Dipole Moment

Introduction

Dipole moments can be developed both in ionic and covalent compounds. It measures the separation of positive and negative charges in the molecule, and hence, can be considered for the measurement of polarity in molecules.

The separation of charges in any molecule results in a dipole moment. This separation of charges takes place due to the difference in electronegativity of the bonded atoms in a compound.

Between two bonded atoms in a bond, if one is more electronegative than the other, then the atom with more electronegativity attracts the bonded electron cloud towards itself. As a result of which the electron density on one bonded atom becomes more than on the other bonded atom. It also happens with the atom having lone pair electrons on it. To measure the polar character of a chemical bond between two atoms in a molecule, the bond dipole moment is considered. A bond dipole moment is a vector quantity because it has both magnitude and direction.

What is a dipole moment?

The dipole moment is the calculation of net polarity in each molecule. It is the separation of charges. When two charges of opposite sign and equal magnitude are separated by a distance (𝑟), an electric dipole is generated which again can be measured by its dipole moment (𝜇). With the help of a dipole moment, we find out the polar nature of a bond. The greater the value of the dipole moment, the more will be the polar

nature of the bond. Molecules with zero dipole moments are non-polar, while molecules with dipole moments are polar.

Let us know more-

• In the case of a polyatomic molecule, the dipole moment of a single bond is the bond dipole moment which is again different from the dipole moment of the whole molecule.

• The arrow used to show dipole moments is slightly different. It is denoted by an arrowhead on the negative centre and a cross on the positive centre. With this arrow, we can understand the shift of electron density in the molecule.

• The magnitude of the dipole moment can also be zero because the two opposite- direction bond dipoles can cancel each other.

• Molecules having zero dipole moment are nonpolar whereas molecules having a value of dipole moment are polar. For example, $\mathrm{CH_{2}O}$ is a polar molecule but 𝐶𝑂 is a nonpolar molecule as polar bonds are present in both of them.

Comparing polar and nonpolar molecules

Where do dipole moments occur?

In general dipole moments occur in any molecule where there is the separation of charges. Considering any molecule, such that, there is at least one such bond where the bonded two atoms are not having the same electronegativity. In this case, the more electronegative element withdraws the electron density towards itself and carries a partial (-ve) charge. At the same time, the other atom carries a partial (+ve) charge, i.e., two poles are generated which is called a dipole. To measure the magnitude of the generated dipole, the parameter dipole moment is used.

Dipole moment formula

The dipole moment is calculated as the product of the Charge (𝑞) and the distance of separation of the charges (𝑟). Dipole moments can be written as the product of the magnitude of electric charge and the internuclear distance between the atoms in a molecule. The unit for dipole moments is Debye. It is denoted by 𝐷. Dipole moment can be represented as -

$$\mathrm{\overline{\mu}\:=\:\Sigma\:q_{n}\:.\:\overline{r}_{n}}$$

Where,

$\mathrm{q_{n}}$ = magnitude of the nth charge.

$\mathrm{\overline{r}_{n}}$ = position of the nth charge.

1 𝐷𝑒𝑏𝑦𝑒 = 3.3356 x10−30 𝐶. 𝑚 where 𝐶 is Coulomb, and 𝑚 is metre.

Considering a simple two-charge separation like in diatomic molecules i.e., a bond dipole within a molecule, the dipole moment can be expressed as-

$$\mathrm{\mu\:=\:\delta\:.\:r}$$

Where

𝛿 = magnitude of partial charges 𝛿+ and 𝛿−.

𝑟 = distance between the partial charges.

Examples of dipole moments

Let us try to understand the meaning of dipole moments with different examples.

$\mathrm{NH_{3}}$ molecule.

$\mathrm{NH_{3}}$ has a pyramidal structure, where 3 𝑁— 𝐻 bonds are pointed downwards as in a tripod stand. The direction of the net dipole moment of the three 𝑁— 𝐻 bonds is upwards and lone pairs are also upward. Hence the resultant dipole moment of the molecule is 1.49𝐷.

Dipole moment of $\mathrm{NH_{3}}$ molecule

$\mathrm{BeF_{2}}$ molecule.

