Van der Waals Forces

Dipole–dipole forces act between molecules that have a permanent electric dipole moment. A dipole arises when two bonded atoms differ significantly in electronegativity, creating a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative atom. For example, in hydrogen chloride (HCl), chlorine is much more electronegative than hydrogen, so the electron density is pulled toward Cl, giving it δ− and leaving H with δ+. When HCl molecules come close together, they align so that the δ+ end of one molecule points toward the δ− end of a neighbouring molecule. This electrostatic attraction lowers the potential energy of the system. Dipole–dipole forces are typically in the range of 0.5–3 kJ mol⁻¹ — much weaker than covalent bonds (~200–400 kJ mol⁻¹) but still significant enough to raise boiling points compared to similar non-polar molecules.

When a polar molecule (one with a permanent dipole) approaches a non-polar molecule, the electric field of the permanent dipole distorts the electron cloud of the non-polar molecule. This creates a temporary, induced dipole in the previously non-polar molecule. The induced dipole then interacts attractively with the permanent dipole that caused it. The strength of this interaction depends on the dipole moment of the polar molecule and the polarisability of the non-polar molecule — larger, more diffuse electron clouds are more easily distorted and thus produce stronger induced dipoles. A classic example is the interaction between a polar molecule like HCl and a non-polar molecule like Ar. These forces are generally weaker than dipole–dipole forces, typically 0.2–2 kJ mol⁻¹.

London dispersion forces, also called instantaneous dipole–induced dipole forces or simply 'London forces', are present between ALL atoms and molecules, including completely non-polar species like noble gases and diatomic halogens. At any given instant, the electrons in an atom or molecule may be unevenly distributed by chance, creating a fleeting instantaneous dipole. This temporary dipole immediately induces a complementary dipole in an adjacent molecule, and the two brief dipoles attract each other. Although each individual LDF is extremely short-lived and very weak, the cumulative effect across many molecules and many instantaneous events is significant. LDF strength increases with: (1) the number of electrons in the molecule (larger molecules are more polarisable); (2) the surface area of the molecule (more contact area means more simultaneous interactions); and (3) the shape of the molecule (long, chain-like molecules have more surface contact than compact, spherical ones). For large molecules like I₂ or long-chain alkanes, LDF can be the dominant intermolecular force, explaining why they are liquid or solid at room temperature despite being non-polar.

Polarisability is the ease with which an atom's or molecule's electron cloud can be distorted by an external electric field. It is the key factor controlling the strength of London dispersion forces. Polarisability increases as: (a) atomic/molecular size increases (electrons are further from the nucleus and less tightly held); (b) the number of electrons increases. This is why, in Group 17 (halogens), boiling point increases from F₂ (−188 °C) to Cl₂ (−34 °C) to Br₂ (+59 °C) to I₂ (+184 °C) — larger atoms have more electrons and greater polarisability, leading to stronger LDF and higher boiling points. Note that for small molecules like HF and H₂O, hydrogen bonding dominates and overrides the LDF trend.

• U_{\text{LDF}} — London dispersion potential energy (J): The potential energy of attraction due to instantaneous dipole–induced dipole interaction; negative indicates attraction.
• \alpha_1,\, \alpha_2 — Polarisability volumes of molecules 1 and 2 (m³): A measure of how easily each molecule's electron cloud is distorted; larger values mean stronger London forces.
• r — Intermolecular separation distance (m): The distance between the centres of the two interacting molecules; LDF falls off very rapidly as r⁶ in the denominator.
• I_1,\, I_2 — First ionisation energies of molecules 1 and 2 (J): The energy required to remove the outermost electron from each molecule; related to how tightly electrons are held.

Problem. Explain why pentane (C₅H₁₂, molar mass 72 g mol⁻¹) has a higher boiling point (36 °C) than neopentane (also C₅H₁₂, molar mass 72 g mol⁻¹, boiling point 9.5 °C), even though both molecules are non-polar and have identical molar masses.

Answer. Pentane has a higher boiling point than neopentane because its straight-chain shape provides greater molecular surface area, allowing stronger cumulative London dispersion forces between molecules despite both having the same molecular formula and electron count.

• Hydrogen Bonding — Hydrogen bonding is a special, stronger type of dipole–dipole interaction. Understanding Van der Waals forces helps you place hydrogen bonding in context as the strongest of the intermolecular forces and explains anomalies in properties of water and HF.
• Electronegativity and Bond Polarity — Dipole–dipole and dipole–induced dipole forces depend on the existence of permanent dipoles, which arise directly from differences in electronegativity between bonded atoms. This concept is prerequisite knowledge for Van der Waals forces.
• States of Matter and Intermolecular Forces — The physical state (solid, liquid, gas) of a molecular substance at a given temperature is determined by the balance between kinetic energy and intermolecular forces, including all three types of Van der Waals forces.
• Real Gases and the Van der Waals Equation — The van der Waals equation of state quantitatively accounts for the attractive intermolecular forces (Van der Waals forces) that cause real gases to deviate from ideal gas behaviour, especially at high pressures and low temperatures.
• Boiling Points of Molecular Substances — Boiling point trends across homologous series of organic compounds and the halogens are directly explained by the increasing strength of London dispersion forces as molecular size and electron count increase.