Hydrogen bond

In chemistry, a hydrogen bond is a specific type of molecular interaction that exhibits partial covalent character and cannot be described as a purely electrostatic force. It occurs when a hydrogen atom, covalently bonded to a more electronegative donor atom or group, interacts with another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor. Unlike simple dipole–dipole interactions, hydrogen bonding arises from charge transfer, orbital interactions, and quantum mechanical delocalization, making it a resonance-assisted interaction rather than a mere electrostatic attraction.
The general notation for hydrogen bonding is Dn−H···Ac, where the solid line represents a polar covalent bond, and the three dots indicate the hydrogen bond. Hydrogen bond donors have a protic hydrogen attached to an electonegative atom such as nitrogen, oxygen, and fluorine. Hydrogen bond acceptors have a lone pair of electrons, such as amines, carboxylates, and water.
The term "hydrogen bond" is generally used for well-defined, localized interactions with significant charge transfer and orbital overlap, such as those in DNA base pairing or ice. In contrast, "hydrogen-bonding interactions" is a broader term used when the interaction is weaker, more dynamic, or delocalized, such as in liquid water, supramolecular assemblies, or weak C-H···O interactions. This distinction is particularly relevant in structural biology, materials science, and computational chemistry, where hydrogen bonding spans a continuum from weak van der Waals-like interactions to nearly covalent bonding.
Hydrogen bonding can occur between separate molecules or within different parts of the same molecule. Its strength varies considerably, depending on geometry, environment, and the donor-acceptor pair, typically ranging from. This places hydrogen bonds stronger than van der Waals interactions but generally weaker than covalent or ionic bonds.
Hydrogen bonding plays a fundamental role in chemistry, biology, and materials science. It is responsible for the anomalously high boiling point of water, the stabilization of protein and nucleic acid structures, and key properties of materials like paper, wool, and hydrogels. In biological systems, hydrogen bonds mediate molecular recognition, enzyme catalysis, and DNA replication, while in materials science, they contribute to self-assembly, adhesion, and supramolecular organization.

Bonding

Definitions and general characteristics

In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named the proton acceptor, whereas the one covalently bound to the hydrogen is named the proton donor. This nomenclature is recommended by the IUPAC. The hydrogen of the donor is protic and therefore can act as a Lewis acid and the acceptor is the Lewis base. Hydrogen bonds are represented as system, where the dots represent the hydrogen bond. Liquids that display hydrogen bonding are called associated liquids.
Hydrogen bonds arise from a combination of electrostatics, covalency, and dispersion.
In weaker hydrogen bonds, hydrogen atoms tend to bond to elements such as sulfur or chlorine ; even carbon can serve as a donor, particularly when the carbon or one of its neighbors is electronegative. Gradually, it was recognized that there are many examples of weaker hydrogen bonding involving donor other than N, O, or F and/or acceptor Ac with electronegativity approaching that of hydrogen. Although weak, "non-traditional" hydrogen bonding interactions are ubiquitous and influence structures of many kinds of materials.
The definition of hydrogen bonding has gradually broadened over time to include these weaker attractive interactions. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. This definition specifies:

Bond strength

Hydrogen bonds can vary in strength from weak to strong. Typical enthalpies in vapor include:

, illustrated uniquely by HF₂^-
, illustrated water-ammonia
, illustrated water-water, alcohol-alcohol
, illustrated by ammonia-ammonia
, illustrated water-amide

The strength of intermolecular hydrogen bonds is most often evaluated by measurements of equilibria between molecules containing donor and/or acceptor units, most often in solution. The strength of intramolecular hydrogen bonds can be studied with equilibria between conformers with and without hydrogen bonds. The most important method for the identification of hydrogen bonds also in complicated molecules is crystallography, sometimes also NMR-spectroscopy. Structural details, in particular distances between donor and acceptor which are smaller than the sum of the van der Waals radii can be taken as indication of the hydrogen bond strength. One scheme gives the following somewhat arbitrary classification: those that are,, and are considered strong, moderate, and weak, respectively.
Hydrogen bonds involving C−H bonds are both very rare and weak.