$\mathrm{BeF_{2}}$ is a linear molecule. The bond angle here is 1800 and the bond dipole of both the 𝐵𝑒— 𝐹 bonds are directly opposite. As a result of which, dipoles cancel each other and the net dipole moment is zero.

Dipole moment of $\mathrm{BeF_{2}}$ molecule.

$\mathrm{NF_{3}}$ molecule.

The $\mathrm{NF_{3}}$ molecule has a pyramidal structure where we have 3 𝑁— 𝐹 bonds and a lone pair on the nitrogen atom. Just like in $\mathrm{NH_{3}}$, the 3 𝑁— 𝐹 bonds in $\mathrm{NF_{3}}$ are pointed downwards as in tripod stands. Fluorine is more electronegative than nitrogen and attracts all the bonded electrons of the 𝑁— 𝐹 bond. The direction of the net dipole moment of the three 𝑁— 𝐹 bonds are downwards but the dipole moment direction due to lone pairs is acting upwards. But the two dipole moments don't cancel each other completely, so the net value of the dipole moment is 0.24D. It is to be noted that the dipole moment of $\mathrm{NF_{3}}$ is smaller than that of 𝑁𝐻3 because the dipoles are in the opposite direction in $\mathrm{NF_{3}}$.

Dipole moment of $\mathrm{NF_{3}}$ and $\mathrm{NH_{3}}$ molecule

Dipole moment of water $\mathrm{(H_{2}O)}$ molecule

Let us understand the dipole moment of water with the help of this diagram

Xzapro4 Dipole Water, CC BY-SA 3.0

According to VSEPR theory, when there are two lone pairs and two bond pairs in a molecule the shape of the molecule is Bent. Hence, the shape of the $\mathrm{(H_{2}O)}$ molecule is bent having two lone pairs and two bond pairs. Since the 𝑶-atom is more electronegative than the 𝑯-atom, the bonded electrons in the 𝑶— 𝑯 bonds are not shared equally. The electron density moved towards the O-atom. As a result, the 𝑶-atom carries the 𝛿− charge and the 𝑯-atom carries the 𝛿+ charge. Since there are two 𝑶— 𝑯 bonds, the net charge on the 𝑶-atom is 𝛿𝟐−. The resultant dipole moment of the two 𝑶— 𝑯 bonds are directed upwards. The value of the dipole moment in the $\mathrm{(H_{2}O)}$ molecule is 𝟏. 𝟒𝑫.

Conclusion

Any molecule where there is the separation of charges possesses a dipole moment. Measurement of dipole moment talks about the polarity of a molecule. It occurs in molecules where we have a chemical bond having two atoms of different electronegativity. It is not true that a molecule having polar bonds will be polar. There must be some value of the dipole moment (𝜇) for a molecule to be polar. If a molecule has polar bonds but the net dipole moment of the molecule is zero, then the molecule is said to be nonpolar. There are several applications of dipole moment. The dipole moment is calculated as the product of the Charge and the distance of separation of the charges. The unit for dipole moments is Debye. It is denoted by 𝐷.

FAQs

1. Discuss some of the applications of dipole moment?

Dipole moment is used to determine several parameters, for example percentage of ionic character in a molecule, the calculation of bond angle and the residual charge of atoms in the molecules, determination of size and shape of molecules etc.

2. Explain why the dipole moment of the following molecules $\mathrm{CO_{2}\:,\:BF_{3}\:and\:CCl_{4}}$ is zero?

In each of these cases, the molecules have symmetrical shapes. As a result, their dipoles are cancelled out, and the net dipole moment of the molecules becomes zero.

3. How do you find the largest dipole moment?

When the electronegativity difference between the bonded atoms in a chemical bond is large, the bond polarity also becomes large and hence the dipole moment is also large.

4. Does the $\mathrm{H_{2}}$ molecule have a dipole moment?

$\mathrm{H_{2}}$ is a linear molecule where both the bonded atoms have the same electronegativity. As a result, there is no separation of charges and hence no dipole moment.

5. Is $\mathrm{XeF_{4}}$ polar or nonpolar?

The shape of 𝑋𝑒𝐹4 is square planar. There are 4 𝑋𝑒— 𝐹bonds which are polar, but the net dipole moment of the molecule is zero due to the cancellation of the opposite direction. Hence 𝑋𝑒𝐹4 is a nonpolar molecule having polar bonds.