States of Matter and Intermolecular Forces
What you must understand
Compare the gas, liquid, and solid states in terms of particle spacing, kinetic energy, and compressibility. For gases, use the ideal gas law PV = nRT and Dalton's law of partial pressures, and understand the kinetic molecular theory assumptions. Know the van der Waals equation as a correction for molecular volume and attractive forces, and be able to say when real gases deviate most (high pressure, low temperature).
For condensed phases, rank intermolecular forces: ion-dipole, hydrogen bonding, dipole-dipole, London dispersion. Dispersion forces strengthen with polarisability and molecular size. Relate these forces to boiling point, vapour pressure, surface tension, viscosity, and solubility. Understand phase diagrams: triple point, critical point, and the meaning of slopes, and use heating curves to identify enthalpy of fusion and vaporisation.
Traps to watch for
- Using temperature in Celsius in the ideal gas law; convert to kelvin.
- Assuming hydrogen bonding requires a bond to hydrogen only, when it must be to N, O, or F.
- Ranking molecules by molar mass alone when hydrogen bonding dominates, as with water versus hydrogen sulfide.
- Confusing vapour pressure with total pressure, or forgetting that boiling occurs when vapour pressure equals external pressure.
Worked example
Rank H2O, H2S, and H2Se by boiling point. All are bent and polar, but only water engages in hydrogen bonding, so it boils highest. Between H2S and H2Se, dispersion forces dominate; H2Se is larger and more polarisable, so it boils higher than H2S. Order: H2O > H2Se > H2S.
Rule of thumb
Identify the strongest force present first; then break ties with molecular size and shape.