Why do certain molecules have zero dipole moment?

Certain molecules have zero dipole moment because their polar bonds are arranged symmetrically, cancelling out each other's effects.

In more detail, a molecule's dipole moment is a measure of its overall polarity. It is determined by both the magnitude of the charge at each end of a molecule and the distance between these charges. A molecule with polar bonds (bonds between atoms of different electronegativities) can still have a zero dipole moment if the geometry of the molecule is such that the bond dipoles cancel each other out.

For instance, consider carbon dioxide (CO2). It has two polar C=O bonds, but the molecule is linear, meaning the two bond dipoles are equal in magnitude but opposite in direction. As a result, they cancel each other out, giving CO2 a dipole moment of zero.

Similarly, methane (CH4) has four polar C-H bonds, but the molecule is tetrahedral. This means the bond dipoles are arranged symmetrically around the central carbon atom, and their effects cancel each other out, resulting in a zero dipole moment.

However, not all molecules with polar bonds have a zero dipole moment. Water (H2O), for example, has two polar O-H bonds, but the molecule is bent, not linear or symmetric. This means the bond dipoles do not cancel each other out, and water has a nonzero dipole moment.

In summary, whether a molecule has a dipole moment depends on both the polarity of its bonds and its geometry. If a molecule has polar bonds and is symmetric, the bond dipoles can cancel each other out, resulting in a zero dipole moment. If the molecule is not symmetric, the bond dipoles may not cancel each other out, and the molecule will have a nonzero dipole moment.