Intermolecular forces (IMFs) are the attractive forces that exist between separate molecules. These forces mediate the interactions between neighboring particles and profoundly influence the physical properties of a substance. Molecular polarity refers to the uneven distribution of electron density within a molecule, creating a positive end and a negative end. This charge separation, or dipole moment, is the primary determinant of both the type and the strength of the intermolecular forces a substance can exhibit.
Understanding Molecular Polarity
The root cause of molecular polarity lies in electronegativity, an atom’s ability to attract shared electrons towards itself in a chemical bond. When two atoms bond, a significant difference in electronegativity causes shared electrons to spend more time near the more attractive atom. This unequal sharing creates a bond dipole, resulting in partial positive charges on the less electronegative atom and partial negative charges on the more electronegative atom.
The distribution of this charge is not a simple binary state but exists along a continuous spectrum ranging from nonpolar covalent, where electrons are shared equally, to highly polar covalent, and eventually to ionic bonding. Even if a molecule contains polar bonds, its overall molecular polarity—or net dipole moment—is determined by its three-dimensional geometry. If the individual bond dipoles cancel each other out symmetrically, as in carbon dioxide, the molecule will be nonpolar despite having polar bonds.
Dipole-Dipole Interactions
The most direct consequence of permanent molecular polarity is the dipole-dipole interaction, which occurs specifically between two polar molecules. These forces are electrostatic attractions where the partial positive end of one molecule aligns with the partial negative end of a neighboring molecule. This alignment minimizes potential energy and draws the molecules closer together.
These attractive forces are generally stronger than London Dispersion Forces but remain considerably weaker than the intramolecular covalent bonds. A classic example is hydrogen chloride (\(\text{HCl}\)), where the more electronegative chlorine atom carries the partial negative charge, attracting the partially positive hydrogen atom of an adjacent \(\text{HCl}\) molecule. Dipole-dipole interactions typically have strengths ranging from 5 to 20 kilojoules per mole (\(\text{kJ/mol}\)).
The Unique Strength of Hydrogen Bonding
Hydrogen bonding represents an exceptionally strong form of dipole-dipole interaction due to its profound influence on chemical and biological systems. This force is not a true chemical bond but a particularly powerful attraction that forms only when a hydrogen atom is covalently bonded to one of three small, highly electronegative atoms: nitrogen (\(\text{N}\)), oxygen (\(\text{O}\)), or fluorine (\(\text{F}\)). This pairing creates an extreme polarity, leaving the hydrogen atom with a large partial positive charge and a small atomic size, allowing it to get very close to a lone pair of electrons on a neighboring \(\text{N}\), \(\text{O}\), or \(\text{F}\) atom.
Water (\(\text{H}_2\text{O}\)) is the primary example, where the oxygen atom strongly pulls electrons away from the two hydrogen atoms. The resulting hydrogen bonds link water molecules together in a network. These interactions, with dissociation energies around 15–25 \(\text{kJ/mol}\), are responsible for water’s unusually high boiling point compared to molecules of similar size.
Interplay with London Dispersion Forces
London Dispersion Forces (LDFs) are the only type of intermolecular force present in all molecules, both polar and nonpolar. These forces arise from the continuous motion of electrons, which can momentarily create an uneven distribution of charge, resulting in an instantaneous, temporary dipole. This temporary dipole can then induce a corresponding dipole in a neighboring molecule, leading to a weak, short-lived attraction.
For nonpolar molecules, LDFs are the sole attractive force, and their strength increases significantly with molecular size and surface area because larger electron clouds are more easily distorted, a property called polarizability. However, in polar molecules, LDFs are always present alongside the stronger dipole-dipole forces or hydrogen bonds. Although LDFs exist in polar substances, their contribution is often overshadowed by the much more powerful polarity-driven interactions.
Macroscopic Properties Influenced by IMFs
The strength of intermolecular forces directly dictates a substance’s observable physical characteristics. Substances with stronger IMFs require significantly more thermal energy to overcome the molecular attractions and change state. This results in higher boiling points and melting points for polar compounds compared to nonpolar compounds of similar molecular mass.
For example, the polar compound iodine monochloride (\(\text{ICl}\)) is a solid at \(0^\circ\text{C}\), while the nonpolar bromine (\(\text{Br}_2\)), which has a very similar molecular weight, is a liquid at the same temperature, illustrating the impact of a permanent dipole moment. Polarity also governs solubility through the principle of “like dissolves like,” where polar solvents, such as water, effectively dissolve polar solutes by forming strong IMFs with them. Liquids with stronger IMFs also exhibit higher viscosity because the molecules resist movement relative to one another.