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Xenon is a chemical element with the symbol Xe and atomic number 54. It is a colourless, heavy, odourless noble gas, that occurs in the Earth’s atmosphere in trace amounts. 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.

Naturally occurring xenon consists of eight stable isotopes. 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. Radioactive xenon-135 is produced from iodine-135 as a result of nuclear fission, and it acts as the most significant neutron absorber in nuclear reactors.

Xenon is used in flash lamps and arc lamps, and as a general aesthetic. The first excimer laser design used a xenon dimer molecule (Xe2) as its lasing medium, and the earliest laser designs used xenon flash lamps as pumps. Xenon is also being used to search for hypothetical weakly interacting massive particles and as the propellant for ion thrusters in spacecraft.


Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers on July 12, 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 [xenon], neuter singular form of ξένος [xenos], meaning ‘foreign(er)’, ‘strange(r)’, or ‘guest’. In 1902, Ramsay estimated the proportion of xenon in the Earth’s atmosphere as one part in 20 million. The current symbol for Xenon is Xe, however historically it was also written as X.

During the 1930s, American 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.

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 aesthetic. Although Russian toxicologist Nikolay V. Lazarev apparently studied xenon anaesthesia in 1941, the first published report confirming xenon anaesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice. Xenon was first used as a surgical aesthetic. in 1951 by American anaesthesiologist Stuart C. Cullen, who successfully operated on two patients.

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 oxidise oxygen gas (O2) to form dioxygenyl hexafluoroplatinate (O2+[PtF6]–). Since O2 and xenon have almost the same first ionization potential, Bartlett realized that platinum hexafluoride might also be able to oxidise xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate. Bartlett thought its composition to be Xe+[PtF6]–, although later work has revealed that it was probably a mixture of various xenon-containing salts. Since then, many other xenon compounds have been discovered, along with some compounds of the noble gases argon, krypton, and radon, including argon fluorohydride (HArF), krypton difluoride (KrF2), and radon fluoride. By 1971, more than 80 xenon compounds were known.


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 surface density of the Earth’s atmosphere, 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. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is higher than the average density of granite, 2.75 g/cm3. Using gigapascals of pressure, xenon has been forced into 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 metalized, xenon looks sky blue because it absorbs red light and transmits other visible frequencies. Such behaviour is unusual for a metal and is explained by the relatively small widths of the electron bands in metallic xenon.

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. 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, but the most intense lines occur in the region of blue light, which produces the colouration.

Occurrence and production

Xenon is a trace gas in Earth’s atmosphere, occurring at 87±1 parts per billion (nL/L), or approximately 1 part per 11.5 million, and is also found in gases emitted from some mineral springs.

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 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. Extraction of a litre of xenon from the atmosphere requires 220 watt-hours of energy.

Worldwide production of xenon in 1998 was estimated at 5,000–7,000 m3. Because of 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; the much more plentiful argon costs less than a cent per litre.

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 planet Jupiter has an unusually high abundance of xenon in its atmosphere; about 2.6 times as much as the Sun. This high abundance remains unexplained and may have been caused by an early and rapid build-up of planetesimals—small, sub planetary bodies—before the presolar disk began to heat up. (Otherwise, xenon would not have been trapped in the planetesimal ices.) The problem of the low terrestrial xenon may potentially be explained by covalent bonding of xenon to oxygen within quartz, hence reducing the out gassing of xenon into the atmosphere.

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 when creating xenon. Instead, xenon is formed during supernova explosions, by the slow neutron capture process ( s-process) of red giant stars that have exhausted the hydrogen at their cores and entered the asymptotic giant branch, in classical nova explosions and from the radioactive decay of elements such as iodine, uranium and plutonium.

Isotopes and isotopic studies

Naturally occurring xenon is made of eight stable isotopes, the most of any element with the exception of tin, which has ten. Xenon and tin are the only elements to have more than seven stable isotopes. The isotopes 124Xe and 134Xe are predicted to undergo double beta decay, but this has never been observed so they are considered to be stable. Besides these stable forms, there are over 40 unstable isotopes that have been studied. The longest lived of these isotopes is 136Xe, which has been observed to undergo double beta decay with a half-life of 2.11 x 1021yr. 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, and therefore used as indicators of nuclear explosions.

Nuclei of two of the stable isotopes of xenon, 129Xe and 131Xe, have non-zero intrinsic angular momenta ( nuclear spins, suitable for nuclear magnetic resonance). The nuclear spins can be aligned beyond ordinary polarization levels by means of circularly polarized light and rubidium vapour. The resulting 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 hyper polarization. The process of hyper polarizing the xenon is called optical pumping (although the process is different from pumping a laser).

