A hydrogen bond is an attractive interaction between a hydrogen atom attached to another atom and an electron-rich atom or group, either within the same molecule or in a different one. Familiar examples involve hydrogen attached to oxygen or nitrogen, although the phenomenon extends beyond these combinations. Hydrogen bonding helps determine the properties of water, the organization of biological macromolecules, and the assembly of molecular materials. The 2011 IUPAC recommendations define it through both the nature of the interaction and evidence that a bond has formed, rather than through a single distance or energy threshold. (publications.iupac.org)
Donors, acceptors, and notation
Hydrogen bonds are commonly represented as X–H···Y. The solid line denotes the covalent bond connecting hydrogen to X; the dots indicate its interaction with Y. The X–H group is the hydrogen-bond donor, while Y is the acceptor. In this terminology, “donor” does not necessarily mean that hydrogen is transferred: an intact X–H group can participate in hydrogen bonding without undergoing a chemical reaction. (goldbook.iupac.org)
In conventional examples, X has greater electronegativity than hydrogen, making the hydrogen end of the bond relatively electron-poor. The acceptor commonly provides a lone electron pair. Oxygen, nitrogen, and fluorine are frequent participants, but these elements do not define an exclusive list. Carbon–hydrogen groups and sulfur-containing groups can also participate under suitable conditions. Whether a particular group acts as a donor or acceptor depends on its bonding and electronic environment, not merely on the identity of an element present. (goldbook.iupac.org)
Physical origin and strength
Electrostatic attraction is an important component of hydrogen bonding. Hydrogen’s small size allows close approach between the interacting groups. However, a hydrogen bond is not adequately described in every case as attraction between fixed partial charges alone. Orbital interactions also contribute: an acceptor’s occupied orbital can interact with the antibonding orbital associated with X–H, producing partial covalent character. This electronic description connects hydrogen bonding with the broader framework of chemical bonding. (goldbook.iupac.org)
There is no universal hydrogen-bond energy. Different chemical environments produce interactions spanning a broad range, so numerical classifications into “weak,” “moderate,” and “strong” depend on context. Many familiar hydrogen bonds are weaker than ordinary covalent bonds, but unusually strong examples blur that distinction. The bifluoride ion, F–H–F⁻, is an example in which a strongly bound hydrogen lies between two fluorine atoms. Consequently, hydrogen bonding cannot be identified or excluded solely by assuming that it must always be weak. (goldbook.iupac.org)
Geometry and structural organization
Hydrogen bonds are characteristically directional. Many favor an approximately straight X–H···Y arrangement, although departures from linearity occur. Their geometry reflects both the donor bond and the distribution of electron density around the acceptor. Directionality makes hydrogen bonding particularly useful for organizing molecules into recognizable patterns rather than merely drawing them together. (goldbook.iupac.org)
An intermolecular hydrogen bond connects separate molecules; an intramolecular one connects groups within a single molecule. These arrangements can generate discrete complexes, chains, rings, or extended networks. In designed assemblies, the number and placement of complementary donor and acceptor sites influence which structures form. Multiple hydrogen bonds can therefore provide both stability and selectivity, although competing interactions may alter the resulting architecture. (goldbook.iupac.org)
Water and ice
Hydrogen bonding occurs in both liquid water and ice. In the liquid, substantial molecular disorder coexists with strong intermolecular attraction. These attractions contribute to water’s unusually high boiling point for such a small molecule, its large heat of vaporization, and its surface tension. Vaporization requires separating molecules from their cohesive liquid environment; it does not require breaking the covalent O–H bonds within each water molecule. (pubs.usgs.gov)
In ordinary ice, hydrogen bonding supports an ordered, relatively open crystal arrangement. Water molecules are farther apart on average than in liquid water, making ice less dense and allowing it to float. This statement concerns ordinary ice under familiar conditions, rather than every possible solid phase of water. The structural distinction between liquid water and ice illustrates how molecular arrangement can strongly affect bulk properties without changing chemical composition. (usgs.gov)
Biological roles
In proteins, hydrogen bonds help stabilize alpha helices and beta sheets, two major forms of secondary structure. They also contribute to protein folding and conformational stability through interactions among protein groups and with surrounding water. Experimental studies of human lysozyme show that the contribution of a particular hydrogen bond depends on its structural setting and on other changes accompanying a substitution. (cdn.rcsb.org)
In DNA, complementary donor–acceptor arrangements support pairing between adenine and thymine and between guanine and cytosine. Exposed patterns of hydrogen-bonding groups in the double helix’s grooves also provide recognition sites for proteins. Hydrogen bonding thus participates in molecular recognition as well as structural association: the location and orientation of interacting groups carry information about molecular identity. (pdb101.rcsb.org)
Molecular materials and solvent competition
In supramolecular chemistry, complementary hydrogen-bonding groups direct molecular self-assembly. They can connect building blocks into supramolecular polymers whose components associate reversibly rather than being joined exclusively by covalent bonds. Such assemblies can grow, shorten, or exchange constituents as conditions change. (doi.org)
The solvent can compete for the same donor and acceptor sites. A hydrogen-bonding pattern that associates effectively in one medium may behave differently in another. Experiments on helical assemblies in water demonstrate that a hydrophobic microenvironment can shield hydrogen-bonded groups from solvent competition. Their organization combines hydrogen bonding with the hydrophobic effect, rather than relying on either interaction in isolation. (pmc.ncbi.nlm.nih.gov)