Intermolecular forces are the attractions that exist between neighbouring molecules. Unlike the strong covalent or ionic bonds that hold atoms together within a molecule, intermolecular forces are significantly weaker. Nevertheless, their cumulative effect on a bulk sample determines the physical properties of a substance — its boiling point, melting point, viscosity, surface tension, and solubility. Understanding these forces is central to explaining why, for example, water is a liquid at room temperature while hydrogen sulfide is a gas, despite both being hydrides of group 16 elements. This essay explores the principal types of intermolecular forces and assesses how each influences measurable physical properties, drawing on established chemical principles.
Types of Intermolecular Forces
Intermolecular forces can be classified into three main categories, ranked by increasing strength: London dispersion forces, permanent dipole–dipole interactions, and hydrogen bonding. All substances experience London dispersion forces, but the presence of permanent dipoles or specific hydrogen bond donors/acceptors dramatically alters physical behaviour.
London Dispersion Forces
London dispersion forces arise from temporary fluctuations in electron distribution within a molecule, creating instantaneous dipoles that induce dipoles in neighbouring molecules. Their magnitude depends on the polarisability of the electron cloud, which increases with molecular mass and surface area. For non‑polar molecules such as noble gases and alkanes, London forces are the only intermolecular attractions. As a result, the boiling points of these substances increase steadily with molar mass — for instance, methane (CH₄) boils at −161 °C, while decane (C₁₀H₂₂) boils at 174 °C (Atkins & de Paula, 2014).
Permanent Dipole–Dipole Interactions
Molecules with a permanent dipole, such as hydrogen chloride (HCl), experience dipole–dipole attractions. The partial positive end of one molecule aligns with the partial negative end of another, adding to the London force present. This interaction raises the boiling point relative to a non‑polar molecule of similar mass — compare HCl (−85 °C) with argon (−186 °C), even though both have similar relative molecular masses (36.5 vs. 40). The strength of dipole–dipole forces is influenced by the magnitude of the dipole moment and the distance between molecules.
Hydrogen Bonding
Hydrogen bonding is a particularly strong type of dipole–dipole interaction that occurs when hydrogen is covalently bonded to a highly electronegative atom (N, O, or F). The small size of hydrogen allows the lone pair of the electronegative atom on a neighbouring molecule to approach closely, resulting in an interaction that is about ten times stronger than a typical dipole–dipole force. The anomalous properties of water — its high boiling point (100 °C), high specific heat capacity, and expansion upon freezing — are direct consequences of extensive hydrogen bonding (Clayden et al., 2012).
How Intermolecular Forces Affect Physical Properties
The macroscopic properties of a substance are the net result of the sum of all intermolecular attractions. The following subsections examine key properties in turn.
Boiling and Melting Points
Boiling and melting points are the temperatures at which the kinetic energy of molecules overcomes the intermolecular forces holding them together. For a homologous series of organic compounds, boiling point increases with chain length because London forces increase with surface area. Introducing polar functional groups adds dipole–dipole and hydrogen bonding contributions, raising boiling points further. For example, ethanol (C₂H₅OH, b.p. 78 °C) boils much higher than ethane (C₂H₆, b.p. −89 °C) of comparable mass due to hydrogen bonding. Melting points are also affected, though the crystal structure and packing efficiency play additional roles. Thus, a substance with stronger intermolecular forces will have higher melting and boiling points.
Viscosity and Surface Tension
Viscosity, the resistance of a liquid to flow, depends on the strength of intermolecular forces. Liquids with extensive hydrogen bonding, such as glycerol, are highly viscous because molecules cannot slide past one another easily. Similarly, surface tension — the energy required to increase the surface area of a liquid — is higher in liquids with strong intermolecular attractions. Water, with its network of hydrogen bonds, has a surface tension of 72.8 mN m⁻¹ at 20 °C, whereas non‑polar hexane has a surface tension of only 18.4 mN m⁻¹ (Atkins & de Paula, 2014). These differences explain why water droplets bead up on waxed surfaces while hexane spreads out.
