Properties of Substances and Mixtures
Intermolecular forces and bulk properties
The strength of attractions between particles sets boiling point, vapor pressure, viscosity, surface tension, and solubility. Rank forces: ion–dipole > hydrogen bonding > dipole–dipole > London dispersion. Hydrogen bonding requires H covalently bonded to N, O, or F; a C–H bond does not qualify. London dispersion grows with electron count and surface area, so larger and less branched molecules have higher boiling points. Polarizability explains why I2 is a solid while Cl2 is a gas.
Gases: the ideal gas law PV = nRT assumes negligible particle volume and no intermolecular attractions. Real gases deviate most at high pressure and low temperature, where particles are close and attractions matter. Kinetic molecular theory connects temperature to average kinetic energy; the Maxwell–Boltzmann distribution broadens and shifts right as T rises, increasing the fraction of collisions with energy above Ea. Effusion and diffusion rates vary inversely with the square root of molar mass.
Solutions and separations: 'like dissolves like,' and ion–dipole interactions drive electrolytes into water. Colligative properties depend on particle count, so use the van't Hoff factor i for strong electrolytes. Vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure all scale with total solute particles. Distillation separates by boiling point; chromatography separates by differential affinity for stationary versus mobile phases.
Traps: a strong electrolyte does not mean strong intermolecular force; a high boiling point can come from any strong force. Check whether a solute is molecular (i = 1) or ionic (i = number of ions).
Worked example: 0.10 mol NaCl in 1.00 kg water gives i = 2 and ΔTf = i·Kf·m = 2(1.86)(0.10) = 0.37 °C.
Rule of thumb: identify the strongest force present, then predict the property.