Dipole Moments and Polar Molecules

In most covalent bonds the electron-attracting powers of the atoms differ and the opposite ends of the bond acquire slight charges.  Such a bond is known as a polar covalent bond.  A diatomic molecule having a polar covalent bond, such as HCl, is said to have a dipole moment. The dipole moment (µ) is defined as the product of the magnitude of the charge (d) at either end of the dipole multiplied by the distance (d) that separates the charges.

µ = d · d

Bond Dipoles and Molecular Dipoles

Carbon dioxide, CO₂, has two C=O bonds.  Each of these bonds is polar covalent and thus has a dipole moment, but the CO₂ molecule is non polar.  To explain this we need to distinguish between bond dipole moments which describe the magnitude of charge separation in individual bonds and molecular dipole moments, which describe charge separation in the whole molecule taking every bond into account.  The following shows why CO₂ is nonpolar.

The cross-based arrows show the directions of the bond moments. Since the molecule is linear, these two bond moments, which are equal in magnitude and opposite in direction, cancel each other and the resultant dipole moment for the molecule is zero. 

Molecular Shapes and Dipole Moments

We can use the following three steps to predict whether a molecule is polar or nonpolar.

• Use electronegativity data to predict the existence of individual bond moments.

• Use VSEPR theory to predict molecular shape.

• From molecular shape, determine whether bond moments cancel to give a nonpolar molecule.

A rule of thumb to use is that any molecule with the maximum symmetry possible for that type is nonpolar.