Hydrogen bonds are a special type of dipole-dipole attraction between molecules, not a covalent bond to a hydrogen atom. They result from the attractive force between a hydrogen atom covalently bonded to a very electronegative atom such as a nitrogen, oxygen, or fluorine atom, and another very electronegative atom. Hydrogen bond strengths range from 4 kJ to 50 kJ per mole of hydrogen bonds. Hydrogen bonds can exist between atoms in different molecules or in parts of the same molecule. One atom of the pair, generally a fluorine, nitrogen, or oxygen atom, is covalently bonded to a hydrogen atom, whose electrons it shares unequally. Its high electron affinity causes the hydrogen to take on a slight positive charge. The other atom of the pair, also typically F, N, or O, has an unshared electron pair, which gives it a slight negative charge. Mainly through electrostatic attraction, the donor atom effectively shares its hydrogen with the acceptor atom, forming a bond.

Hydrogen bonds have important physical consequences. For example, water is liquid over a far greater range of temperatures than would be expected for a molecule of its size because of its extensive hydrogen bonding. Water is also a good solvent for ionic compounds and many others because it readily forms hydrogen bonds with the solute. Hydrogen bonding between amino acids in a linear protein molecule determines the way it folds up into its functional configuration. Hydrogen bonds between nitrogenous bases in nucleotides on the two strands of DNA also play a crucial role in the structure of DNA.