Hydrogen Bonding

hydrogen bonding

Interaction involving a hydrogen atom located between a pair of other atoms having a high affinity for electrons; such a bond is weaker than an ionic bond or covalent bond but stronger than van der Waals forces. Hydrogen bonds can exist between atoms in different molecules or in parts of the same molecule. One atom of the pair (the donor), generally a fluorine, nitrogen, or oxygen atom, is covalently bonded to a hydrogen atom (−FH, −NH, or −OH), whose electrons it shares unequally; its high electron affinity causes the hydrogen to take on a slight positive charge. The other atom of the pair, also typically F, N, or O, has an unshared electron pair, which gives it a slight negative charge. Mainly through electrostatic attraction, the donor atom effectively shares its hydrogen with the acceptor atom, forming a bond. Because of its extensive hydrogen bonding, water (H₂O) is liquid over a far greater range of temperatures that would be expected for a molecule of its size. Water is also a good solvent for ionic compounds and many others because it readily forms hydrogen bonds with the solute. Hydrogen bonding between amino acids in a linear protein molecule determines the way it folds up into its functional configuration. Hydrogen bonds between nitrogenous bases in nucleotides on the two strands of DNA (guanine pairs with cytosine, adenine with thymine) give rise to the double-helix structure that is crucial to the transmission of genetic information.

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Hydrogen Bonding 

a form of chemical interaction of atoms in molecules characterized by the essential involvement of a hydrogen atom (H) already linked by a covalent bond to another atom (A). The A—H group acts as a proton donor (electron acceptor), while another group (or atom) B acts as an electron donor (proton acceptor). In other words, the A—H group behaves like an acid and group B like a base. Unlike the ordinary valence bond, which is denoted by a dash, the hydrogen bond is shown by dots, that is, A—H ··· B (in the limiting case of a symmetrical hydrogen bond, for example, in the acid potassium bifluoride K⁺(F ··· H ··· F)⁻, the difference between the two bonds disappears).

Groups able to form a hydrogen bond are A—H, where A is one of the atoms O, N, F, Cl, Br, and to a lesser extent C and S. The electron donor center B can be the same atoms O, N, S in various functional groups, such anions as F⁻ and Cl⁻, and to a lesser degree aromatic rings and multiple bonds. If A—H and B belong to separate (different or identical) molecules, the hydrogen bond is called intermolecular, and if they are located in different parts of the same molecule, it is called intramolecular.

Hydrogen bonding differs from the van der Waals forces of mutual attraction common to all molecules in being directional and having a capacity for saturation, that is, by having the characteristics of ordinary (valence) chemical bonds. Unlike what was previously thought, hydrogen bonding does not reduce to electrostatic attraction between polar groups A—H and B but is considered as a donor-acceptor chemical bond. In respect to its binding energy, usually 3-8 kilocalories per mole, the hydrogen bond is intermediate between van der Waals interactions (equal to fractions of a kilocalorie per mole) and typical chemical bonds (tens of kilocalories per mole; 1 kilocalorie ≈ 4.19 × 10³ joules).

The most widespread intermolecular bonding is hydrogen bonding. It leads to an association of identical or different molecules into various aggregate-complexes via hydrogen bonds, or to H-complexes, which under ordinary conditions undergo rapid equilibration. By this process we obtain both binary complexes (acid-base and cyclic dimers) and large structures (chains, rings, spirals, and plane and three-dimensional networks of linked molecules). The presence of such hydrogen bonds determines the properties of various solutions and liquids (primarily water and aqueous solutions, a series of industrial polymers—capron, nylon, and so forth) as well as the crystalline structures of many molecular crystals and crystal hydrates of inorganic compounds, among them, of course, ice. In exactly the same way hydrogen bonding substantially determines the structures of proteins, nucleic acids, and other biologically important compounds, and hence plays a major part in the chemistry of all vital processes. Because of the universal occurrence of hydrogen bonds, their role is also substantial in many other areas of chemistry and technology (distillation, extraction, adsorption, acid-base equilibrium processes, chromatography, catalysis, and so forth).

Since the formation of hydrogen bonds specifically alters the properties of the A—H and B groups, it is also reflected in the molecular properties, and this is shown, in part, in the vibrational spectra and proton magnetic resonance spectra. Hence spectroscopy, especially infrared spectroscopy, is an important method for studying hydrogen bonds and the processes depending on them.

REFERENCES

Pimentel, G., and A. McClellan. Vodorodnaia sviaz’. Moscow, 1964. (Translated from English.)
Vodorodnaia sviaz’: sb. st. Moscow, 1964.
Pauling, L. The Chemical Bond. New York, 1967.

A. V. IOGANSEN