Xenon Difluoride – Structure, Properties, Applications

Xenon Difluoride

There are a total of 18 groups in the periodic table. All the elements are divided into different blocks according to their properties. Most of the properties of all elements are dependent on their valence shell electrons. Hence, according to the number of valence electrons, some group elements are highly reactive while some are least.

Even group 18 elements are known for their non-reactive and exceptional behavior. They have complete octets that let them be the least chemically active. As a result, they do not easily form any compound or bond with other atoms.

Xenon is one among these group 18 inert gasses. Although xenon is chemically inert, it forms compounds with atoms of fluorine, oxygen, etc., small in size and electronegative atoms. It forms XeF₂, XeF₄, and XeF₆ type molecules based on the number of reactive orbitals.

This article is about xenon difluoride formula XeF₂, its structure of XeF₂, properties, and uses. Have a look.

History of Xenon Difluoride

Sir William Ramsay is known for discovering ‘inert gasses or Noble gasses’ in Chemistry. For this great work, he received the Nobel Prize in 1904. They were also termed ‘inert gasses’ due to their poor reactivity. Xenon is also one of the members of the inert gas family, but because of its large size and more reactivity than others, it easily forms a bond with other elements.

It is believed that, in 1962, German Chemist Rudolf Hoppe first developed xenon difluoride, XeF2. He produced it by mixing fluorine and xenon in an electrical discharge process. However, the records showed that in 1962, xenon difluoride was synthesized by C. L. Chernick, J. L. Weeks, and M. S. Matheson is a petrochemical reaction between xenon and fluorine.

Bonding in XeF₂

In the past, many chemists believed that noble gasses could not proceed. But producing the first accurate chemical compound with xenon promptly raised many questions. This compound was [Xe]⁺[PtF₆]^– and made people wonder whether the previous bonding models were still valid. After proper research, it became understandable that similar representations are claimed for halogen oxy and interhalogen species. Two methods of bonding that plays a vital role are

• The molecular orbital model
• The valence shell electron pair repulsion theory (VSEPR).

After this discovery, many xenon compounds have been prepared, mainly with the most electronegative elements like fluorine and oxygen.

XeF₂ Structure and Geometry

A Xenon Difluoride atom has eight electrons in its valence shell, i.e., 5s² 5p⁶ 5d⁰. It uses only two electrons to form two Xe-F 𝜎-bonds. The remaining six electrons remain as three lone pairs of electrons on the Xe-atom.

The electron dot structure shows that Xe-atom is surrounded by two 𝜎-bond pairs and three lone pairs. Therefore, Xe-atom is sp³d hybridized in the XeF₂ molecule.

The ground state electronic configuration of xenon is [Kr] 4d¹⁰ 5s² 5p⁶.

The excited state electronic configuration of xenon is [Kr] 4d¹⁰ 5s² 5p⁵ 5d¹.

This way the xenon difluoride formula becomes XeF₂.

In the exciting state, xenon has two single electrons, one in the 5p subshell and one in the 5d subshell. While each fluorine atom has one unpaired electron in its 2p orbital, and it needs one electron to get an inert gas configuration. Therefore, two F-atoms share their electrons with electrons of the Xe-atom and get bonded with it covalently.

According to VSEPR theory, the three lone pairs occupy the equatorial hybrid orbitals because, in this case, the electron pair repulsion is minimum, and the molecule gets maximum stability. Therefore, the spatial arrangement of five sp3d hybrid orbitals around the central Xe-atom is trigonal bipyramidal. The structure of the XeF2 molecule, due to the presence of three lone pairs in the equatorial positions, gets distorted. And XeF₂ structure becomes linear.

As in sp3d hybridization, two hybridizations are involved, i.e., sp² and pd. sp² hybridized orbitals form equatorial bonds, while pd hybridized orbitals form axial bonds. The percentage of s-character in sp² is 33.33%, while in pd, it is zero.

Because there is immense repulsion between the lone pairs on the Xe-atom, they attain larger bond angles and possess equatorial positions. The bond angle formed by equatorial bonds is 120°. Hence, lone pairs of xenon form vertices of the triangle at equatorial positions. At the same time, fluorine atoms are placed at the axial position from Xe-atom. Therefore, the structure of XeF₂ is linear, while the geometry becomes trigonal bipyramidal.

Synthesis of Xenon Difluoride

Xenon reacts directly with fluorine under 673 K temperature and 1 bar atomic pressure to form xenon difluoride formula XeF₂. The chemical reaction that takes place is

Xe (g) (in excess) + F₂ (g) → XeF₂ (s)

Here, xenon and fluoride are taken in a 2:1 ratio. If the ratio of the atoms gets different, the product obtained will change. Therefore, tracking the ratio of xenon and fluorine atoms during the reaction is necessary.

