54
  Xe  
131.290000
Xenon

Name: Xenon
Symbol: Xe
Atomic Number: 54
Atomic Weight: 131.290000
Family: Noble Gases
CAS RN: 7440-63-3
Description: A colorless inert gas
State (25C): Gas
Oxidation states: 0

Molar Volume: 42.9 cm3/mole
Valence Electrons: 5p6

Boiling Point:  165.18K, -107.97C, -162F
Melting Point:
161.45K, -111.7C, -169.1F
Electrons Energy Level: 2, 8, 18, 18, 8
Isotopes: 30 + 8 Stable
Heat of Vaporization: 12.636 kJ/mol
Heat of Fusion: 2.297 kJ/mol
Density: 5.9 g/L @ 273K & 1atm
Specific Heat: 0.158J/gK
Atomic Radius: 1.24
Ionic Radius: unknown
Electronegativity: 0 (Pauling); 2.4 (Allrod Rochow)
Xenon (from Greek, meaning "strange") is the rarest of the stable noble gases in the air.  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.  It is still recovered by liquefaction techniques and is widely used in strobe lamps.

Xenon is a member of the zero-valence elements that are called noble or inert gases, however, "inert" is not a completely accurate description of this chemical series since some noble gas compounds have been synthesized.  In 1962 the first noble gas compound was produced by Neil Bartlett, combining xenon, platinum and fluorine.  It is now possible to produce xenon compounds in which the oxidation states range from +2 to +8.

In a gas filled tube, xenon  emits a blue glow when the gas is excited by electrical discharge. Using tens of gigapascals of pressure, xenon has been forced into a metallic phase.  Xenon can also form clathrates with water when atoms of it are trapped in a lattice of the water molecules.

Like the noble gas krypton, xenon can also be extracted by fractional distillation or liquefaction of liquid air and by selective adsorption on activated carbon.

Xenon is a trace gas in Earth's atmosphere, occurring in one part in twenty million.   The element is obtained commercially through extraction from the residues of liquefied air.  This noble gas is naturally found in gases emitted from some mineral springs.  133Xe and 135Xe are synthesized by neutron irradiation within air-cooled nuclear reactors.

2
He
4.002
10
Ne
20.17
18
Ar
39.94
36
Kr
83.80
54
Xe
131.3
86
Rn
222.0
118
Uuo
293.0

1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p6

Applications

This gas is most widely and most famously used in light-emitting devices called Xenon flash lamps, which are used in photographic flashes and stroboscopic lamps, to excite the active medium in lasers which then generate coherent light, to produce laser power for inertial confinement fusion, in bactericidal lamps (rarely), and in certain dermatological uses. Continuous, short-arc, high pressure Xenon arc lamps have a color temperature closely approximating noon sunlight and are used in solar simulators, some projection systems, automotive HID headlights and other specialized uses. They are an excellent source of short wavelength ultraviolet radiation and they have intense emissions in the near infrared, which are used in some night vision systems. Other uses of Xenon:

1s2
2s2 2p6
3s2 3p6 3d10
4s2 4p6 4d10
5s2 5p6

Compounds

Xenon and the other noble gases had for a long time been considered to be completely chemically inert and not able to form compounds.  However, in 1962 at the University of British Columbia, the first xenon compound, Xenon Hexafluoroplatinate, was synthesized by Neil Bartlett.  Now, many compounds of xenon are known, including xenon difluoride, xenon tetrafluoride, xenon hexafluoride, xenon tetroxide, xenon hydrate, xenon deuterate, and sodium perxenate.  A highly explosive compound xenon trioxide has also been made.  There are at least 80 xenon compounds in which fluorine or oxygen is bonded to xenon. Some compounds of xenon are colored but most are colorless.

Recently at the University of Helsinki in Finland, a group of scientists (M. Rsnen et al.) prepared HXeH, HXeOH, and HXeCCH (xenon dihydride, xenon hydride-hydroxide, and hydroxenoacetylene).  They are stable up to 40oK.

Isotopes

Naturally occurring xenon is made of eight stable and two slightly radioactive isotopes.  Beyond these stable forms, there are 30 unstable isotopes that have been studied. 129Xe is produced by beta decay of 129I (half-life: 16 million years); 131mXe, 133Xe, 133mXe, and 135Xe are some of the fission products of both 235U and 239Pu, and therefore used as indicators of nuclear explosions.

The artificial isotope 135Xe is of considerable significance in the operation of nuclear fission reactors.  135Xe has a huge cross section for thermal neutrons, 2.65x106 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).

Relatively high concentrations of radioactive xenon isotopes are also found emanating from nuclear reactors due to the release of this fission gas from cracked fuel rods or fissioning of uranium in cooling water. The concentrations of these isotopes are still usually low compared to naturally occurring radioactive noble gases such as 222Rn.

Because xenon is a tracer for two parent isotopes, Xe isotope ratios in meteorites are a powerful tool for studying the formation of the solar system.  The I-Xe method of dating gives the time elapsed between nucleosynthesis and the condensation of a solid object from the solar nebula.  Xenon isotopes are also a powerful tool for understanding terrestrial differentiation. 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.

