Hydrogen Bonding: Hydrogen bonding is the unusually strong dipole-dipole
interaction that occurs when a highly electronegative atom (N, O, or F) is bonded
to a hydrogen atom. This bond nearly strips the hydrogen atom of its electrons
leaving, essentially, a naked proton. This proton is highly attracted to the
electron pairs on nearby molecules.
Hydrogen bonding is significantly stronger than the dipole-dipole interactions
which are in turn stronger than London dispersion forces. Hydrogen bonding
exists only in molecules with an N-H, O-H, or F-H bond.
- occur between molecules that have a permanent net dipole resulting from hydrogen being covalently bonded to either fluorine, oxygen or nitrogen. For example, hydrogen bonds operate between water (H2O) molecules, ammonia (NH3) molecules, hydrogen fluoride (HF) molecules, hydrogen peroxide (H2O2) molecules, alkanols (alcohols) such as methanol (CH3OH) molecules, and between alkanoic (caboxylic) acids such as ethanoic (acetic) acid (CH3COOH) and between organic amines such as methanamine (methyl amine, CH3NH2).
- are a stronger intermolecular force than either Dispersion forces or dipole-dipole interactions since the hydrogen nucleus is extremely small and positively charged and fluorine, oxygen and nitrogen being very electronegative so that the electron on the hydrogen atom is strongly attracted to the fluorine, oxygen or nitrogen atom, leaving a highly localized positive charge on the hydrogen atom and highly negative localized charge on the fluorine, oxygen or nitrogen atom. This means the electrostatic attraction between these molecules will be greater than for the polar molecules that do not have hydrogen covalently bonded to either fluorine, oxygen or nitrogen.