Intermolecular forces (IMFs) are the attractive and repulsive forces that exist between neighboring molecules. These forces differ fundamentally from the much stronger intramolecular forces, such as covalent bonds, which hold atoms together within a single molecule. While covalent bonds require hundreds of kilojoules of energy to break, IMFs are significantly weaker. However, their collective effect dictates the physical properties of matter observed on a macroscopic scale. Understanding these forces explains why substances behave the way they do, from the temperature at which water boils to how a cell membrane functions.
Governing Physical States
Intermolecular forces determine a substance’s physical state—solid, liquid, or gas—at a given temperature and pressure. The state depends on a balance between the kinetic energy of the molecules and the strength of the attractive forces holding them together. Stronger IMFs require more energy to overcome, leading to higher temperatures for phase transitions.
For example, water molecules exhibit strong hydrogen bonding. This attraction forces water to remain a liquid until it reaches 100°C, a temperature much higher than other molecules of similar size. Methane (\(text{CH}_4\)), a nonpolar molecule, only experiences the weakest IMFs, called London Dispersion Forces. Because these forces are weak, methane is a gas that boils at \(-161.5^{circ}text{C}\).
Boiling requires providing enough thermal energy to break the molecules free from their neighbors’ attractive forces. Similarly, the melting point is the temperature at which molecules gain enough energy to overcome the forces that lock them into fixed positions. The relative strengths of the three main IMFs—London Dispersion, Dipole-Dipole, and Hydrogen Bonding—directly correlate with a substance’s boiling and melting points.
Driving Solubility and Mixing
The ability of one substance to dissolve in another, known as solubility, is governed by the principle: “like dissolves like.” This rule means the intermolecular forces of the solute and the solvent must be comparable for mixing to occur. Dissolution is a competition between three sets of forces: solute-solute, solvent-solvent, and the newly formed solute-solvent interactions.
For a solute to dissolve, the energy required to break the original solute-solute and solvent-solvent attractions must be compensated by the energy released from forming new solute-solvent attractions. Polar solvents, such as water, dissolve polar solutes, like ethanol, because both can form strong hydrogen bonds and dipole-dipole interactions. Conversely, nonpolar substances, such as oil, do not dissolve in water because the weak London Dispersion Forces between oil and water cannot compensate for the strong hydrogen bonds that must be broken between water molecules.
When an ionic compound, like table salt (\(text{NaCl}\)), dissolves in water, the strong ionic bonds are overcome by ion-dipole forces. The polar water molecules surround the charged ions in a process called solvation. The negative oxygen end surrounds the positive sodium ions, and the positive hydrogen end surrounds the negative chloride ions, effectively pulling the crystal lattice apart.
Shaping Biological Structures
Intermolecular forces construct the complex, functional structures required for life, providing stability without sacrificing necessary flexibility. In proteins, the specific three-dimensional folding pattern (tertiary structure) is stabilized by various IMFs involving amino acid side chains. The hydrophobic effect is a major driving force, causing nonpolar side chains to cluster in the protein’s interior, shielded from the surrounding water.
Hydrogen bonds are key to forming the protein’s secondary structures, such as the \(alpha\)-helix and the \(beta\)-sheet, where they form between the carbonyl oxygen and the amide hydrogen atoms of the backbone. In the DNA double helix, hydrogen bonds form the rungs of the ladder, pairing adenine with thymine (two bonds) and guanine with cytosine (three bonds). Although individual hydrogen bonds are weak, the cumulative effect provides stability for the genetic material.
The DNA helix is further stabilized by base-stacking interactions, which combine van der Waals forces and hydrophobic interactions between the flat, aromatic rings of adjacent bases. Cell membranes are also products of IMFs, as phospholipid molecules spontaneously arrange into a lipid bilayer in water. This arrangement is driven by the hydrophobic effect, forcing nonpolar fatty acid tails into the interior while polar head groups remain exposed to the aqueous exterior, stabilized by hydrogen bonding.
Practical Applications in Engineered Materials
Scientists and engineers manipulate intermolecular forces to control material properties and develop new technologies. In polymer science, the mechanical properties of plastics and fibers result directly from the IMFs between the long molecular chains. Materials like Nylon 6,6, used in durable fabrics, exhibit high tensile strength because their repeating units maximize hydrogen bonding between adjacent polymer chains.
Conversely, flexible materials, such as certain types of polyethylene, have less ordered chains where only weak London Dispersion Forces operate, allowing the chains to slip past one another easily. In drug design, a pharmaceutical compound’s efficacy depends on its ability to bind specifically to a target receptor site. Drug molecules are engineered with complementary shape and polarity to maximize non-covalent attractions, such as hydrogen bonding, dipole-dipole forces, and van der Waals interactions, ensuring they lock into the receptor site.
The function of common adhesives is a macroscopic display of concentrated IMFs. Adhesion relies on the liquid adhesive flowing across the substrate surface to achieve intimate molecular contact, a process called wetting. Once set, the cohesive forces within the adhesive and the adhesive forces between the adhesive and the surface, primarily van der Waals forces, are maximized across the large surface area, resulting in a strong bond.