Is XeF2 Polar or Nonpolar?

The chemical name of XeF2 is xenon difluoride which is a dense crystalline solid. XeF2 is one of the most stable compounds of xenon but discomposed when comes in contact with light or water vapor otherwise it is stable to storage.

Xenon difluoride is synthesized simply by reacting xenon and fluorine gas which results in the formation of a solid XeF2.

Xe + F2 → XeF2

This reaction is an endothermic reaction because it required external heat or electric energy to proceed. The produced solid product undergoes fractional distillation for purification.

Here in this article, we are going to know whether XeF2 is a polar or nonpolar molecule with a detailed explanation. But before entering into XeF2 polarity, let's have some ideas about what polar and nonpolar molecule are:

Polar Molecules

Polar molecules are those molecules that have a net dipole moment and this is possible only if there is an electronegativity difference between the associated atoms as well as asymmetrical geometry so that the induced dipole charges do not cancel each other.

There are some certain conditions for the molecule to be polar i.e. have asymmetrical structure, presence of lone pair of electrons or electronegativity difference between central atom to bonded atoms.

Polar molecules have polar covalent bonds or ionic bonds. They are soluble in water (called hydrophilic) and other polar solvents like hydrochloric acid (HCl), ammonia (NH3), Nitrogen dioxide (NO2) sulfur dioxide (SO2) etc.

Nonpolar Molecules

Nonpolar molecules are considered as a pure covalent bond because it forms by equal sharing of electrons between the atoms in a compound and this is what a covalent bond is known for.

One more case for the nonpolar molecule if the molecules have symmetrical structure so that the induced partial charges on the individual atoms nullify each other and the net dipole moment becomes zero.

As these are a pure covalent bond so they are insoluble in water (called hydrophobic) and other polar solvents but are soluble in nonpolar solvents like carbon tetrachloride (CCl4), carbon dioxide (CO2), benzene (C6H6) etc.

Now, is XeF2 polar or nonpolar? XeF2 is a nonpolar molecule despite two Xe-F bonds are polar. The XeF2 has a linear molecular geometry and Xe-F bonds are symmetrical to each other, as a result, the net dipole moment becomes zero.

Why is XeF2 a nonpolar molecule? [Detailed Explanation]

As we have already known there are certain conditions for a molecule to be polar or nonpolar in the above definitions of the polar molecule and nonpolar molecule. The same things apply here which makes XeF2 a molecule.

In XeF2, there are two fluorine atoms are bonded with a central xenon atom. Also, there are three lone pair of nonbonding electrons are localized on the central Xe atom arranged symmetrically.

The Xe-F bonds are polar as the electronegativity difference between xenon (2.6) and fluorine (4.0) is 1.4. But we know molecular polarity is a vector quantity which means the net polarity is the vector sum of individual induced bond dipoles.

The molecule is symmetrical so that all the Xe-F bond induced dipoles are opposing each other, as a result, net dipole moment on the molecule becomes zero.

But as shown in the above diagram, there are still three pairs of non-bonding electrons are present on central Xe atom which is aligned in equatorial position so that their effects are also nullified. And there is not any dipole charge presence on the molecule which makes XeF2 a nonpolar molecule.

Also, Checkout: Is XeF4 Polar or Nonpolar?

XeF2 Polar or Nonpolar (Based on polarity determining factors)

These are some of the major factors which are used to determine the polarity of the compounds:

Electronegativity difference

Electronegativity is a kind of force exerted by atoms on their bonding partners at the time of bond formation. All the atoms have their own electronegativity value which varies from atom to atom. Some have a higher EN value and some have lower EN values.

The atoms having higher electronegativity value attract the electrons pair from the bonding partners. It means higher the electronegativity difference between atoms, the more uneven sharing of electrons takes place. In the periodic table, fluorine (F) atom has the highest electronegativity value 4.0, and francium (Fr) has the lowest 0.7.

In the XeF2 compound, the electronegativity difference between fluorine (3.98) and Xe (2.6) is 3.98-2.6= 1.38 which shows the Xe-F bond is polar according to the Pauli scale.

Molecular geometry

The central atom of XeF2 (Xe) has a total of 10 electrons localized in the form of 4 as Xe-F bonding and 6 as non-bonding pairs.

And due to the symmetrical structure, the polarity of the Xe-F bond cancels each other and no dipole charge induced on the molecule same likes in the XeCl2 molecule.

Dipole moment

The dipole moment is defined as the products of induced charge and distance of separation between the atoms. It is given as,

Dipole moment = Charge (Q) * distance of separation (d)

Its unit is Debye and denoted by ‘D’. 1D = 3.33564*10-30 C.m, where C is Coulomb and m is meter.

The dipole moment is the major aspect for any compound to be polar or nonpolar. All the polar molecules have a net dipole moment and all nonpolar have zero dipole moment.

XeF2 is a nonpolar molecule so that its dipole moment is 0D (Zero).

Properties of XeF2

• Xenon difluoride has a density of 4.32 g/cm3, molecular weight 169.29 g/mol, boiling point 155 °C, and melting point 128.6 °C.
• As XeF2 is a nonpolar molecule, it is insoluble in water but soluble in some of the nonpolar solvents like BrF5, BrF3, IF5, anhydrous HF, etc.
• The oxidation state of xenon in xenon difluoride is +2

Uses of XeF2

• XeF2 is used as a fluorinating agent because it is selective about which atom it fluorinates.
• It is used in microelectromechanical system (MEMS) as an isotropic gaseous etchant for silicon.
• It is also used as antiseptic, used to kill bacteria and fuel for ion engines.
• Sometimes it is used as tracing a very little amount of iodine and iodine-based compounds.

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