Why liquids exist: the delicate balance of intermolecular forces

Science · Published 28 October 2026 Swipe more like this →

In short: If molecules only bonded when they were chemically forced to, the world would be a chaotic soup of independent gases. Instead, substances like water and alcohol stay liquid because of subtle, persistent attractions that keep molecules huddled together without locking them into permanent structures.

Intermolecular forces are the attractive interactions between neighboring molecules. While they are significantly weaker than the covalent or ionic bonds that hold atoms together within a molecule, they are the primary reason matter takes on different physical states. Proof of such forces lies in the liquefaction of gases, the substantial energy needed to vaporise common liquids, and the solid state of numerous molecular substances.

There are three primary types of these interactions. Permanent dipole-dipole forces occur in polar molecules, where atoms with different electronegativities create partial charges—such as in hydrogen chloride. Hydrogen bonds are a specialized, stronger version of this, occurring when a hydrogen atom is bonded to fluorine, oxygen, or nitrogen. Despite having only about 5% the strength of a covalent bond, hydrogen bonds are powerful enough to keep water and alcohol from becoming gases at room temperature.

Then there are van der Waals' forces, which are induced dipole-dipole interactions. These occur in all molecules, including non-polar ones like noble gases, and arise from the movement of electrons. As the number of electrons in a molecule increases, these forces strengthen. At very close range, these attractive forces are eventually overtaken by strong repulsive forces, which arise from the overlap of wave functions and the Pauli exclusion principle, preventing molecules from occupying the same space.

The relationship between these forces and energy is visualized through potential energy curves. The energy minimum for intermolecular forces is much shallower than that of chemical bonds. This distinction explains why substances can transition between solid, liquid, and gas phases relatively easily compared to the massive energy required to break the internal bonds of the molecules themselves.

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