Xenon in Periodic Table
Xenon (Xe), chemical element, inert gas, or noble gas of Group 18 of the periodic table occur in trace amounts in the earth’s atmosphere, used in flash or arc lamps, and an anesthetic in medicine. It is a colorless, dense, odorless, tasteless, monoatomic gas that forms true chemical compounds with fluorine.
A heavy and extremely rare gas, xenon was discovered by the British chemists Sir William Ramsay and Morris W. Travers and name the gas derived from the Greek latter xenos meaning strange or foreign. In solid-state, xenon forms a face-centered cubic crystal lattice, and the cubic close-packed structure being the most space-economizing.
The rare gas, xenon has the chemical symbol Xe, atomic number 54, atomic weight 131.30, melting point -111.75 °C, boiling point -108.01 °C, valence shell electron configuration [Kr] 4d10 5s2 4p6, and heat capacity ratio (Cp/Cv) close to 1.66. Due to filled valence orbital and high ionization energy, the oxidation number or state is zero but with fluorine, xenon shows a higher oxidation state also. It occurs in the earth’s atmosphere to the extent of 0.00001 percent by volume and is manufactured on a small scale by the fractional distillation of liquid air.
Isotopes and Uses
Natural xenon is a mixture of nine stable isotopes like 124Xe, 126Xe, 128Xe, 129Xe, 130Xe, 131Xe, 132Xe, 134Xe, and 136Xe. Different types of artificial radioactive isotopes are obtained by the nuclear fission reaction whose life is very small. Xenon-135 is produced in the reactor of the nuclear power plant by nuclear transmutation of uranium. Xenon-135 has a huge cross-section for thermal neutrons since it can be considered as a neutron absorber in a nuclear reaction or chain reaction. A huge number of radioactive isotopes of xenon are also released from the nuclear reactor during the nuclear fission reaction. Xenon isotopes are powerful tools for studying solar renewable energy systems and nuclear magnetic resonance spectroscopy.
Xenon is rare in the earth’s atmosphere and expensive than the other noble gases like helium, neon, argon, and krypton but the gas widely used in light-emitting devices like flash lamps, photographic flashes, high-pressure arc lamps, shorter-lived carbon arc lamps, and as a starter gas in high-pressure sodium lamps.
Chemistry of Xenon
The chemistry of xenon and other noble gases like helium, neon, argon, krypton has different from other periodic table elements due to their stable valence shell electron configuration. They cannot participate in covalent bonding due to very high electron promotional energy. Formation of chemical bonding to form an ionic compound like Xe+F– is similarly difficult due to very high ionization energy. Xe2+ ion is stable in a vacuum with a bond order of 0.5 and usually high bond energy. Prior to 1962, no true chemical compound of xenon is known. Clathrates or cage compounds formed with para-quinol and water.
As early as 1933, Linus Pauling attempt to made the chemical compound XeF6 but failed until 1962. In 1962, British chemist Neil Bartlett produced the first yellow-orange solid xenon compound that contains a mixture of [XeF+][PtF6−], [XeF+][Pt2F11−]. He observed that PtF6 changes color on exposure to air. It was subsequently established that PtF6 is a very strong oxidizing agent capable of oxidizing oxygen and xenon. This was the beginning of noble gas chemistry. Very soon different types of xenon compounds, including halides, oxides, oxofluorides, oxo salts, and numerous covalent derivatives are discovered.
Xenon fluorides like XeF2, XeF4, and XeF6 are the thermodynamically stable compounds obtained by heating xenon and fluorine in a sealed nickel vessel. Halides such as XeCl2, XeClF, XeBr2, and XeCl4 are thermodynamically unstable and have been detected only in small amounts. All three xenon fluorides are colorless volatile crystalline solids but in the liquid or gas phase XeF6 shows yellow color. All the fluorides are readily hydrolyzed to form Xe or XeO3. The formation of XeO3 is highly explosive which caused many accidents during the study of XeF4 and XeF6.
All three fluorides are strong oxidizing and fluorinating agents, reactivity increasing from XeF2 to XeF6. Therefore, XeF6 combines with hydrogen and attacks glass at ordinary temperature but others attack by heating. XeF2 used as a fluorinating agent in alky or aryl hydrocarbon synthesis and oxidizer of silver (I) to silver (II) and chromium (III) to chromium (IV) and BrO3– to BrO4–. Xenon fluorides react with fluoride ion acceptors like pentafluoride (MF5), where M = As, Sb, Bi, Nb, Ta, Ru, Os, Ir, Pt. Here halides act as a base to strong Lewis acid like MF5.
XeOF4 (melting point -46 °C) and XeO2F2 (melting point 31 °C) are the colorless crystalline oxyfluoride of xenon. The square pyramidal, XeOF4 is formed in the controlled hydrolysis of XeF6. The oxyfluoride XeO2F2 is the same structural formula, formed by the reaction of XeO3 with XeOF4 or XeF6. Other oxofluorides of xenon like XeOF2 and XeOF3 are unstable formed in a low-temperature matrix by reacting xenon with OF2 or by hydrolysis of XeF4 at low temperature. XeOF2 disproportionates in presence of CsF to give anion [XeO2F3]–. Stable white salts of [XeO3F]– are formed by reacting aqueous XeO3 with CsF or KF.
Xenon Oxides, Xenates, and Perxenates
The xenon trioxide, XeO3 is the end product of hydrolysis of XeF6, and the violence of the reaction controlled by sweeping XeF6 vapor with dry nitrogen into the water. XeO3 crystals consist of trigonal pyramidal units with xenon atom at the apex. The tetraoxide, XeO4 is formed by the action of concentrated sulfuric acid on sodium and barium perxenates. A perxenate of lithium, potassium are formed by disproportionation of XeO3 and XeF6 in the respective alkaline solution.
The dichloride, XeCl2 is not formed under usual conditions but may be formed when xenon and chlorine mixture passed through a microwave discharge and rapidly cool. XeCl2, XeBr2, and XeCl4 have also been detected by Mossbauer spectroscopy in the beta-particle decay. A chemical compound (BXeF3) containing boron xenon bond is formed by the reaction of O2BF4 and xenon at -100 °C contains a planner structure in which one fluorine atom in XeF2 is replaced by BF2 unit.