Hydrogen bonding occurs when hydrogen is bonded to a strongly electronegative element: #F#; #O#; and #N# are typically the heteroatoms to which hydrogen is bound.
Because the heteroatom tends to polarize the electron density of the bond towards itself, some charge separation occurs, which is the very definition of #"polarity"#.
So for each molecule, we could depict this charge separation, this polarity, as #""^(delta+)H-F^(delta-)#, #""^(delta+)H-O^(delta-)-H^(delta+)#, and #""^(delta-)NH_3^(delta+)#. In the bulk solutions of water, ammonia, hydrogen fluoride, these dipoles line up to form an additional intermolecular force, that is responsible for the anomalously high boiling points of these solvents.
Note here the distinction between intermolecular force, bonds between molecules, versus intramolecular force, bonds between atoms within molecules.
By way of contrast, we could examine non-polar molecules. #CH_4# is a carbon hydride, however, here the difference in electronegativity between carbon and hydrogen is fairly marginal (certainly it is less than that between hydrogen and oxygen). The #C-H# bond is fairly non-polar, and as a result does not exhibit any significant intermolecular interaction. The result? Methane is a room temperature gas with low melting and boiling points.
