Hydrogen bonding is a relatively strong intermolecular attraction that occurs when a hydrogen atom covalently bonded to a highly electronegative atom is attracted to a lone pair of electrons on another electronegative atom. In ordinary molecular chemistry, the atoms most strongly associated with hydrogen bonding are nitrogen, oxygen and fluorine because their high electronegativities make the hydrogen in bonds such as O-H, N-H and F-H substantially electropositive.
The hydrogen bond is therefore different from an ordinary covalent bond within a molecule. A covalent bond involves sharing electrons between atoms, whereas a hydrogen bond is an attractive interaction between a hydrogen donor and an electron rich acceptor. Hydrogen bonding can occur between separate molecules, called intermolecular hydrogen bonding, or between different parts of the same molecule, called intramolecular hydrogen bonding. It is responsible for many important physical and chemical properties of water, alcohols, amines, proteins, nucleic acids and numerous other substances.
The molecular basis of hydrogen bonding arises from unequal electron distribution. When hydrogen is covalently bonded to oxygen, nitrogen or fluorine, the electronegative atom attracts the shared electrons more strongly. The electronegative atom consequently develops a partial negative charge, while the bonded hydrogen develops a partial positive charge. A nearby electronegative atom possessing an available lone pair can then attract the partially positive hydrogen.
The interaction is directional because the strongest hydrogen bonding generally occurs when the donor hydrogen, acceptor atom and relevant electron density are arranged approximately along a favourable geometric direction. This directional character distinguishes hydrogen bonding from a simple nonspecific attraction between opposite charges. The strength of an individual hydrogen bond depends on the identities of the donor and acceptor, their electronic environments, distance, angle and surrounding molecular environment. Hydrogen bonds are weaker than typical covalent bonds but can be strong enough collectively to dominate the behaviour of molecular liquids and biological macromolecules.
Hydrogen bonding is particularly important in water. Each water molecule contains two O-H bonds in which oxygen is more electronegative than hydrogen, giving the molecule partial charges. Oxygen also has two lone pairs capable of accepting hydrogen bonds. Consequently, water molecules form an extensive and constantly changing network of hydrogen bonds.
At any instant, a water molecule can interact with several neighbouring molecules, although the exact number and geometry continually fluctuate because thermal motion continually breaks and reforms individual hydrogen bonds. This network contributes to water's unusually high boiling point compared with many molecules of similar molecular mass, its high specific heat capacity, substantial heat of vaporisation and unusual surface tension. Hydrogen bonding also contributes to the density anomaly of water, because the relatively open structure associated with hydrogen bonding in ice differs from the more compact average molecular arrangements found in liquid water.
Hydrogen bonding has major effects on phase changes because energy is required to disrupt intermolecular attractions. A liquid containing strong hydrogen bonding can require considerable energy to convert into a gas because molecules must separate sufficiently to escape the condensed phase. This helps explain why substances such as water and ethanol have much higher boiling points than might be expected from molecular mass alone. Hydrogen bonding also affects melting points, viscosity, surface tension, evaporation, solubility and molecular association.
The strength and prevalence of hydrogen bonding can change with temperature, pressure and molecular environment. In a gas, hydrogen bonded clusters may still occur under suitable conditions, but their average abundance is much lower than in the liquid. In a solid, molecules can become organised into relatively persistent hydrogen bonded structures. The macroscopic properties of a substance therefore emerge from the collective behaviour of enormous numbers of microscopic interactions.
Hydrogen bonding is fundamental to biological structure and function. In proteins, hydrogen bonds help stabilise secondary structures such as alpha helices and beta sheets by linking appropriate peptide groups along the polypeptide chain. They also contribute to tertiary and quaternary structure through interactions involving side chains and surrounding molecules. In nucleic acids, hydrogen bonding between complementary bases contributes to the organisation of DNA into its double helical structure and helps RNA adopt functional folded structures.
Hydrogen bonds are not the only forces involved in biological structure. Covalent bonds, ionic interactions, van der Waals interactions, hydrophobic effects and other contributions are also important. Hydrogen bonding is especially useful biologically because it can be sufficiently strong to provide structural organisation while remaining reversible enough to permit molecular recognition, conformational changes, binding and dynamic molecular processes.
Hydrogen bonds can be intramolecular or intermolecular. Intramolecular hydrogen bonding occurs when donor and acceptor groups within the same molecule are positioned appropriately to interact. This can create a favourable ring like arrangement and can influence molecular shape, polarity, acidity and solubility. Intermolecular hydrogen bonding occurs between separate molecules and can produce chains, networks, dimers or larger molecular assemblies. A molecule can simultaneously act as a hydrogen bond donor and acceptor, allowing extensive networks to develop.
Carboxylic acids, for example, can form strongly associated pairs through hydrogen bonding, while alcohol molecules can form extended intermolecular networks. The presence or absence of hydrogen bonding can substantially alter boiling point and solubility. Molecular geometry is therefore critical. Merely containing hydrogen and oxygen does not guarantee strong hydrogen bonding because the relevant atoms must have appropriate bonding and electron arrangements.
Hydrogen bonding also influences acid base chemistry, molecular recognition and solvation. Water molecules can hydrogen bond with ions and polar solutes, helping stabilise dissolved species. Hydrogen bonding between solvent and solute can therefore influence whether a substance dissolves readily. In organic chemistry, hydrogen bonding can affect reaction mechanisms because reactants, intermediates and transition states may be stabilised or destabilised by hydrogen bond interactions. It can also influence proton transfer, where a hydrogen atom effectively moves between interacting molecular sites.
Hydrogen bonds can vary substantially in strength, and modern chemistry recognises a broader range of hydrogen bonding interactions than the simplest textbook definition suggests. Some systems involve unusual donor and acceptor atoms or electronically unusual environments. Nevertheless, the central idea remains an attractive interaction involving hydrogen attached to an electronegative donor and electron density associated with an acceptor.
Hydrogen bonding is therefore best understood as an important directional intermolecular or intramolecular interaction arising from molecular polarity and electron distribution. It is stronger and more structured than many ordinary van der Waals interactions but much weaker than the covalent bonds that construct the molecules themselves. Its effects are visible on many scales, from the boiling behaviour of water to the three dimensional structures of proteins and nucleic acids.
The concept also explains why chemically similar substances can have very different physical properties when their hydrogen bonding capacities differ. Hydrogen bonding is dynamic rather than permanently fixed in liquids and many biological systems. Bonds continuously form, distort and break as molecules move. Its importance comes not simply from the strength of an individual interaction but from the large number of hydrogen bonds that can operate cooperatively within molecular networks.
