Intermolecular forces

Intermolecular forces (IMFs), sometimes referred to as secondary forces, are the interactions that occur between molecules or between molecules and neighbouring particles such as atoms or ions. These forces arise from electromagnetic attractions or repulsions and are significantly weaker than the intramolecular forces—such as covalent, ionic, or metallic bonds—that hold atoms together within a molecule. Although comparatively weak, intermolecular forces are fundamental to understanding the physical properties of matter, governing phenomena such as boiling and melting points, viscosity, solubility, and the behaviour of biological macromolecules.

Nature and Historical Development

The concept of microscopic forces emerged gradually through the work of early scientists. Alexis Clairaut’s Théorie de la figure de la Terre (1743) marks one of the earliest discussions, followed by significant contributions from Pierre-Simon Laplace, Carl Friedrich Gauss, James Clerk Maxwell, Ludwig Boltzmann, and Linus Pauling. With the rise of molecular physics and quantum chemistry in the twentieth century, the understanding of intermolecular forces became firmly linked to potential energy functions and statistical mechanics. These forces can be analysed through macroscopic properties such as viscosities and pressure–volume–temperature measurements, and modelled using pair potentials including the Lennard–Jones, Buckingham, and Mie potentials.
Intermolecular forces operate between any combination of molecules, atoms, ions, or molecular ions without generating new covalent or ionic structures. Although they do not fundamentally alter the electron configuration of interacting species, IMFs play major roles in catalysis, biomolecular recognition, and enzyme–substrate binding. In biological systems, numerous weak interactions with precise spatial alignment can lead to substantial changes in energy states and facilitate subsequent covalent reactions.

Hydrogen Bonding

Hydrogen bonding is a particularly important type of intermolecular attraction. It occurs when a hydrogen atom covalently bonded to a highly electronegative atom—typically nitrogen, oxygen, or fluorine—is attracted to another electronegative atom bearing a lone pair of electrons. Hydrogen bonds have an electrostatic character but also display partial covalency, reflected in their directionality, shorter-than-expected interatomic distances, and limited coordination patterns.
In water, each molecule can form up to four hydrogen bonds: two via its lone pairs and two through its hydrogen atoms. This extensive hydrogen-bond network is responsible for water’s high boiling point relative to other hydrides of the chalcogens. Hydrogen bonds are central to the structural organisation of proteins (in α-helices and β-sheets), nucleic acids (in base pairing), and numerous natural and synthetic polymers.

Salt Bridges

Salt bridges are noncovalent attractions between positively and negatively charged sites, commonly described as ion pairing. Driven largely by electrostatic forces, their behaviour in aqueous environments is influenced by entropy and often shows endothermic characteristics. In crystalline solids, salt bridges display characteristic interionic distances governed by van der Waals radii, and they tend not to be directional.
In solution, the free energy associated with formation of a simple 1:1 ion pair is typically around 5–6 kJ mol⁻¹, largely independent of ion size or polarizability. The strength of these interactions increases with ionic charge and depends on ionic strength according to the Debye–Hückel treatment. Salt bridges are significant in protein stability, molecular recognition, and supramolecular chemistry.

Dipole–Dipole and Related Interactions

Dipole–dipole interactions, or Keesom forces, arise between molecules possessing permanent dipole moments. These interactions are stronger than London dispersion forces but weaker than ion–ion attractions. Molecules align such that opposite partial charges attract, reducing potential energy. Examples include interactions between hydrogen chloride molecules.
Symmetric molecules may contain polar bonds yet exhibit no net dipole moment; in such cases, dipole–dipole interactions cancel out. Carbon dioxide and tetrachloromethane exemplify this behaviour. Since atoms rarely possess permanent dipoles, atom–atom dipole interactions are usually negligible.
Dipole–dipole forces are part of the broader family of van der Waals interactions and exhibit characteristic temperature dependence. Their interaction energy decreases with the inverse sixth power of the distance.

Ion–Dipole and Ion-Induced Dipole Forces

Ion–dipole forces occur between ions and polar molecules. These interactions are stronger than hydrogen bonds due to the full ionic charges involved. They play crucial roles in solvation processes, particularly the hydration of ions in water. When water molecules orient around ions, the energy released is termed hydration enthalpy, an important factor in explaining the stability of ions such as Cu²⁺ in aqueous media.
Ion-induced dipole interactions involve an ion and a nonpolar molecule. The ion’s electric field distorts the electron cloud of the nonpolar species, creating a temporary dipole and allowing attraction. These interactions underpin many solubility and complexation processes involving weakly polar or nonpolar molecules near charged centres.

Van der Waals Forces

Van der Waals forces encompass several types of interactions between uncharged particles. These forces contribute to the cohesion of condensed phases, the physisorption of gases, and the general attraction between macroscopic bodies.

• Keesom interactions: Permanent dipole–dipole attractions, averaged over molecular rotations.
• Debye interactions: Attractions between a permanent dipole and an induced dipole.
• London dispersion forces: Interactions arising from instantaneous fluctuations in electron distribution, producing temporary dipoles even in nonpolar molecules.

Dispersion forces are universal and often dominate intermolecular attraction, particularly in large, nonpolar molecules. They depend on molecular size, polarizability, and intermolecular distance, decreasing with the inverse sixth power of separation.
Keesom interactions, in contrast, require permanent dipole moments and are sensitive to temperature. Their average interaction energy is expressed as an inverse sixth-power function of distance, similar to dispersion forces, but includes explicit dependence on rotational averaging and thermal conditions.

Importance in Chemistry and Biology

Intermolecular forces determine the physical behaviour of materials, influencing boiling and melting points, solubility, crystal structures, and viscosity. In molecular biology, IMFs govern enzyme–substrate recognition, protein folding, membrane integrity, and nucleic acid pairing. Weak individually, but powerful collectively, these forces underpin mechanisms of catalysis, molecular self-assembly, and supramolecular organisation.