Boiling Points of the Covalent Halides
Hydrogen bonding is a very special intermolecular force that occurs in polar molecules when a hydrogen atom is bonded to a highly electronegative atom. The only possible atoms with a high enough electronegativity are N, O, and F. When a hydrogen is bonded to one of these, the sole electron on the hydrogen is shared to such an extent that the hydrogen atom has a much smaller radius. This allows it to approach very close to the N, O, or F atom on another molecule. Since the energy of attraction depends on the separation distance, this particularly close approach leads to a very low energy (very stable) configuration. Thus the effect of hydrogen bonds can be very large. They are so strong that they seem like weak covalent bonds. Thus, the name hydrogen bond. Below is a graph of the boiling point of some compounds like CH4, NH3, H2O, and HF as a function of row in the periodic table with the non-hydrogen element being replaced with every higher mass elements moving down the periodic table. Note the general trend that can be seen in the methane series (CH4, SiH4,....) as the compounds have larger and larger central atoms with more electrons farther from the nucleus they are more and more polarizable. Thus their boiling points increase due to increasing dispersion forces. However, note the clear exceptions for NH3, HF, and H2O. These are the only three compounds that can form hydrogen bonds. This gives them much stronger intermolecular interactions than might be expected from dipole-dipole and dispersion forces alone. Thus they have much higher boiling points than the trend would predict. You can also see that CH4, which has no hydrogen bonds, is not an exception and follows the trend to have the lowest boiling point in the series.
Boiling Points of the Covalent Halides