Water boils at 100 °C. Methane boils at −161 °C. The two molecules have similar molecular weights — water is 18, methane is 16 — yet their boiling points differ by more than 250 degrees. The reason is not the chemistry within molecules (the covalent bonds holding atoms together) but the chemistry between molecules: the intermolecular forces that determine how strongly molecules stick to one another. Intermolecular forces are typically one to two orders of magnitude weaker than covalent bonds, but they control nearly every macroscopic property of substances — phase transitions, solubility, surface tension, viscosity, the structure of biological macromolecules, DNA's double helix; without them there would be no liquids, no condensed matter, and no life.
Intermolecular forces are attractive interactions between separate molecules arising from electrostatic interactions between partial charges and induced dipoles, falling into a few categories of decreasing strength. Hydrogen bonds (5–30 kJ/mol) are strong dipole-dipole interactions in which a hydrogen covalently bound to N, O, or F is attracted to a lone pair on another N, O, or F — water's anomalously high boiling point, surface tension, and heat capacity all derive from extensive hydrogen bonding (each water can form up to four), DNA is held in its double helix by hydrogen bonds between complementary bases (A–T uses two, G–C uses three), and protein secondary structure (α-helices, β-sheets) is stabilized by hydrogen bonds along the peptide backbone. Ion-dipole interactions (40–600 kJ/mol) drive dissolution of ionic compounds in water, dipole-dipole interactions (5–25 kJ/mol) act between molecules with permanent dipoles, and London dispersion forces (0.05–40 kJ/mol) are the universal attraction between all atoms and molecules arising from transient electron fluctuations — individually weak but cumulative, scaling with molecular size (methane is gas, hexane is liquid, octadecane is wax). The hydrophobic effect — the apparent attraction between nonpolar molecules in water — is driven by water's preference for hydrogen-bonding to itself: nonpolar surfaces disrupt the network, and minimizing nonpolar-water contact frees water to bond more, driving protein folding, lipid-bilayer formation, and many drug-target interactions. Macroscopic consequences run through every property of matter — boiling and melting points, viscosity, surface tension, solubility (like dissolves like).
Drug design is centrally about optimizing intermolecular interactions between drug and target — typically a protein binding pocket — with hydrogen-bond donors and acceptors tracked during medicinal chemistry, and Lipinski's Rule of Five (1997) explicitly limits them. Molecular docking software (AutoDock, Glide, Vina) computes binding energies of candidate drugs to target proteins, and AlphaFold-derived protein structures now provide high-quality starting points across nearly all human proteins. In materials science, liquid crystals (LCD displays) are organized by intermolecular forces, adhesives (gecko-inspired, super glue) exploit van der Waals and capillary forces, and self-assembled monolayers in nanotechnology are organized similarly. The forces Linus Pauling first systematized in The Nature of the Chemical Bond (1939) remain the foundation of structural biology, materials chemistry, and pharmaceutical design.