Intermolecular Forces: London Dispersion, Dipole-Dipole, and Hydrogen Bonding
A gecko can run up a vertical glass wall using van der Waals forces in its toe hairs. Water striders walk on water surfaces held together by hydrogen bonding. DNA’s double helix is stabilized by hydrogen bonds between base pairs. The forces between molecules, while individually weak, collectively determine whether a substance is gas, liquid, or solid at room temperature.
Intermolecular forces explain why propane (C3H8, boiling point -42°C) is a gas while butane (C4H10, boiling point -1°C) is a liquid, despite both being nonpolar hydrocarbons. They explain why water has such a high boiling point for its molecular mass. Understanding these forces is essential for predicting physical properties and phase behavior.
Types of Intermolecular Forces
Intermolecular forces range from relatively strong (hydrogen bonding) to very weak (London dispersion). All are electrostatic in origin, involving attractions between partial charges or temporary charge distributions.
London Dispersion Forces
London dispersion forces arise from temporary fluctuations in electron distribution within molecules. At any instant, the electron cloud may be unevenly distributed, creating an instantaneous dipole. This dipole induces a dipole in a neighboring molecule, creating a brief attraction.
Dispersion forces are present between all molecules, polar and nonpolar. They are the only intermolecular force between nonpolar molecules. Their strength increases with molecular size and surface area — larger molecules have more electrons and greater polarizability.
This explains the boiling point trend among alkanes. Methane (CH4) boils at -164°C, ethane (C2H6) at -89°C, propane (C3H8) at -42°C, and butane (C4H10) at -1°C. Each additional CH2 group increases dispersion forces and raises the boiling point by about 40°C.
Molecular shape also matters. Linear molecules have greater surface area for intermolecular contact than branched molecules. Pentane (linear, boiling point 36°C) boils 8°C higher than neopentane (branched, boiling point 28°C), despite having the same molecular formula.
Dipole-Dipole Interactions
Polar molecules have permanent dipoles — a partial positive charge on one end and partial negative on the other. The positive end of one molecule attracts the negative end of another. These dipole-dipole interactions are stronger than London dispersion forces for small molecules.
The strength of dipole-dipole interactions depends on the magnitude of the molecular dipole moment. Molecules with large dipole moments, like acetone (dipole moment 2.88 D, boiling point 56°C), have stronger attractions than molecules with smaller dipole moments, like formaldehyde (dipole moment 2.33 D, boiling point -19°C).
In mixtures, dipole-dipole interactions affect solubility. Polar molecules dissolve in polar solvents (like dissolves like). Nonpolar molecules dissolve in nonpolar solvents. This principle is fundamental to solution chemistry.
Hydrogen Bonding
Hydrogen bonding is a special case of dipole-dipole interaction that occurs when hydrogen is bonded to one of three highly electronegative elements: nitrogen, oxygen, or fluorine. The small size of hydrogen allows the electronegative atom to approach very closely, creating a strong interaction.
Hydrogen bonds are about 5-30 kJ/mol — much weaker than covalent bonds (200-400 kJ/mol) but significantly stronger than other intermolecular forces (1-10 kJ/mol). They are responsible for water’s exceptional properties.
Water (H2O) has a boiling point of 100°C, while hydrogen sulfide (H2S, no hydrogen bonding) boils at -60°C. The difference of 160°C is almost entirely due to hydrogen bonding. Each water molecule can form up to four hydrogen bonds, creating a three-dimensional network that gives water its high surface tension, heat capacity, and boiling point.
Ion-Dipole Forces
Ion-dipole forces occur between ions and polar molecules. When sodium chloride dissolves in water, water molecules surround the Na+ ions (oxygen end pointing toward the ion) and Cl- ions (hydrogen end pointing toward the ion). These ion-dipole attractions overcome the ionic bonds in the crystal lattice.
The strength of ion-dipole forces depends on ion charge density. Small, highly charged ions (Mg2+) have stronger ion-dipole attractions than large, singly charged ions (Cs+). This explains why Mg2+ has a larger hydration shell than Cs+.
Effects on Physical Properties
Boiling Point
Boiling requires overcoming intermolecular forces to separate molecules into the gas phase. Substances with stronger intermolecular forces have higher boiling points. The type and strength of intermolecular forces determine the boiling temperature.
Among molecules of similar size, hydrogen-bonding substances have the highest boiling points (water, ethanol), followed by dipole-dipole substances (acetone), then nonpolar substances (hexane). Among nonpolar molecules, larger molecules boil higher due to stronger London dispersion forces.
