Dipole Moment - Practice Questions & MCQ

When the covalent bond forms between two different atoms then because of the difference in electronegativity the electrons get shifted towards the more electronegative atom and thus forms the polar covalent bond and this polar molecule is known as a dipole molecule. The more electronegative atom occupies a partial negative charge (δ-) and the other atom possesses a partial positive charge (δ+). This separation of charge gives rise to a bond dipole moment. The magnitude of a bond dipole moment is represented by the Greek letter mu (µ) and is given by the formula as shown below, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges:

μ=Qr

This bond moment can be represented as a vector, a quantity having both direction and magnitude as shown in the figure. Dipole vectors are shown as arrows pointing along with the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.

(a) There is a small difference in electronegativity between C and H, represented as a short vector. (b) The electronegativity difference between B and F is much larger, so the vector representing the bond moment is much longer.

The dipole moment measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.

For diatomic molecules, there is only one bond, so it's bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br₂ and N₂ have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.

When a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their dipole moments cancel, then the molecule is non-polar. For example, in the case of CO₂ as shown in the figure given below. Each of the bonds is polar, but the molecule as a whole is non-polar. The dipole moments cancel each other because they are pointed in opposite directions. 

In the case of the water molecule, the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).