Intermolecular forces (IMFs) are the attractive or repulsive forces that exist between individual molecules in a substance. These forces differ fundamentally from intramolecular bonds (covalent or ionic bonds) which hold atoms together within a single molecule. Although intramolecular bonds are much stronger, the collective influence of IMFs dictates a substance’s physical properties, such as whether it exists as a solid, liquid, or gas. Identifying which IMFs are present requires a systematic analysis of a molecule’s composition and three-dimensional structure.
The Universal Force
The London Dispersion Force (LDF) is present in every atom and molecule, regardless of structure or polarity. Sometimes called induced dipole-induced dipole forces, LDFs arise from the continuous, random motion of electrons. Electrons may be momentarily distributed unevenly, creating a transient, or instantaneous, dipole moment. This imbalance in one molecule can induce a corresponding dipole in a neighbor, resulting in a weak, short-lived attraction.
The strength of LDFs is directly related to polarizability, which is the ease with which a molecule’s electron cloud can be distorted to form a temporary dipole. Larger molecules, and those with greater molecular mass, have more electrons farther from the nucleus. This makes the electron cloud more diffuse and less tightly held. Therefore, larger, heavier molecules are more easily polarized, resulting in significantly stronger LDFs. For instance, iodine (\(text{I}_2\)) is a solid at room temperature, while the much lighter fluorine (\(text{F}_2\)) is a gas, due to the difference in their LDF strength.
Determining Molecular Polarity
Identifying the remaining, typically stronger IMFs depends entirely on whether a molecule possesses a permanent net dipole moment (polarity). This determination is a two-step process beginning with analyzing the individual bonds. Bond polarity is established by the difference in electronegativity between the two bonded atoms. Electrons are pulled toward the more electronegative atom, creating a partial negative charge (\(delta^-\)) and leaving the other atom with a partial positive charge (\(delta^+\)).
The second step involves considering the molecule’s overall three-dimensional geometry. Even if a molecule contains polar bonds, the overall molecule may be nonpolar if the individual bond dipoles cancel out due to symmetry. For example, in carbon dioxide (\(text{CO}_2\)), the linear geometry places the two polar \(text{C-O}\) bond dipoles in opposite directions, causing them to nullify one another, resulting in a nonpolar molecule. Conversely, water (\(text{H}_2text{O}\)) has two polar \(text{O-H}\) bonds, but its bent geometry prevents the dipoles from canceling, giving the molecule a net dipole moment and making it polar. Only polar molecules proceed to the next level of IMF analysis.
Identifying Specific Polar Forces
Molecules confirmed to be polar exhibit dipole-dipole interactions. These are electrostatic attractions between the permanent positive end of one molecule and the permanent negative end of a neighbor. Dipole-dipole forces are generally stronger than LDFs in molecules of comparable size because they involve fixed, rather than transient, partial charges. The magnitude of this force depends on the size of the molecule’s permanent dipole moment.
A special, highly potent form of dipole-dipole interaction is hydrogen bonding, the strongest of the common intermolecular forces. Hydrogen bonding occurs only when a hydrogen atom is covalently bonded directly to nitrogen (\(text{N}\)), oxygen (\(text{O}\)), or fluorine (\(text{F}\)). The extreme electronegativity of these atoms creates a large partial positive charge on the hydrogen nucleus. This positive hydrogen is then strongly attracted to a lone pair of electrons on a neighboring \(text{N}\), \(text{O}\), or \(text{F}\) atom. The small size of the hydrogen atom allows it to approach the electron pair very closely, significantly amplifying the strength of the electrostatic attraction beyond a typical dipole-dipole force.
Linking Forces to Physical Properties
Identifying a molecule’s IMFs is not merely an academic exercise, as the type and strength of these forces directly govern a substance’s macroscopic physical properties. The most direct consequence is the effect on melting and boiling points. To change a substance from a liquid to a gas (or solid to liquid), energy must be supplied to overcome the attractive forces between the molecules.
Stronger intermolecular forces require a greater amount of energy to disrupt, leading to higher melting and boiling temperatures. Water’s unusually high boiling point of \(100^{circ}text{C}\) is a direct result of its ability to form extensive networks of strong hydrogen bonds. In contrast, nonpolar molecules relying only on weak LDFs have lower boiling points, requiring little energy to separate them. IMFs also influence solubility, following the rule that “like dissolves like.” This means polar substances mix well with other polar substances (like water), and nonpolar substances mix well with other nonpolar substances (like oil).