Hydrogen bond

A hydrogen bond is a molecular attraction that is not purely electrostatic, as it involves some covalent character through charge transfer and orbital overlap. It forms when a hydrogen atom, already covalently linked to a more electronegative donor atom (like nitrogen, oxygen, or fluorine), interacts with a different electronegative atom that has a lone pair of electrons—the acceptor. This interaction is written as Dn−H···Ac, where the solid line is the covalent bond and the dots represent the hydrogen bond. Unlike simple dipole–dipole forces, hydrogen bonding relies on quantum mechanical delocalization and resonance, making it a stronger, more directed interaction.

The term "hydrogen bond" is reserved for well-defined, localized interactions with significant charge transfer, such as those in DNA base pairs or ice. A broader term, "hydrogen-bonding interactions," covers weaker, more dynamic, or delocalized cases, like those in liquid water, supramolecular assemblies (e.g., lipid membranes or protein-protein contacts), or weak C-H···O bonds. This distinction matters in structural biology, materials science, and computational chemistry, where hydrogen bonding spans a continuum from van der Waals-like attractions to nearly covalent bonds.

Hydrogen bonds can occur between separate molecules (intermolecular) or within the same molecule (intramolecular). Their strength varies widely—from about 4.2 to 167.4 kJ/mol (1 to 40 kcal/mol)—depending on geometry, environment, and the donor-acceptor pair. This places them stronger than van der Waals forces but generally weaker than covalent or ionic bonds. For example, the bifluoride ion (HF₂⁻) has an exceptionally strong hydrogen bond at 161.5 kJ/mol, while typical O−H···O bonds in water or alcohols are around 21 kJ/mol, and N−H···O bonds are about 8 kJ/mol.

The IUPAC defines the hydrogen bond as an attractive interaction between a hydrogen atom from a fragment X−H (where X is more electronegative than H) and an atom or group in the same or another molecule, provided there is evidence of bond formation. The donor hydrogen is protic and acts as a Lewis acid, while the acceptor is a Lewis base. Liquids that display hydrogen bonding, like water, are called associated liquids.

Hydrogen bonding arises from a mix of electrostatics (multipole interactions), covalency (charge transfer via orbital overlap), and dispersion forces. Weaker hydrogen bonds can involve donors like sulfur, chlorine, or even carbon (e.g., in chloroform or terminal acetylenes), especially when the carbon or a neighbor is electronegative. Though weak—around 4.2 kJ/mol—these "non-traditional" hydrogen bonds are widespread and influence many material structures. The definition has broadened over time to include these weaker interactions, as recommended by an IUPAC Task Group in 2011.

The strength of intermolecular hydrogen bonds is often measured by equilibria between donor/acceptor molecules in solution. Intramolecular hydrogen bonds are studied through conformer equilibria. Crystallography and NMR spectroscopy are key methods for identifying hydrogen bonds, even in complex molecules. A structural clue is a donor-acceptor distance smaller than the sum of their van der Waals radii, which indicates bond strength. One classification scheme groups hydrogen bonds as strong (63 to 167 kJ/mol), moderate, or weak.

Hydrogen bonding is fundamental in chemistry, biology, and materials science. It explains water's high boiling point, stabilizes protein and nucleic acid structures, and gives key properties to materials like paper, wool, and hydrogels. In biology, it drives molecular recognition, enzyme catalysis, and DNA replication; in materials science, it enables self-assembly, adhesion, and supramolecular organization.