Hydrogen Bond
A hydrogen bond is an attractive intermolecular or intramolecular force between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom. With strengths typically ranging from 5 to 40 kJ/mol, hydrogen bonds are much weaker than covalent bonds but far stronger than other van der Waals forces. They are responsible for water's unusually high boiling point, the structure of DNA base pairs, and the secondary and tertiary structures of proteins.
Hydrogen Bond Strengths in Common Systems
| System | Donor–Acceptor | Bond Energy (kJ/mol) | Significance |
|---|---|---|---|
| Water (liquid) | O–H···O | ~21 | High boiling point, surface tension |
| DNA base pairs (A–T) | N–H···O / O–H···N | ~10–25 | Genetic information storage |
| DNA base pairs (G–C) | N–H···O / N–H···N | ~25–40 | Three H-bonds, greater stability |
| Protein alpha-helix | N–H···O=C | ~8–20 | Maintains secondary structure |
| HF (liquid) | F–H···F | ~29 | Strongest per molecule |
| Ammonia (liquid) | N–H···N | ~5 | Weaker due to lower electronegativity |
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Related Terms
Van der Waals Forces
Van der Waals forces are a collective term for weak, short-range intermolecular attractions that arise from transient or permanent electric dipoles, including London dispersion forces, dipole-dipole interactions, and dipole-induced dipole interactions. Named after Dutch physicist Johannes Diderik van der Waals, these forces explain deviations of real gases from ideal behaviour and govern properties such as boiling points, surface tension, and the adhesion of geckos to surfaces. Their strength scales with molecular size and polarisability, making them significant for large molecules and noble gases.
Dipole-Dipole Interaction
Dipole-dipole interactions are intermolecular forces that occur between polar molecules, where the partially positive end (δ+) of one molecule is attracted to the partially negative end (δ–) of a neighbouring molecule. These forces are stronger than London dispersion forces but weaker than hydrogen bonds, typically ranging from 1–20 kJ/mol. They are responsible for the elevated boiling points of polar molecules such as HCl, SO₂, and acetone compared to nonpolar molecules of similar molecular weight.
Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a shared pair of electrons toward itself within a covalent bond, expressed on a dimensionless scale. The most widely used scale is the Pauling scale, where fluorine is assigned the highest value of 3.98, making it the most electronegative element, and caesium the lowest at 0.79. Electronegativity determines bond polarity, the character of chemical bonds (ionic vs. covalent), and influences molecular properties such as reactivity, acid strength, and solubility.
The term was coined by Linus Pauling in his 1931 paper and later elaborated in "The Nature of the Chemical Bond" (1939). "Hydrogen" refers to the central atom involved, and "bond" reflects its attractive nature, though it is weaker than a true covalent bond.