Hydrogen bond

A hydrogen bond is a molecular interaction that involves partial covalent character, not just electrostatic attraction. It forms when a hydrogen atom, already covalently linked to a more electronegative donor atom or group, interacts with another electronegative atom that has a lone pair of electrons—the acceptor. This interaction arises from charge transfer, orbital overlap, and quantum mechanical delocalization, making it a resonance-assisted phenomenon rather than a simple dipole–dipole force. The standard notation is Dn−H···Ac, where the solid line marks the polar covalent bond and the dots represent the hydrogen bond itself.

Donors are protic hydrogens attached to electronegative atoms like nitrogen, oxygen, or fluorine. Acceptors possess lone pairs, such as the nitrogen in amines and amides or the oxygen in carboxylates and water. The term "hydrogen bond" is reserved for well-defined, localized interactions with significant charge transfer and orbital overlap—for example, in DNA base pairing or ice. "Hydrogen-bonding interactions" is a broader term for weaker, more dynamic, or delocalized cases, such as in liquid water, supramolecular assemblies like lipid membranes and protein-protein contacts, or weak C-H···O interactions. This distinction matters in structural biology, materials science, and computational chemistry, where hydrogen bonding spans a continuum from van der Waals-like forces to nearly covalent bonds.

Hydrogen bonds can occur between separate molecules (intermolecular) or within different parts of the same molecule (intramolecular). Their strength varies widely with geometry, environment, and the donor-acceptor pair, typically ranging from 4.2 to 167.4 kJ/mol (1 to 40 kcal/mol)—stronger than van der Waals interactions but generally weaker than covalent or ionic bonds.

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

The electronegative atom not covalently attached to hydrogen is called the proton acceptor, while the one covalently bound is the proton donor—nomenclature recommended by IUPAC. The donor’s hydrogen is protic and acts as a Lewis acid; the acceptor is a Lewis base. Hydrogen bonds are represented as H···Y systems, and liquids that display them (like water) are termed associated liquids. These bonds arise from electrostatics (multipole-multipole and multipole-induced multipole interactions), covalency (charge transfer via orbital overlap), and dispersion (London forces).

Weaker hydrogen bonds can involve donors like sulfur, chlorine, or even carbon—especially when the carbon or a neighbor is electronegative (e.g., in chloroform, aldehydes, or terminal acetylenes). Over time, it became clear that many examples exist where donors are not limited to N, O, or F, or where acceptors have electronegativity close to hydrogen’s. Though weak (around 4.2 kJ/mol or 1 kcal/mol), these "non-traditional" hydrogen bonds are widespread and influence many material structures. The definition broadened accordingly, and in 2011 an IUPAC Task Group recommended a modern evidence-based definition: the hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or fragment X−H (where X is more electronegative than H) and an atom or group in the same or another molecule, where there is evidence of bond formation.

Hydrogen bond strength ranges from weak (1–2 kJ/mol) to strong (161.5 kJ/mol in the bifluoride ion, HF₂⁻). Typical vapor-phase enthalpies include: F−H···F⁻ (161.5 kJ/mol), O−H···N (29 kJ/mol), O−H···O (21 kJ/mol), N−H···N (13 kJ/mol), N−H···O (8 kJ/mol), and OH₃⁺···OH₂ (18 kJ/mol). Intermolecular hydrogen bond strength is often measured by equilibria between donor/acceptor molecules in solution; intramolecular strength is studied via conformer equilibria. The most important identification method, even in complex molecules, is crystallography, sometimes NMR spectroscopy. Structural details—especially donor-acceptor distances smaller than the sum of van der Waals radii—indicate bond strength. One classification scheme groups hydrogen bonds as those from 63 to 167 kJ/mol (15 to 40 kcal/mol).