Intermolecular forces (IMFs) are the attractive or repulsive forces that arise between molecules, separate from the chemical bonds that hold atoms together within a molecule. These relatively weaker forces govern how individual molecules interact, holding them in close proximity. The collective strength of these attractions determines the physical state of matter—whether a substance is a solid, liquid, or gas at a given temperature. Understanding these four types of forces explains many observable properties.
London Dispersion Forces
London Dispersion Forces (LDFs) are the weakest intermolecular attractions, yet they are present in every atom and molecule, including nonpolar substances like noble gases or hydrocarbons. This force results from the constant, random movement of electrons, which can create a fleeting, momentary charge separation known as a temporary dipole. This temporary dipole in one molecule induces a corresponding, short-lived dipole in a neighbor, resulting in a weak, attractive force. The strength of LDFs increases significantly with the size and mass of the molecule because molecules with more electrons have a larger, more easily distorted electron cloud. This ease of distortion is called polarizability, and higher polarizability leads to stronger LDFs.
Dipole-Dipole Interactions
Dipole-Dipole interactions occur exclusively between polar molecules, which have a permanent separation of charge. A permanent dipole is created by the uneven sharing of electrons between atoms due to differences in electronegativity, provided the molecule’s geometry prevents the charges from canceling out. This results in one end of the molecule having a slight negative charge (\(\delta-\)) and the opposite end having a slight positive charge (\(\delta+\)). The attraction arises as the positive end of one polar molecule electrostatically attracts the negative end of an adjacent polar molecule. Because these dipoles are permanent, the resulting attraction is generally stronger than the temporary LDFs between molecules of similar size.
Hydrogen Bonding
Hydrogen bonding is a particularly strong type of Dipole-Dipole interaction. For this force to form, a hydrogen atom must be covalently bonded to one of three small, highly electronegative atoms: nitrogen (\(\text{N}\)), oxygen (\(\text{O}\)), or fluorine (\(\text{F}\)). The high electronegativity of these atoms pulls the bonding electrons strongly away from the hydrogen atom, leaving the hydrogen with a large partial positive charge and a highly exposed nucleus. This exposed, partially positive hydrogen is then strongly attracted to a lone pair of electrons on an adjacent \(\text{N}\), \(\text{O}\), or \(\text{F}\) atom in a separate molecule. Water (\(\text{H}_2\text{O}\)) is the most well-known example, where the powerful attraction between molecules is responsible for properties such as its exceptionally high boiling point.
Ion-Dipole Forces
Ion-Dipole forces represent the electrostatic interaction between a fully charged ion and a neutral, polar molecule. This interaction involves a full, integer charge (either a cation or an anion) interacting with a partial charge on a dipole. Since the charge on an ion is much greater than the partial charge on a dipole, Ion-Dipole forces are often significantly stronger than the other types of intermolecular forces between neutral molecules. This force is important in the process of solvation, which is how ionic compounds dissolve in polar solvents. When sodium chloride (\(\text{NaCl}\)) dissolves in water, the polar water molecules orient themselves around the separated ions. The partially negative oxygen end attracts the positive sodium ion (\(\text{Na}^+\)), while the partially positive hydrogen ends attract the negative chloride ion (\(\text{Cl}^-\)).
How Intermolecular Forces Shape Matter
The collective strength of intermolecular forces directly influences a substance’s observable macroscopic properties. Stronger IMFs mean that molecules are more tightly held together, requiring more energy to separate them. This energy requirement is most clearly seen in the boiling point, the temperature needed to overcome molecular attractions and transition into a gas. Substances with stronger intermolecular forces, such as those that exhibit hydrogen bonding, have higher boiling and melting points compared to those held together only by weaker LDFs. Viscosity, which is a liquid’s resistance to flow, is also increased by stronger IMFs because the molecules resist moving past one another. The type of force present dictates solubility, following the principle that polar substances tend to dissolve well in other polar solvents.