Diagram showing the net effect of symmetrical polar bonds (direction of yellow arrows show the migration of electrons) within boron trifluoride cancelling out to give a net polarity of zero. ?- shows an increase in negative charge and ?+ shows an increase in positive charge.

The symmetrical polar bonds within boron trifluoride cancel out to give a net polarity of zero.

It is easy to become accustomed to thinking only about whether or not the bonds within a molecule are polar in deciding the intermolecular force relationships it can engage in. Typically, you examine the electronegativity differences between bonded atoms and assign the polarity of the bonds and go from there to decide whether the molecule is nonpolar or polar.

However, the polarity of bonded atoms is not the only factor in determining the overall polarity of a molecule, and by extension, intermolecular force. Molecular geometry as predicted by VSEPR is also important.

For example, if the various dipole moments within a molecule are oppositely directed in space, even though the particular bonds are polar, the molecule may have no net dipole moment at all, in which case the degree of intermolecular force is weak. Carbon dioxide is the quintessential example, a linear molecule in which the dipole moments of the carbon-oxygen bonds are oppositely directed.

From a test-writer's perspective, this kind of exception makes a good multiple choice question. Keep VSEPR in mind on the MCAT if you are faced with interpreting the overall polarity of a molecule. Don't forget to ask yourself whether the dipole moments of individual bonds cancel each other out.