Xenon Tetrafluoride (XeF4) was the first discovered binary chemical compound of a noble gas (Xenon). An exothermic chemical reaction of Xenon (Xe) with Fluorine (F), leads to the production of XeF4. A colorless, crystalline substance, its structure has been established by two methods: NMR Spectroscopy and X-Ray crystallography.
Is XeF4 polar or non-polar? Xenon tetrafluoride (XeF4) is a non-polar chemical compound, owing to its symmetrical square planar structure. The individual Xe-F bonds are polar due to unequal electronegativity of Xe and F atoms, but the net vector sum of the polarities of Xe-F bonds is zero as they cancel out each other. So, the net dipole moment of XeF4 is 0 Debye.
Polarity of XeF4
Bond Polarity of Xe-F
The electronegativity of Xenon is 2.6 and that of fluorine is 3.98.
Electronegativity difference = |3.98 – 2.60| = 1.38
The difference is pretty significant. This means that the individual Xenon-Fluorine (Xe-F) bonds are polar in nature.
So, for a single Xe-F bond, Fluorine will attract electrons very strongly towards itself and this will lead to build-up of a positive charge on the Xenon side and negative charge on the Fluorine side.
The polarity of the complete XeF4 structure
The molecular geometry of XeF4 is square planar. With the two lone pairs placed axially opposite to each other, the overall structure of the concerned molecule is symmetric.
The molecular polarity is determined by the vector sum (magnitude and direction both are considered) of all individual bond dipoles.
In XeF4, there are four Xe-F bonds, each having an individual dipole moment with a magnitude and a direction. When we take the vector sum of these dipoles, it adds up to be a net-zero Debye.
There are twelve atoms around Xenon- eight from the four fluorine bonds (two from each bond) and two lone pairs of electrons. Of these, four are non-bonding electrons.
Non-bonding orbitals tend to diffuse more than bonding orbitals. Therefore, the localized non-bonding electron pairs are placed opposite to each other for minimum electron-pair repulsion.
Owing to the symmetric arrangement of the Xe-F bonds and the non-bonding electron pairs, the net effective dipole across the whole molecule is zero. Therefore, XeF4 is a non-polar molecule (or compound) by nature.
Factors influencing the Polarity of a molecule
There are numerous factors that help decide whether a molecule is polar or non-polar. The polarity of a molecule is basically the measure of its dipole moment.
If the dipole is a considerable number (not equal to 0 Debye), the molecule is said to be polar.
The dipole moment, and hence the polarity of any molecule largely depends on:
1.) Structure: Molecules with the symmetric arrangement of atoms that form them, tend to have a net-zero vector sum of dipole moments.
This is because, for every polar bond, there will be another polar bond with a dipole that is equal in magnitude and opposite in direction, making the net addition equal to zero.
Thus, symmetric molecules are non-polar in nature while those without symmetricity are polar.
2.) Type of bond: Molecules that are based on weak Vander Waal forces are normally non-polar in nature. This is because Vander Waal forces are so fragile that they contribute to a nil charge build-up at the poles of the molecule.
On the other hand, the existence of strong Hydrogen bonds leads to an amassing of charges and result in the polarity of the molecule.
Electronegativity of atoms: If the atoms that form any bond are similar or the difference in their electronegativity is less than 0.4, then the bond so formed is considered non-polar.
Check out the article for the non-polarity ofCBr4. Any dipole moment developed in such a molecule is insignificant and does not help establish polarity.
Charge build-up: When there is an accumulation of negative charges at one end of the molecule and positive charges on the other, along with a significant dipole moment, the molecule is said to have some polarity.
For example, NCl3 is a polar molecule due to the difference between chlorine and nitrogen atoms. Here is the article for the polarity of NCl3. Otherwise, it is considered non-polar.
XeF4 – Structure and Formation
Structure & Hybridization
The hybridization in XeF4 takes place in the central atom Xenon (Xe). The valence shell of Xe has six electrons in 5p orbital and two in 5s, while the f- and d- orbitals are empty.
So, two of the electrons (in an excited state) from the 5p orbital move to 5d orbital to fill the vacancies. This results in an sp2d2 hybridization.
Four fluorine atoms then bond with the four half-filled orbitals and are placed on either side of the central atom. The hybridization of Xenon and Fluorine atoms suggests the molecule to have octahedral geometry.
According to the VSEPR (Valence Shell Electron Pair Repulsion) Theory, the repulsion has to be minimized between the two lone pairs of electrons.
This is possible when they are placed opposite to each other, and the Fluorine atoms occupy the equatorial positions.
Hence, the overall molecular geometry of XeF4 comes out to be square planar, with a bond angle equal to 90° or 180°.
Formation of XeF4
The process of formation of Xenon Tetrafluoride is an exothermic process that releases net energy of about 251 kJ/mol.
The chemical equation for the same can be written as:
Xe + 2F2 ——> XeF4
On heating, a mixture of Xenon (Xe) and Fluorine (F2) in the molecular ratio 1:5, in a nickel (Ni) container or tube, at 6 atm., and a temperature of 400°C, the two elements react to form XeF4 (Xenon tetrafluoride).
Nickel in this reaction does not act as a catalyst. It is merely used because the Ni in the containers reacts with Fluorine to form a protective, non-peeling layer of NiF2 (Nickel fluoride) on the inner walls of the tube or the container.
Properties of XeF4
- Molecular weight = 207.29 g/mol
- Density = 4.10 g/cm3
- Vapour pressure = 3mm (at room temperature)
- Bond angle = 90° or 180°
- Boiling point = 115.7°C
- Melting point = 116°C. The m.p. is higher in comparison to non-polar molecules as it is a compound consisting of a noble gas, which is unreactive in nature.
- XeF4 is sparingly soluble in anhydrous HF (Hydrogen fluoride) and reacts readily with water (even small traces of moisture in the air) to form Xenon Oxide.
- The reaction of XeF4 with water can be depicted by the following chemical equation:6XeF4 + 12H2O ———-> 2XeO3 + 24HF + 4Xe + 3O2
- The appearance of XeF4 is in the form of colorless crystals and colorless vapors at room temperature.
- Xenon tetrafluoride is stable in its pure form, which must be kept free from any traces of moisture. It can be stockpiled indefinitely in Nickel or Monel containers.
Applications and Uses of XeF4
In general, XeF4 has very limited applications. Some of them are as listed below:
- To analyze the trace metal impurities in silicone rubber, it has to be degraded. For the degradation of silicone rubber, XeF4 is a reliable reagent.
Its reaction with silicone generates simple gaseous products and metal impurities are left behind.
- Noble gases are considered to be completely inert i.e. under no circumstance, they will react with another element to form a compound.
Since XeF4 is a compound involving a noble gas (Xenon), it is extremely fascinating to people dealing with chemistry. Hence, it is widely used for research purposes.
- It is used as an oxidizing agent (for the conversion of iodide to iodine) as well as a fluorinating agent.
- For the formation of higher fluorides of Xenon (XeF6), a reaction between F2 and XeF4 can be carried out.XeF4 + F2 ——-> XeF6
As discussed, the XeF4 molecule has a symmetrical square planar shape due to which all the XeF4 bonds have an equal and opposite dipole. Xe and F forms a covalent polar bond due to the difference in electronegativity of both atoms and also result in a net dipole. But all Xe-F bonds cancel out dipole of each other making the overall molecule a nonpolar.
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