Xenon

What is Xenon?

Xenon is a chemical element or inert or noble gas of Group 18 of the periodic table with the symbol Xe and atomic number 54. It is found in trace amounts in the earth’s environment. Due to filled valence orbital and high ionization energy, the oxidation number or state is zero, but with fluorine, xenon also shows a higher oxidation state. Xenon is used in flash or arc lamps and as an anesthetic in medicine. It is a colorless, dense, odorless, tasteless, monoatomic gas that forms true chemical compounds with fluorine.

Xenon in the Periodic Table

Xenon is placed in group 18 and period 5 in the periodic table. It is a noble gas that lies between krypton and radon.

Who Discovered Xenon?

A heavy and extremely rare gas, xenon was discovered by the British chemists Sir William Ramsay and Morris W. Travers. The name of the gas was derived from the Greek letter Xenos meaning strange or foreign.

Where is Xenon Found?

It occurs in the earth’s atmosphere to the extent of 0.00001 percent by volume. Xenon is manufactured on a small scale by the fractional distillation of liquid air. Xenon is rare in the earth’s atmosphere and more expensive than other noble gases like helium, neon, argon, and krypton.

The gas is widely used in light-emitting devices like flash lamps, photographic flashes, high-pressure arc lamps, and shorter-lived carbon arc lamps. It is also used as a starter gas in high-pressure sodium lamps.

Properties of Xenon

The rare gas, xenon has valence shell electron configuration [Kr] 4d¹⁰ 5s² 4p⁶, and a heat capacity ratio (C_p/C_v) close to 1.66.

In solid-state, xenon forms a face-centered cubic crystal lattice, and the cubic close-packed structure is the most space-economizing.

Isotopes

Natural xenon is a mixture of nine stable isotopes. The atomic mass and natural abundance of these xenon isotopes are given below in the table,

Different types of artificial radioactive isotopes are obtained by the nuclear fission reaction whose life is very small.

Xe-135 is produced in the reactor of the nuclear power plant by nuclear transmutation of uranium. Xe-135 has a huge cross-section for thermal neutrons since it can be considered 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.

Uses of Xenon

Xenon is rare in the earth’s atmosphere and more expensive than other noble gases like helium, neon, argon, and krypton.

• The gas is used in specific light-emitting devices like flash lamps, photographic flashes, high-pressure arc lamps, and shorter-lived carbon arc lamps.
• It is a starter gas in high-pressure sodium lamps.
• Xenon isotopes are powerful tools for studying solar renewable energy systems and nuclear magnetic resonance spectroscopy.

Facts About Xenon

The chemistry of xenon and other noble gases like He, Ne, Kr, and Rn is 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.

The 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 are formed with para-quinol and water.

Chemical Compounds

As early as 1933, Linus Pauling attempted to make 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 Fluoride

Xenon fluorides like XeF₂, XeF₄, and XeF₆ are thermodynamically stable compounds. They are obtained by heating Xe and fluorine in a sealed nickel vessel. Halides such as XeCl₂, XeClF, XeBr₂, and XeCl₄ are thermodynamically unstable. Therefore, they 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 is used as a fluorinating agent in alky or aryl hydrocarbon synthesis. It also acts as an 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 = arsenic, antimony, bismuth, niobium, tantalum, ruthenium, osmium, iridium, and platinum. Here halides act as a base for strong Lewis acid like MF₅.

Xenon Oxyfluorides

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 and formed in a low-temperature matrix by reacting Xe with OF₂ or by hydrolysis of XeF₄ at low temperatures.

XeOF₂ disproportionates in the presence of CsF to give anion [XeO₂F₃]⁻. Stable white salt of [XeO₃F]⁻ is formed by reacting aqueous XeO₃ with CsF or KF.

Oxides, Xenates, and Perxenate

The xenon trioxide, XeO₃ is the end product of hydrolysis of XeF₆. The violence of the reaction is controlled by sweeping XeF₆ vapor with dry nitrogen into the water. XeO₃ crystals consist of trigonal pyramidal units with xenon atoms at the apex.

The tetraoxide, XeO₄ is formed by the action of concentrated sulfuric acid on sodium and barium perxenates. A perxenate of lithium and potassium is formed by the disproportionation of XeO₃ and XeF₆ in the respective alkaline solution.

The dichloride, XeCl₂ is not formed under usual conditions. However, it may be formed when the Xe and chlorine mixture passes through a microwave discharge and rapidly cools. XeCl₂, XeBr₂, and XeCl₄ have also been detected by Mossbauer spectroscopy in the beta-particle decay.

A chemical compound (BXeF₃) containing a boron xenon bond is formed by the reaction of O₂BF₄ and xenon at −100 °C. It contains a planner structure in which one fluorine atom in XeF₂ is replaced by a BF₂ unit.