Solubility
The principle “like dissolves like” is a direct consequence of intermolecular forces. Polar and ionic solutes dissolve in polar solvents because the solvent molecules can overcome the solute–solute attractions through dipole–dipole or ion–dipole interactions. Non‑polar solutes dissolve in non‑polar solvents through London dispersion forces. The role of hydrogen bonding is exemplified by the miscibility of ethanol with water — both are capable of hydrogen bonding. However, longer‑chain alcohols become increasingly immiscible as the non‑polar hydrocarbon chain dominates. This contrast governs many practical applications, from pharmaceutical formulation to the design of industrial solvents.
Vapour Pressure and Enthalpy of Vaporisation
The strength of intermolecular forces inversely correlates with vapour pressure — the pressure exerted by a vapour in equilibrium with its liquid. Stronger forces reduce the tendency of molecules to escape into the gas phase, so a liquid with strong hydrogen bonding has a lower vapour pressure at a given temperature. For example, water has a vapour pressure of only 2.34 kPa at 20 °C, while diethyl ether (non‑hydrogen bonding) has a vapour pressure of 58.9 kPa. Consequently, the enthalpy of vaporisation (ΔHᵥₐₚ) is higher for water (40.7 kJ mol⁻¹) than for non‑polar liquids of similar molar mass (Atkins & de Paula, 2014). This explains why water is an excellent coolant — it absorbs much heat without a large temperature rise.
Conclusion
Intermolecular forces are the fundamental interactions that dictate the physical state and behaviour of all substances. London dispersion forces, dipole–dipole interactions, and hydrogen bonding each contribute to a network of attractions that determines boiling and melting points, viscosity, surface tension, solubility, vapour pressure, and enthalpy changes. The relative importance of each force depends on the molecular structure and the functional groups present. Mastery of this topic is essential for any advanced chemistry student because it underpins predictions about chemical behaviour — from the design of pharmaceuticals to the engineering of new materials.
For students seeking to strengthen their knowledge of intermolecular forces in the context of A Level Chemistry, revision materials such as A Levels Economics Revision Notes and Essays provide structured guidance. Although this resource is titled for economics, the essay‑writing techniques it teaches are directly transferable to rigorous chemistry essays.
Furthermore, the principles discussed here link directly to other key A Level topics. Studying equilibrium principles reveals how intermolecular forces affect the position of equilibrium in solution reactions. The reactivity of organic molecules is also heavily influenced by the same forces that determine physical properties. Understanding these connections enables students to write essays that demonstrate coherent chemical reasoning.
Frequently Asked Questions
What are intermolecular forces?
Intermolecular forces are attractive forces that exist between separate molecules. They are weaker than covalent or ionic bonds but collectively determine physical properties such as boiling point, viscosity, and solubility.
How do intermolecular forces affect boiling point?
Stronger intermolecular forces require more energy to overcome, raising the boiling point. For instance, water’s high boiling point relative to its molecular mass is due to hydrogen bonding.
Why is water a liquid at room temperature but hydrogen sulfide is a gas?
Water forms extensive hydrogen bonds because oxygen is highly electronegative and small. Hydrogen sulfide has weaker dipole–dipole forces and no hydrogen bonding, so it remains a gas despite similar molar mass.
What is the difference between intramolecular and intermolecular forces?
Intramolecular forces (covalent, ionic) hold atoms together within a molecule, while intermolecular forces attract different molecules to each other. The former are much stronger.
How does surface tension relate to intermolecular forces?
Surface tension arises from the net inward pull on molecules at the surface, caused by unbalanced intermolecular attractions. Liquids with stronger forces have higher surface tension.
References
Atkins, P. W., & de Paula, J. (2014). Atkins’ Physical Chemistry (10th ed.). Oxford University Press.
Clayden, J., Greeves, N., & Warren, S. (2012). Organic Chemistry (2nd ed.). Oxford University Press.
Housecroft, C. E., & Sharpe, A. G. (2018). Inorganic Chemistry (5th ed.). Pearson Education.
