Hydrogenbonding

Hydrogen bonding is a noncovalent interaction that occurs when a hydrogen atom covalently bonded to a highly electronegative atom (the donor) interacts with a lone pair on a second electronegative atom (the acceptor). By convention, the strongest cases involve nitrogen, oxygen, or fluorine as donors and acceptors, though other electronegative atoms such as sulfur can participate. Hydrogen bonds are a type of dipole-dipole interaction that is highly directional, favoring near-linear arrangements of donor–hydrogen–acceptor.

The strength of a hydrogen bond depends on the partners and the environment, typically ranging from about 5 to 30 kilojoules per mole, and in some cases higher in specialized systems. They are weaker than covalent bonds but stronger than most van der Waals forces, and they can form transient networks that adapt to conditions. The angle and distance between donor, hydrogen, and acceptor influence the bond energy, with near 180-degree D–H···A geometry being most favorable.

Hydrogen bonding has profound consequences for the properties of many substances. Water’s high boiling point, surface tension, and solvent abilities arise from extensive hydrogen-bond networks. In biology, hydrogen bonds stabilize DNA base pairing (A–T and G–C) and contribute to the secondary and tertiary structures of proteins. Ice expands upon freezing due to an open H-bond network. In many systems, hydrogen bonds are cooperative and can be intramolecular, influencing conformations.

Techniques such as infrared spectroscopy, NMR, and X-ray or neutron diffraction are used to study hydrogen bonds. Common donors include N–H and O–H groups; common acceptors include carbonyl, ether, amine, and hydroxyl oxygens.