Polar or nonpolar is XeF4?

Xenon Tetrafluoride (XeF4) was the first noble gas binary compound identified (Xenon). An exothermic chemical reaction between Xenon (Xe) and Fluorine (F) generates XeF4. Two approaches, NMR spectroscopy and X-ray crystallography, were used to determine the structure of a crystalline, colourless material.

Is XeF4 polar or non-polar? Because of its symmetrical square planar structure, xenon tetrafluoride (XeF4) is a non-polar chemical molecule. Due to the uneven electronegativity of Xe and F atoms, the individual Xe-F bonds are polar, but the net vector sum of the polarities of Xe-F bonds is zero, as they cancel each other out. Therefore, XeF4 has a net dipole moment of 0 Debye.

Orientation of XeF4

Polarity of Xe-F Bonds

Xenon’s electronegativity is 2,6 while fluorine’s is 3,98.

Difference in electronegativity = (3.98 – 2.60) = 1.38

The difference is rather substantial. This indicates that Xenon-Fluorine (Xe-F) connections are polar in nature.

Therefore, for a single Xe-F bond, Fluorine will strongly draw electrons towards itself, leading to the accumulation of a positive charge on the Xenon side and a negative charge on the Fluorine side.

Polarity of the entire XeF4 structure

XeF4 has a square-planar molecular geometry. The overall structure of the concerned molecule is symmetric, as the two lone pairs are arranged axially opposite one another.

Polarity of a molecule is determined by the vector sum (both magnitude and direction are considered) of all bond dipoles.

In XeF4, there are four Xe-F bonds, each with a magnitude and a direction-specific dipole moment. When the vector total of these dipoles is calculated, the result is a net-zero Debye.

Xenon is surrounded by twelve atoms: eight from the four fluorine bonds (two from each bond) and two lone pairs of electrons. Four of these electrons are non-bonding.

Non-bonding orbitals have a greater tendency to spread than bonding orbitals. To minimise electron-pair repulsion, the localised non-bonding electron pairs are positioned opposite each other.

As a result of the symmetric arrangement of the Xe-F bonds and the non-bonding electron pairs, the dipole moment of the entire molecule is zero. Therefore, XeF4 is naturally a non-polar molecule (or compound).

Factors influencing a molecule’s polarity

Numerous factors contribute to determining whether a molecule is polar or non-polar. Polarity of a molecule is determined mostly by its dipole moment.

If the dipole is significant (more than or equal to 0 Debye), the molecule is polar.

The dipole moment of a molecule, and hence its polarity, is substantially determined by:

1.) Structure: Molecules with a symmetric arrangement of atoms likely to have a vector sum of dipole moments that is net zero.

This is due to the fact that for every polar bond, there will be another polar bond with a dipole that is equal in magnitude and opposite in direction, resulting in a net addition of zero.

Thus, symmetric molecules are non-polar, whereas asymmetric molecules are polar.

Molecules whose bonds are based on weak Vander Waal forces are often non-polar in nature. This is because Vander Waal forces are so weak that they do not contribute to the accumulation of charge at the molecule’s poles.

In contrast, the presence of strong Hydrogen bonds results in the accumulation of charges and polarity of the molecule.

Electronegativity of atoms: If the atoms that create a bond have similar electronegativity or if the difference between their electronegativity is less than 0.4%, then the bond is considered non-polar.

Check out the article on the nonpolarity of CBr4 for further information. Any dipole moment produced by such a molecule is negligible and does not contribute to polarity formation.

Charge accumulation: When there is an accumulation of negative charges at one end of the molecule and positive charges at the other end, as well as a significant dipole moment, the molecule is characterised as polar.

For instance, NCl3 is a polar molecule due to the difference in atomic number between chlorine and nitrogen. This page discusses the polarity of NCl3. Otherwise, the substance is non-polar.

XeF4 – Formation and Structure

Structure & Hybridization

The hybridization of XeF4 occurs at the Xenon central atom (Xe). Xe possesses six electrons in the 5p orbital and two in the 5s orbital, while the f- and d- orbitals are unoccupied.

To fill the vacancies, two electrons (in an excited state) from the 5p orbital travel to the 5d orbital. This leads in the hybridization of sp2d2.

Four fluorine atoms are then inserted on either side of the centre atom and bonded to the four half-filled orbitals. The hybridization of Xenon and Fluorine atoms indicates that the molecule’s shape is octahedral.

According to the VSEPR (Valence Shell Electron Pair Repulsion) Theory, the repulsion between the two lone electron pairs must be reduced.

This is conceivable when the Fluorine atoms occupy the equatorial locations and the atoms are arranged opposite one another.

Consequently, the overall molecular geometry of XeF4 is square and planar, with bond angles of 90° or 180°.

You must read an article on XeF4 lewis structure and geometrical structure for in-depth knowledge.

Synthesis of XeF4

The production of Xenon Tetrafluoride is an exothermic reaction that releases approximately 251 kJ/mol of net energy.

The corresponding chemical equation can be expressed as:

Xe + 2F2 ——> XeF4

On heating a mixture of Xenon (Xe) and Fluorine (F2) with a molecular ratio of 1:5 in a nickel (Ni) container or tube at 6 atmospheres and 400 degrees Celsius, the two elements react to generate XeF4 (Xenon tetrafluoride).

In this reaction, nickel does not function as a catalyst. Ni in the containers combines with Fluorine to generate a protective, non-peeling layer of NiF2 (Nickel fluoride) on the interior walls of the tube or container.

Qualities of XeF4

Molecular mass equals 207.29 g/mol

Density = 4.10 g/cm3

Vapour pressure = 3mm (at room temperature)

Bond angle = 90 or 180 degrees

Boiling point = 115.7°C

Point de fusion = 116°C. It has a greater m.p. than non-polar molecules because it is composed of a noble gas, which is naturally unreactive.

XeF4 is sparingly soluble in anhydrous HF (Hydrogen fluoride) and readily interacts with water to create Xenon Oxide.

The following chemical equation illustrates the reaction between XeF4 and water.

6XeF4 + 12H2O ———-> 2XeO3 + 24HF + 4Xe + 3O2

At room temperature, XeF4 appears in the form of colourless crystals and colourless fumes.

In its pure form, xenon tetrafluoride is stable, but it must be protected from moisture. It can be stored in Nickel or Monel containers indefinitely.

XeF4 Applications and Utilizations

XeF4 has relatively few applicability in general. Some of them include the following:

Degradation is required to assess the trace metal impurities in silicone rubber. XeF4 is a reliable reagent for the decomposition of silicone rubber.

Its interaction with silicone produces simple gaseous products and leaves behind metal contaminants.

Noble gases are regarded as fully inert, meaning they will never react with another element to form a combination.

Since XeF4 is a compound consisting of a noble gas (Xenon), it is incredibly intriguing to chemists. Therefore, it is commonly employed for study.

It is utilised as both an oxidising agent (to convert iodide to iodine) and a fluorinating agent.

For the synthesis of higher fluorides of Xenon (XeF6), F2 and XeF4 can undergo a reaction.

XeF4 + F2 ——-> XeF6


As previously explained, the XeF4 molecule has a symmetrical square planar structure, resulting in all XeF4 bonds possessing an equal and opposing polarity. Due to the difference in electronegativity of both atoms, Xe and F form a covalent polar bond and a net dipole. However, all Xe-F bonds cancel each other’s dipoles, rendering the total molecule nonpolar.

If you have questions about polarity, please ask. Ask them in the comment area below. I will also clarify your confusion in this regard.

Read more: Structure, Molecular Geometry, and Hybridization of H2O

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.


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