Intermolecular Forces

Introduction to Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion between molecules, which play a crucial role in determining the physical properties of substances, such as boiling point, melting point, solubility, and state of matter. These forces explain why some solids dissolve in water, why noble gases can liquefy, and why some substances are gases at room temperature while others are liquids or solids.

• Intermolecular forces may be attractive or repulsive.

• Johannes D. van der Waals was the first to postulate intermolecular forces to account for the properties of real gases.

Types of Intermolecular Forces

• van der Waals forces include:

  • London (dispersion) forces

  • Dipole-dipole forces

  • Dipole-induced dipole forces

• Other intermolecular forces:

  • Ion-dipole interactions

  • Ion-induced dipole interactions

  • Hydrogen bonding

London (Dispersion) Forces

London forces, also known as dispersion forces, arise from temporary variations in electron density in atoms and molecules. At any instant, the electron distribution may be unsymmetrical, producing an instantaneous dipole that can induce a transient dipole in neighboring molecules.

• Present between all molecules, whether polar or nonpolar.

• Stronger in molecules that are easily polarizable.

• Larger and heavier molecules exhibit stronger dispersion forces than smaller and lighter ones.

Example: Neon atoms do not have a permanent dipole, but can exhibit dispersion forces due to temporary dipoles.

Magnitude of Dispersion Forces

• The shape of molecules affects the magnitude of dispersion forces.

• Cylindrical shapes (e.g., n-pentane) have greater surface area and stronger dispersion forces than spherical shapes (e.g., neopentane).

Dipole-Dipole Forces

Dipole-dipole forces occur between molecules with permanent dipoles. The positive end of one molecule is attracted to the negative end of another.

• Strength depends on the magnitude of the dipole moment and the distance between molecules.

Dipole-Induced Dipole Forces

These forces operate between polar molecules (with permanent dipoles) and nonpolar molecules. The polar molecule induces a dipole in the neighboring nonpolar molecule.

• Interaction energy depends on the dipole moment of the polar molecule.

• Example: Oxygen gas can dissolve in water because the permanent dipole in water induces a dipole in oxygen.

• Example: Nonpolar iodine dissolves in polar ethanol due to dipole-induced dipole forces.

Ion-Dipole and Ion-Induced Dipole Interactions

These interactions occur between ions and polar molecules. The strength depends on the charge and size of the ion and the magnitude of the dipole moment of the polar molecule.

• Example: Hydration of sodium ions (Na+) and chloride ions (Cl-) in water.

Hydrogen Bonding

Hydrogen bonding is an electrostatic force of attraction between a covalently bonded hydrogen atom and an electronegative atom (such as fluorine, oxygen, or nitrogen) of another molecule.

• Conditions for hydrogen bond formation:

  • Hydrogen atom attached to a highly electronegative atom (F, O, N).

  • The electronegative atom acts as a hydrogen bond donor or acceptor.

• Hydrogen bonding causes:

  • Higher boiling point in ethanol compared to diethyl ether.

  • Higher viscosity of sulfuric acid and glycerol.

  • Water is a liquid, whereas H2S is a gas.

  • Dimerization of carboxylic acid molecules.

  • Cage-like structure of water molecules in ice, causing lower density of ice.

Comparison of Intermolecular Forces

Solutions: Effect of Intermolecular Forces

Solute-Solvent Interactions

For a solute and solvent to mix, the attractive forces between solute-solute and solvent-solvent must be overcome. The process is often endothermic, but the formation of new solute-solvent attractions is exothermic.

• Oil and water do not mix because oil is nonpolar and water is polar. "Like dissolves like" is a guiding principle for solubility.

Common Types of Solutions

Solution Process and Enthalpy of Solution

The process of solution formation involves three steps:

1. Separating the solute into its constituent particles (, endothermic)

2. Separating the solvent particles from each other (, endothermic)

3. Mixing the solute particles with the solvent particles (, exothermic)

The overall enthalpy change () is:

• If the sum of endothermic terms is approximately equal to the exothermic term, is about zero.

• If the sum of endothermic terms is less than the exothermic term, is negative (exothermic).

• If the sum of endothermic terms is greater than the exothermic term, is positive (endothermic).

Solubility Limit

• A saturated solution contains the maximum amount of solute that can dissolve at a given temperature.

• An unsaturated solution contains less solute than saturation.

• A supersaturated solution contains more solute than saturation and is unstable.

Temperature Dependence of Solubility

• For most solids, solubility increases with increasing temperature (when is endothermic).

• For gases, solubility decreases with increasing temperature.

Pressure Dependence of Gas Solubility: Henry's Law

The solubility of a gas () in a liquid is directly proportional to its partial pressure () above the liquid:

• is Henry's law constant, which is specific for each gas and temperature.

Solution Concentration Terms

Summary Table: Common Laboratory Solvents

Key Concepts and Examples

• "Like dissolves like": Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.

• Enthalpy of solution () determines whether the solution process is exothermic or endothermic.

• Hydrogen bonding leads to unique properties in water, such as high boiling point and lower density of ice.

• Solubility curves can be used to predict saturation, unsaturation, or supersaturation at different temperatures.

• Henry's Law explains the effect of pressure on the solubility of gases in liquids.

Additional info: Some explanations and tables have been expanded for clarity and completeness based on standard General Chemistry curriculum.