Intermolecular forces (IMFs) are the attractions between individual particles at the molecular level. These non-chemical forces cause molecules to associate with one another, influencing whether a substance exists as a solid, liquid, or gas at a given temperature. The relative strength of these forces dictates a substance’s macroscopic properties, such as its melting point and solubility.
Defining Intermolecular Forces
Intermolecular forces are the attractive or repulsive forces that act between neighboring molecules or atoms. These forces are generally electrostatic, arising from the attraction between positive and negative charges, but they do not involve the sharing or transfer of electrons. While IMFs are present in all states of matter, their influence is most apparent in liquids and solids, where molecules are held closely together.
It is important to distinguish IMFs from intramolecular forces, which are the chemical bonds that exist within a single molecule. Intramolecular forces, such as covalent and ionic bonds, are significantly stronger than intermolecular forces. For example, the covalent bonds holding the two hydrogen atoms and one oxygen atom together in a single water molecule are intramolecular forces.
In contrast, intermolecular forces are responsible for the physical properties of substances, such as the transition from liquid water to steam. When water boils, the covalent bonds within the H₂O molecule do not break. Instead, the weaker intermolecular attractions between separate water molecules are overcome, allowing the molecules to escape into the gaseous phase.
The Three Primary Types of Intermolecular Forces
Intermolecular forces are classified into three major types, ranging in strength from weakest to strongest.
London Dispersion Force (LDF)
The weakest type of attraction, present in all atoms and molecules, is the London Dispersion Force (LDF). These forces arise from the continuous, random motion of electrons, which creates a temporary, instantaneous dipole moment. This fleeting charge imbalance induces a corresponding dipole in a neighboring molecule, resulting in a weak, momentary attraction. The strength of LDFs increases with the size and molecular weight of the substance because larger molecules are more easily polarized.
Dipole-Dipole Force
A stronger type of attraction is the Dipole-Dipole Force, which occurs only between molecules that possess a permanent dipole. These molecules have unequally shared electrons, creating a permanent partial positive end (\(\delta+\)) and a permanent partial negative end (\(\delta-\)). The opposing partial charges on adjacent molecules then align and attract each other. This constant attraction makes dipole-dipole forces stronger than LDFs.
Hydrogen Bonding
The third and strongest type of IMF is Hydrogen Bonding, a specialized form of dipole-dipole interaction. This force occurs when a hydrogen atom is covalently bonded to nitrogen (N), oxygen (O), or fluorine (F). Because these highly electronegative atoms strongly pull electrons away from the hydrogen atom, they create an extremely polarized bond. The partially positive hydrogen atom on one molecule is then strongly attracted to the partially negative N, O, or F atom on a neighboring molecule.
How Intermolecular Forces Govern Physical States
The collective strength of a substance’s intermolecular forces directly determines its physical state at a given temperature and pressure. In a gas, molecules move fast enough to overcome all IMFs, allowing them to separate and fill any volume. In a liquid, IMFs are strong enough to keep molecules close together but weak enough to allow them to slide past one another. In a solid, the IMFs lock the molecules into fixed, rigid positions.
Phase changes, such as melting or boiling, require supplying enough energy to overcome these intermolecular attractions. A substance with strong IMFs, like water with its extensive hydrogen bonding, requires significant heat energy to transition from liquid to gas, resulting in a higher boiling point. Conversely, substances like methane, which only have weak London Dispersion Forces, require little energy to separate molecules, resulting in extremely low boiling points.
IMFs also dictate the solubility of substances, a concept summarized by the principle “like dissolves like.” Polar solvents, which have strong IMFs, are effective at dissolving polar solutes because new, similarly strong forces can form between the solvent and solute molecules. For instance, water (polar) readily dissolves sugar (polar) because both can form hydrogen bonds. Nonpolar substances, which rely only on London Dispersion Forces, dissolve best in nonpolar solvents. When a nonpolar substance encounters a polar solvent, the polar molecules preferentially attract each other, excluding the nonpolar solute.