3.6.1—Hydrogen bonding

Hydrogen bonding is a strong intermolecular attraction and a special case of permanent dipole–permanent dipole attraction. For the assessed N–H and O–H examples, the donor molecule has a highly polar bond to hydrogen and the neighbouring molecule supplies a lone pair on an electronegative nitrogen or oxygen atom.

The electronegative atom pulls bonding electrons away from hydrogen, giving H a substantial δ+ charge. That hydrogen is attracted to a lone pair on N or O in another molecule; the attraction is strongest when the covalent and hydrogen bonds are close to linear. Ammonia and water therefore form intermolecular hydrogen-bond networks, with the number of possible links limited by available N–H/O–H hydrogens and lone pairs.

In water, hydrogen bonds hold molecules together, so more energy is needed to separate them: melting and boiling points are relatively high. At the surface, neighbouring molecules pull surface molecules inward, producing high surface tension. In ice, a rigid, more open hydrogen-bonded arrangement leaves molecules farther apart than in liquid water, so ice has a lower density.

Do not call the O–H or N–H covalent bond itself a hydrogen bond: the hydrogen bond is the intermolecular attraction to a neighbouring lone pair. Do not infer hydrogen bonding from the presence of hydrogen alone, and do not replace the network/packing explanation with the claim that all solids are denser than their liquids.