Resonance assisted hydrogen bond

The resonance assisted hydrogen bond is a strong type of hydrogen bond. It is characterized by the π-delocalization that involves the hydrogen and cannot be properly described by the electrostatic model alone. This description of the hydrogen bond has been proposed to describe unusually short distances generally observed between or.

Structural details

The distance is typically ≈110 pm, whereas the distance is ≈160 to 200 pm. The typical length of a hydrogen bond in water is 197 pm. The ideal bond angle depends on the nature of the hydrogen bond donor. The following hydrogen bond angles between a hydrofluoric acid donor and various acceptors have been determined experimentally:

	VSEPR geometry	Angle
	C≡N···H angle: linear	180°
	C=O···H angle: trigonal planar	120°
	H-O···H angle: pyramidal	46°
	H-S···H angle: pyramidal	89°
	S=O···H angle: trigonal	145°

Spectroscopy

Strong hydrogen bonds are revealed by downfield shifts in the ¹H NMR spectrum. For example, the acidic proton in the enol tautomer of acetylacetone appears at 15.5, which is about 10 ppm downfield of a conventional alcohol.
In the IR spectrum, hydrogen bonding shifts the stretching frequency to lower energy. This shift reflects a weakening of the bond. Certain hydrogen bonds - improper hydrogen bonds - show a blue shift of the stretching frequency and a decrease in the bond length. H-bonds can also be measured by IR vibrational mode shifts of the acceptor. The amide I mode of backbone carbonyls in α-helices shifts to lower frequencies when they form H-bonds with side-chain hydroxyl groups. The dynamics of hydrogen bond structures in water can be probed by this OH stretching vibration. In the hydrogen bonding network in protic organic ionic plastic crystals, which are a type of phase change material exhibiting solid-solid phase transitions prior to melting, variable-temperature infrared spectroscopy can reveal the temperature dependence of hydrogen bonds and the dynamics of both the anions and the cations. The sudden weakening of hydrogen bonds during the solid-solid phase transition seems to be coupled with the onset of orientational or rotational disorder of the ions.

Theoretical considerations

Hydrogen bonding is of persistent theoretical interest. According to a modern description integrates both the intermolecular O:H lone pair ":" nonbond and the intramolecular polar-covalent bond associated with repulsive coupling.
Quantum chemical calculations of the relevant interresidue potential constants revealed large differences between individual H bonds of the same type. For example, the central interresidue hydrogen bond between guanine and cytosine is much stronger in comparison to the bond between the adenine-thymine pair.
Theoretically, the bond strength of the hydrogen bonds can be assessed using NCI index, non-covalent interactions index, which allows a visualization of these non-covalent interactions, as its name indicates, using the electron density of the system.
Interpretations of the anisotropies in the Compton profile of ordinary ice claim that the hydrogen bond is partly covalent. However, this interpretation was challenged and subsequently clarified.
Most generally, the hydrogen bond can be viewed as a metric-dependent electrostatic scalar field between two or more intermolecular bonds. This is slightly different from the intramolecular bound states of, for example, covalent or ionic bonds. However, hydrogen bonding is generally still a bound state phenomenon, since the interaction energy has a net negative sum. The initial theory of hydrogen bonding proposed by Linus Pauling suggested that the hydrogen bonds had a partial covalent nature. This interpretation remained controversial until NMR techniques demonstrated information transfer between hydrogen-bonded nuclei, a feat that would only be possible if the hydrogen bond contained some covalent character.

History

The concept of hydrogen bonding once was challenging. Linus Pauling credits T. S. Moore and T. F. Winmill with the first mention of the hydrogen bond, in 1912. Moore and Winmill used the hydrogen bond to account for the fact that trimethylammonium hydroxide is a weaker base than tetramethylammonium hydroxide. The description of hydrogen bonding in its better-known setting, water, came some years later, in 1920, from Latimer and Rodebush. In that paper, Latimer and Rodebush cited the work of a fellow scientist at their laboratory, Maurice Loyal Huggins, saying, "Mr. Huggins of this laboratory in some work as yet unpublished, has used the idea of a hydrogen kernel held between two atoms as a theory in regard to certain organic compounds."