Intermolecular forces (IMFs) are the attractive or repulsive electrostatic forces that exist between individual molecules or atoms. These forces arise from the interactions between positive and negative charges on neighboring particles. While significantly weaker than the forces holding atoms together within a single molecule, IMFs are important. Their presence and strength determine many physical properties of substances, such as whether a compound is a gas, liquid, or solid at room temperature.
Intermolecular vs. Intramolecular Bonds
The distinction between intermolecular and intramolecular forces is based on their location relative to the molecule. Intramolecular forces are the strong chemical bonds that exist within a single molecule, holding the atoms together. Examples include the covalent bonds in a water molecule or the ionic bonds in table salt. These bonds are extremely strong and require large amounts of energy to break, providing molecules with chemical stability.
Intermolecular forces, by contrast, are the much weaker attractions that occur between separate, neighboring molecules. They are responsible for molecules sticking together, leading to the formation of liquids and solids. When a substance changes physical state, such as boiling water, the process involves overcoming these weak intermolecular forces, not breaking the strong intramolecular bonds. Intramolecular bonds are typically hundreds of times stronger than intermolecular forces.
The Three Primary Intermolecular Forces
Intermolecular interactions are categorized into three principal types, generally ranked by relative strength. The strongest is hydrogen bonding, a special type of dipole-dipole interaction. This attraction occurs when a hydrogen atom is covalently bonded to a highly electronegative atom (nitrogen, oxygen, or fluorine). The large electronegativity difference creates a strong partial positive charge on the hydrogen, which is then strongly attracted to the lone pair of electrons on a neighboring N, O, or F atom.
Dipole-dipole forces are the next strongest type and occur between molecules with a permanent, uneven distribution of electric charge (a permanent dipole). In these polar molecules, one end is partially positive and the other is partially negative, causing them to align and attract electrostatically. The strength of this attraction is proportional to the magnitude of the molecule’s dipole moment.
The weakest of the three primary forces are the London Dispersion Forces (LDFs), which are present in all molecules, polar and nonpolar. LDFs result from temporary, momentary dipoles that arise from the constant movement of electrons. These fleeting dipoles in one molecule can induce a temporary, opposing dipole in an adjacent molecule, creating a weak attraction. LDFs increase in strength with the size and molecular weight of a molecule because larger molecules have more electrons and more easily deformable electron clouds.
How Intermolecular Forces Shape Matter
The strength of intermolecular forces dictates many physical properties observed in substances. These forces govern the state of matter a substance exists in at a given temperature. Substances relying only on weak London Dispersion Forces require little energy to separate and are typically gases at room temperature. Conversely, substances with medium or strong intermolecular forces, like water with its hydrogen bonding, are more often liquids or solids because the molecules are held together more tightly.
A direct consequence of intermolecular force strength is the boiling and melting points of a substance. To convert a liquid into a gas or a solid into a liquid, energy must be supplied to overcome the attractions between molecules. Therefore, substances with stronger forces require a higher input of heat to achieve a phase change, resulting in higher boiling and melting points. Water, for instance, has an unusually high boiling point for its small size because of the relatively strong hydrogen bonds between its molecules.
Intermolecular forces also determine a substance’s solubility, summarized by the principle “like dissolves like.” Polar solvents, like water, readily dissolve polar solutes because the strong attractions between the solvent and solute molecules can overcome the forces holding the solute molecules together. Nonpolar substances, which only exhibit weak London Dispersion Forces, dissolve best in other nonpolar solvents where similar, weak attractions can form.