A molecule has a net dipole moment if it contains polar bonds and its molecular geometry is such that the individual bond dipoles do not cancel each other out. This usually happens when the molecule is asymmetric or has lone pairs on the central atom.

Ammonia (${\mathrm{NH}}_3$): The central nitrogen atom in ${\mathrm{NH}}_3$ has three bond pairs with hydrogen atoms and one lone pair of electrons. This gives it a trigonal pyramidal geometry. The N-H bonds are polar, and due to the pyramidal shape and the presence of the lone pair, the individual bond dipoles do not cancel out. Instead, they add up, resulting in a net dipole moment for the molecule.

• (A) ${\mathrm{CO}}_2$: Carbon dioxide has a linear geometry. Although the C=O bonds are polar, the two bond dipoles are equal in magnitude and opposite in direction, causing them to cancel each other out. Thus, ${\mathrm{CO}}_2$ has no net dipole moment.
• (B) $p$-dichlorobenzene: In $p$-dichlorobenzene, the two chlorine atoms are attached to opposite carbon atoms on the benzene ring. The C-Cl bonds are polar, but due to the symmetrical para-positioning, the two C-Cl bond dipoles are equal in magnitude and opposite in direction, leading to their cancellation. Therefore, $p$-dichlorobenzene has no net dipole moment.
• (D) ${\mathrm{CH}}_4$: Methane has a tetrahedral geometry. The C-H bonds are only slightly polar, but more importantly, due to the highly symmetrical tetrahedral arrangement, the four C-H bond dipoles cancel each other out. Thus, ${\mathrm{CH}}_4$ has no net dipole moment.

Correct Answer: (C)