Which of the following molecules would you expect to have a resultant dipole moment \((\mu) ;\) (a) \(\mathrm{F}_{2}\), (b) \(\mathrm{NO}_{2}\) (c) \(\mathrm{BF}_{3},\) (d) \(\mathrm{HBr},\) (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\), (f) \(\mathrm{SiF}_{4},\) (g) OCS? Explain.

The first thing to do for each of these molecules is to determine their molecular geometry. (a) \(\mathrm{F}_{2}\) is a diatomic molecule, so its geometry is linear. (b) \(\mathrm{NO}_{2}\) is a linear molecule due to the presence of an unpaired electron on the nitrogen atom. (c) \(\mathrm{BF}_{3}\) has a trigonal planar geometry as boron forms three bonds and does not have any lone pairs. (d) \(\mathrm{HBr}\) is diatomic, so it is linear. (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\) is tetrahedral as the carbon atom is bonded to four other atoms. (f) \(\mathrm{SiF}_{4}\) is also tetrahedral with silicon at the center. (g) OCS is a linear molecule.

Next, analyze the polarity of the bonds. In molecules where atoms have different electronegativities, the electrons aren't shared equally and the bond is polar. (a) In \(\mathrm{F}_{2}\), the two fluorine atoms have the same electronegativity, so the bond is not polar. (b) \(\mathrm{NO}_{2}\), both nitrogen and oxygen have different electronegativities and the bonds are polar. (c) For \(\mathrm{BF}_{3}\), boron and fluorine have different electronegativities, so the bonds are polar. (d) In \(\mathrm{HBr}\), hydrogen and bromine have different electronegativities, so the bond is polar. (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\), carbon-hydrogen and carbon-chlorine bonds are polar due to the differences in electronegativities. (f) \(\mathrm{SiF}_{4}\), the silicon-fluorine bonds are polar due to differences in electronegativity. (g) For OCS, oxygen-carbon and carbon-sulfur bonds are polar.

Now, consider the arrangement of the polar bonds (if any) in the molecule. If they cancel out due to symmetry, the molecule does not have a net dipole moment.(a) Since there are no polar bonds in \(\mathrm{F}_{2}\), it does not have a net dipole moment.(b) As \(\mathrm{NO}_{2}\) has a linear shape, the dipole moments do not cancel out and it has a net dipole moment.(c) The polar bonds in \(\mathrm{BF}_{3}\) cancel out due to its symmetrical shape and it does not have a net dipole moment.(d) \(\mathrm{HBr}\) has a net dipole moment as it's a linear molecule with a polar bond.(e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\) has a net dipole moment due to the difference between hydrogen and chlorine in electronegativity.(f) The polar bonds in \(\mathrm{SiF}_{4}\) cancel out due to its tetrahedral shape, so it does not have a net dipole moment.(g) OCS has a net dipole moment due to its linear shape and the polarity of its bonds.