Because a 129Xe nucleus has a spin of 1/2, and therefore a zero electric quadrupole moment, the 129Xe nucleus does not experience any quadrupolar interactions during collisions with other atoms, and thus its hyper polarization. can be maintained for long periods of time even after the laser beam has been turned off and the alkali vapour. removed by condensation on a room-temperature surface. 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 non-zero 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, 2.6×106 barns, 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). 135Xe reactor poisoning played a major role in the Chernobyl disaster. A shut-down or decrease of power of a reactor can result in build-up of 135Xe and getting the reactor into the iodine pit.

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, or fissioning of uranium in cooling water.

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. In 1960, physicist John H. Reynolds discovered that certain meteorites contained an isotopic anomaly in the form of an over-abundance of xenon-129. He inferred that this was a decay product of radioactive iodine-129. This isotope is produced slowly by cosmic ray spallation and nuclear fission, but is produced in quantity only in supernova explosions. As the half-life of 129I is comparatively short on a cosmological time scale, only 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, as the 129I isotope was likely generated before the Solar System was formed, but not long before, and seeded 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 also a powerful tool for understanding planetary differentiation and early out gassing. For example, The atmosphere of Mars shows a xenon abundance similar to that of Earth: 0.08 parts per million, 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. In another example, 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.


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.


Three fluorides are known: XeF2, XeF4, and XeF6. XeF is theorized to be unstable. The fluorides are the starting point for the synthesis of almost all xenon compounds.

The solid, crystalline difluoride XeF2 is formed when a mixture of fluorine and xenon gases is exposed to ultraviolet light. Ordinary daylight is sufficient. Long-term heating of XeF2 at high temperatures under an NiF2 catalyst yields XeF6. Pyrolysis of XeF6 in the presence of NaF yields high-purity XeF4.

The xenon fluorides behave as both fluoride acceptors and fluoride donors, forming salts that contain such cations as XeF+ and Xe


3, and anions such as XeF−5, XeF−7, and XeF2−8. The green, paramagnetic Xe+

2 is formed by the reduction of XeF2 by xenon gas.

XeF2 is also able to form coordination complexes with transition metal ions. Over 30 such complexes have been synthesized and characterized.

Whereas the xenon fluorides are well-characterized, the other halides are not known, the only exception being the dichloride, XeCl2. Xenon dichloride is reported to be an endothermic, colourless, crystalline compound that decomposes into the elements at 80°C, formed by the high-frequency irradiation of a mixture of xenon, fluorine, and silicon or carbon tetrachloride. However, doubt has been raised as to whether XeCl2 is a real compound and not merely a van der Waals molecule consisting of weakly bound Xe atoms and Cl2 molecules. Theoretical calculations indicate that the linear molecule XeCl2 is less stable than the van der Waals complex.

Oxides and oxohalides

Three oxides of xenon are known: xenon trioxide (XeO3) and xenon tetroxide (XeO4), both of which are dangerously explosive and powerful oxidizing agents, and xenon dioxide (XeO2), which was reported in 2011 with a coordination number of four. XeO2 forms when xenon tetrafluoride is poured over ice. Its crystal structure may allow it to replace silicon in silicate minerals. The XeOO+ cation has been identified by infra-red spectroscopy in solid argon.

Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF6: XeF6 + 3 H2O → XeO3 + 6 HF

XeO3 is weakly acidic, dissolving in alkali to form unstable xenate salts containing the HXeO−4 anion. These unstable salts easily disproportionate into xenon gas and perxenate salts, containing the XeO4−6 anion.

Barium perxenate, when treated with concentrated sulphuric acid, yields gaseous xenon tetroxide:

Ba2XeO6 + 2 H2SO4 → 2 BaSO4 + 2 H2O + XeO4

To prevent decomposition, the xenon tetroxide thus formed is quickly cooled to form a pale-yellow solid. It explodes above −35.9 °C into xenon and oxygen gas.

A number of xenon oxyfluorides are known, including XeOF2, XeOF4, XeO2F2, and XeO3F2. XeOF2 is formed by the reaction of OF2 with xenon gas at low temperatures. It may also be obtained by the partial hydrolysis of XeF4. It disproportionates at −20 °C into XeF2 and XeO2F2. XeOF4 is formed by the partial hydrolysis of XeF6, or the reaction of XeF6 with sodium perxenate, Na4XeO6. The latter reaction also produces a small amount of XeO3F2. XeOF4 reacts with CsF to form the XeOF−5 anion, while XeOF3 reacts with the alkali metal fluorides KF, RbF and CsF to form the XeOF−4 anion.

Synopsis of https://en.wikipedia.org/wiki/Xenon


Posted 2018/04/20 by Stelios in Education

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