XeF₂ can be prepared by bypassing xenon and fluorine through a sealed Ni tube at 400℃.

It can also be prepared by photochemical combination, i.e., by irradiating a mixture of fluorine and xenon with sunlight or light from a high-pressure mercury arc lamp or in the presence of Hg vapors.

Therefore, the xenon difluoride obtained is solid. It is purified with the help of fractional distillation or selective condensation using a vacuum tube. It can also be synthesized by using dioxygen difluoride.

Properties of Xenon Difluoride

1. Physical properties of XeF₂:

• XeF₂ is soluble in solvents like BrF₃, BrF₅, anhydrous hydrogen fluoride, and IF₅.
• The molar mass of XeF₂ is 169.290 g.mol^-1.
• It is a dense white crystalline solid and sublime readily at 298 K.
• XeF₂ converts into higher fluoride when subjected to heat and pressure.
• The melting point of XeF₂ is 140℃.
• It is an exothermic compound and has a heat of formation of -125.5 kJ/mol.
• Its shape is linear.
• The solubility of XeF₂ in water is 25g/L at 0℃.
• It is corrosive to exposed tissues.
• It has a low vapor pressure and a disgusting odor.
• The Xe-F bond length in solid XeF₂ is 200 picometres, while in the vapor state, it is 197.73 picometres.

2. Chemical properties of XeF₂:

• XeF₂ is moisture sensitive and the most stable compound.
• In the presence of H₂O, XeF₂ is reduced to Xe.

2XeF₂ + 2H₂O → 2Xe + 4FH + O₂

• An aqueous solution of a base rapidly hydrolyses XeF₂.

2XeF₂ + 4OH^– → 2Xe + 4F^– + 2H₂O + O₂

• XeF₂ oxidizes Cl^– to Cl₂ and Ce(III) to Ce(IV).

XeF₂ + 2HCl → 2HF + Xe + Cl₂

SO₄^-2 + XeF₂ + Ce₂(SO₄)₃ → 2Ce(SO₄)₂ + Xe + F₂

• It is a mild fluorinating agent.
• It reacts with fluoride ion acceptors to form cationic species and with fluoride ions to form fluoroanion.

XeF₂ + PF₅ → [XeF]+[PF₆]^–

• In the presence of moisture, it releases toxic compounds.

Applications of Xenon Difluoride

Some applications of XeF₂ are

• It operated as a strong fluorinating agent.
• It can also work as an oxidizing agent.
• It can detect and determine the amount of iodine.
• XeF₂ can analyze selenium, sulfur, and tellurium in a number of compounds.
• It produces toxic and explosive substances. Therefore, caution must be taken while using XeF₂ when it reacts with moisture.
• It is an isotropic gaseous notch for silicon, especially in manufacturing microelectromechanical systems.
• When reacted with uracil, the product obtained is used as an anticancer drug 5-fluorouracil.
• It has a high imprint rate and does not require extrinsic energy or ion bombardment for the engraving of silicon.

Conclusion

XeF₂ is an example of an AB₂(lp) 3-type species. Lewis’s structure of XeF₂ shows that since five electron pairs surround Xe-atom, the spatial arrangement of five electron pairs around the central atom is trigonal bipyramidal.

Due to the presence of three lone pairs, the XeF₂ molecule assumes a linear shape with an F-Xe-F bond angle equal to 180°. The length of each of the Xe-F bonds is equal to 2.0 Å.

Frequently Asked Questions

1. Why do noble gases not participate in chemical combinations?

Noble gases are chemically inert due to the following reasons:

• They have a fulfilled valence shell.
• They have high ionization enthalpies.
• Their electron gain enthalpies are positive.

Therefore, they neither tend to lose nor gain electrons. Hence, they do not enter into chemical combinations.

2. Why do noble gases have comparatively large atomic sizes?

The atomic radii of noble gases are by far the largest in their respective periods. Noble gasses have only van der Waals radii, while other molecules have covalent radii. By definition, van der Waals radii are larger when compared to covalent radii. Hence, noble gases have equivalently large atomic sizes.

3. Being a noble gas, Xe forms a XeF₂ molecule. Why?

Although Xe-atom is an inert gas and has fulfilled valence shell orbital, its inner electrons shield the outer electrons from the nucleus because of the large size of the xenon. Hence, its valence electrons experience a weaker attraction force from the nucleus. As a result, the highly electronegative and small-sized atoms target the valence electrons of the xenon atom. Therefore, valence electrons of xenon are attracted by fluorine, and they bond and form a XeF₂ molecule.