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Isotope Atomic Mass Half-Life
Xe110 109.944 0.6 us
Xe111 110.942 0.74 seconds
Xe112 111.936 2.7 seconds
Xe113 112.9334 2.74 seconds
Xe114 113.928 10 seconds
Xe115 114.927 18 seconds
Xe116 115.922 59 seconds
Xe117 116.921 61 seconds
Xe118 117.92 3.8 minutes
Xe119 118.916 5.8 minutes
Xe120 119.9122 40 minutes
Xe121 120.9114 40.1 minutes
Xe122 121.9085 20.1 hours
Xe123 122.9085 2.08 hours
Xe124 123.9059 Stable
Xe125 124.9064 16.9 hours
Xe126 125.9043 Stable
Xe127 126.9052 36.4 days
Xe128 127.9035 Stable
Xe129 128.9048 Stable
Xe130 129.9035 Stable
Xe131 130.9051 Stable
Xe132 131.9042 Stable
Xe133 132.9059 5.243 days
Xe134 133.9054 Stable
Xe135 134.9072 9.14 hours
Xe136 135.9072 >2.36E 21 years
Xe137 136.9116 3.818 minutes
Xe138 137.914 14.08 minutes
Xe139 138.9188 39.68 seconds
Xe140 139.9216 13.6 seconds
Xe141 140.9266 1.73 seconds
Xe142 141.93 1.22 seconds
Xe143 142.935 0.3 seconds
Xe144 143.938 1.15 seconds
Xe145 144.944 0.9 seconds
Xe146 145.947 >150 ns
Xe147 146.953 >150 ns

Precautions

40px-Skull_and_crossbones.svg.jpg (1420 bytes) The gas can be safely kept in normal sealed glass containers at standard temperature and pressure.  Xenon is non-toxic, but many of its compounds are toxic due to their strong oxidative properties.

The speed of sound in xenon is slower than that in air (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 vs 131), is generally used in this stunt, although it too is an asphyxiant.

A myth exists that xenon is too heavy for the lungs to expel unassisted, and that after inhaling xenon, it is necessary to bend over completely at the waist to allow the excess gas to "spill" out of the body.  In fact, the lungs mix gases very effectively and rapidly, such that xenon would be purged from the lungs within a breath or two. 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.


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Xenon Data

 

Atomic Structure

Atomic Radius (): 1.24
Atomic Volume cm3/mol : 37.3cm3/mol
Covalent Radius: 1.31
Crystal Structure: Cubic face centered
Ionic Radius: unknown

Chemical Properties

Electrochemical Equivalents: unknown
Electron Work Function: unknown
Electronegativity: 0 (Pauling); 2.4 (Allrod Rochow)
Heat of Fusion: 2.297 kJ/mol
Incompatibilities: unknown
First Ionization Potential: 12.13
Second Ionization Potential: 21.21
Third Ionization Potential: 32.1
Valence Electron Potential: unknown
Ionization Energy (eV): 12.130 eV

Physical Properties

Atomic Mass Average: 131.29
Boiling Point: 165.18K, -107.97C, -162F
Melting Point: 161.45K, -111.7C, -169.1F
Heat of Vaporization: 12.636 kJ/mol
Coefficient of Lineal Thermal Expansion/K-1: N/A
Electrical Conductivity: unknown
Thermal Conductivity: 0.0000569 W/cmK
Density: 5.9 g/L @ 273K & 1atm
Enthalpy of Atomization: unknown
Enthalpy of Fusion: 2.3 kJ/mole
Enthalpy of Vaporization: 12.64 kJ/mole
Flammability Class: unknown
Molar Volume: 42.9 cm3/mole
Optical Refractive Index: 1.000702
Relative Gas Density (Air=1): unknown
Specific Heat: 0.158J/gK
Vapor Pressure: unknown
Estimated Crustal Abundance: 310-5 milligrams per kilogram
Estimated Oceanic Abundance: 510-5 milligrams per liter


(Gr. xenon, stranger) Discovered by Ramsay and Travers in 1898 in the residue left after evaporating liquid air components. Xenon is a member of the so-called noble or "inert" gases.  It is present in the atmosphere to the extent of about one part in twenty million.  Xenon is present in the Martian atmosphere to the extent of 0.08 ppm. the element is found in the gases evolved from certain mineral springs, and is commercially obtained by extraction from liquid air.  Natural xenon is composed of eight stable isotopes.  In addition to these, 30 unstable isotopes have been characterized.  Before 1962, it had generally been assumed that xenon and other noble gases were unable to form compounds.  Evidence has been mounting in the past few years that xenon, as well as other members of zero valance elements, do form compounds.  Among the "compounds" of xenon now reported are sodium perxenate, xenon deuterate, xenon hydrate, difluoride, tetrafluoride, and hexafluoride.   Xenon trioxide, which is highly explosive, has been prepared.  More than 80 xenon compounds have been made with xenon chemically bonded to fluorine and oxygen.   Some xenon compounds are colored. Metallic xenon has been produced, using several hundred kilobars of pressure.  Xenon in a vacuum tube produces a beautiful blue glow when excited by an electrical discharge.  The gas is used in making electron tubes, stoboscopic lamps, bactericidal lamps, and lamps used to excite ruby lasers for generating coherent light.  Xenon is used in the nuclear energy field in bubble chambers, probes, and other applications where a high molecular weight is of value.  The perxenates are used in analytical chemistry as oxidizing agents. 133Xe and 135Xe are produced by neutron irradiation in air cooled nuclear reactors.  133Xe has useful applications as a radioisotope.  The element is available in sealed glass containers of gas at standard pressure.  Xenon is not toxic, but its compounds are highly toxic because of their strong oxidizing characteristics.

Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond