Lawrence Stein, and Moshe H. Zirin (1962). "Radon Fluoride". Journal of the American Chemical Society. 84 (21): 4164&ndash, 4165. doi:10.1021/ja00880a048.CS1 maint: Multiple names: authors list (link)
34. ↑ Hwang, Shuen-Cheng (2005). "Noble Gases". Kirk-Othmer Encyclopedia of Chemical Technology (5th edition ed.). Wiley. doi:10.1002/0471238961.0701190508230114.a01. ISBN 047148511X. Unknown parameter `|coauthors=` ignored (help)CS1 maint: Extra text (link)
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## External links

[edit | edit source]

* WebElements.com - Xenon
* USGS Periodic Table - Xenon
* EnvironmentalChemistry.com - Xenon
* Xenon as an anesthetic

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