The net dipole moment of a molecule depends on its geometry and the polarity of its individual bonds. If the vector sum of all bond dipoles is zero, the molecule has a net dipole moment of zero.

The molecule ${\mathrm{BF}}_3$ has a trigonal planar geometry. Boron is at the center, and three fluorine atoms are bonded to it. The B-F bonds are polar, but due to the symmetrical trigonal planar arrangement (bond angles of 120°), the three bond dipoles cancel each other out vectorially. Therefore, ${\mathrm{BF}}_3$ has a net dipole moment of zero.

• A) $\mathrm{HF}$: This is a diatomic molecule with a polar covalent bond (due to the electronegativity difference between H and F). It has a linear geometry and a net dipole moment.
• B) ${\mathrm{H}}_2\mathrm{O}$: Water has a bent (V-shaped) geometry due to two lone pairs on the oxygen atom. The O-H bonds are polar, and their dipoles do not cancel out, resulting in a significant net dipole moment.
• C) ${\mathrm{BF}}_3$: As explained above, its trigonal planar geometry leads to the cancellation of bond dipoles, resulting in a net dipole moment of zero.
• D) ${\mathrm{CHCl}}_3$: Chloroform has a tetrahedral geometry. The C-Cl bonds are polar, and the C-H bond is also slightly polar. Due to the different bond polarities and the tetrahedral arrangement, the bond dipoles do not perfectly cancel out, leading to a net dipole moment.

Correct Answer: (C)