Viscosity
Viscosity measures resistance to flow. Stronger intermolecular forces make liquids more viscous. Glycerol, with three OH groups capable of extensive hydrogen bonding, is very viscous. Motor oil is viscous due to long hydrocarbon chains with strong dispersion forces between them.
Temperature reduces viscosity by increasing molecular motion, making it easier to overcome intermolecular attractions. This is why warm oil flows more easily than cold oil.
Surface Tension
Surface tension is the energy required to increase liquid surface area. Molecules at the surface experience net inward attraction because they have fewer neighbors. This creates a “skin” effect that allows water striders to walk on water.
Hydrogen bonding gives water an unusually high surface tension (72.8 mN/m at 20°C). Most organic solvents have much lower surface tensions (acetone: 23.7 mN/m, hexane: 18.4 mN/m). This is why water beads up on surfaces while organic solvents spread out.
Capillary Action
Capillary action — water rising in a narrow tube against gravity — combines adhesion (attraction between water and the tube wall) and cohesion (attraction between water molecules). Adhesion pulls water up the tube walls, and cohesion pulls the rest of the water along.
Plants use capillary action to transport water from roots to leaves through xylem tubes. Trees over 100 meters tall rely on this mechanism combined with transpiration pull.
Intermolecular Forces in Biological Systems
Hydrogen bonding determines protein structure. Alpha helices and beta sheets are stabilized by hydrogen bonds between backbone amide and carbonyl groups. The specificity of base pairing in DNA (A-T and G-C pairs) comes from hydrogen bonding patterns.
Enzyme-substrate recognition relies on complementary intermolecular forces. The substrate must fit the enzyme’s active site like a key in a lock, with hydrogen bonds, ionic interactions, and dispersion forces all contributing. Molecular geometry determines whether these forces can align optimally.
Phase Changes and Critical Phenomena
Intermolecular forces determine the temperatures and pressures at which phase changes occur. The normal boiling point is the temperature at which vapor pressure equals atmospheric pressure. Stronger intermolecular forces mean lower vapor pressure and higher boiling point.
The critical temperature is the temperature above which a gas cannot be liquefied no matter how much pressure is applied. Water’s critical temperature is 374°C, reflecting its strong hydrogen bonding. Helium’s critical temperature is -268°C — its extremely weak dispersion forces make it nearly impossible to liquefy.
Supercritical Fluids
Above the critical temperature and pressure, a substance becomes a supercritical fluid with properties between those of liquids and gases. Supercritical CO2 (Tc = 31°C, Pc = 73 atm) is widely used as a solvent for decaffeinating coffee and extracting natural products. Supercritical fluids have liquid-like density with gas-like viscosity and diffusion, making them excellent extraction solvents.
Supercritical water oxidation destroys hazardous organic wastes by oxidizing them in water above 374°C and 218 atm. The process converts organic compounds completely to CO2 and water without the harmful byproducts of incineration.
Intermolecular Forces in Materials Science
Polymer properties are largely determined by intermolecular forces. Polyethylene consists of long nonpolar hydrocarbon chains held together by London dispersion forces. High-density polyethylene has more crystalline regions with closer chain packing, making it stronger and more rigid than low-density polyethylene.
Nylon incorporates hydrogen bonding between polymer chains through amide groups. These hydrogen bonds increase tensile strength and melting point, making nylon suitable for fibers and engineering plastics. Kevlar, with extensive hydrogen bonding between aligned polymer chains, has exceptional strength-to-weight ratio used in bulletproof vests.
Frequently Asked Questions
What is the strongest intermolecular force? Hydrogen bonding is the strongest type of intermolecular force, ranging from 5-30 kJ/mol. Ion-dipole forces can be stronger but are usually classified separately.
Why does water have such a high boiling point? Water molecules form extensive hydrogen bonds. Each water molecule can form four hydrogen bonds, creating a network that requires significant energy to break.
Do all molecules have London dispersion forces? Yes. London dispersion forces exist between all molecules, whether polar or nonpolar. They are the only intermolecular force for nonpolar molecules.
How do intermolecular forces affect solubility? “Like dissolves like” — polar and ionic substances dissolve in polar solvents. Nonpolar substances dissolve in nonpolar solvents. Solubility requires similar types and strengths of intermolecular forces.
Why does water have high surface tension? Strong hydrogen bonding between water molecules creates a cohesive force at the surface that resists penetration. This allows water striders to walk on water and causes water to form droplets rather than spreading out.