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FY13CE Chemistry study materials
FY13CE Chemistry study resources

FY13CE Chemistry: States of Matter and Intermolecular Forces

FY13CE Chemistry · Work through your exam, one topic at a time.

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States of Matter and Intermolecular Forces

This domain tests the gas laws, the assumptions and failures of the ideal-gas model, phase behaviour, and the ranking of melting and boiling points from the strength of intermolecular forces.

Gas behaviour

  • Ideal gas equation PV = nRT with P in Pa, V in m3, T in kelvin and R = 8.314 J mol−1 K−1.
  • Dalton's law: total pressure is the sum of partial pressures; each partial pressure is mole fraction × total pressure.
  • Kinetic theory assumes negligible molecular volume, no intermolecular forces, and elastic collisions. Real gases deviate most at high pressure and low temperature, when molecules are close together.
  • van der Waals equation (P + an²/V²)(V − nb) = nRT corrects for attraction (a) and molecular volume (b).

Intermolecular forces

  • London dispersion forces act between all particles and strengthen with more electrons and greater surface contact; they are the only force in non-polar molecules.
  • Dipole–dipole forces act between permanent dipoles. Hydrogen bonding is the strong special case requiring H directly bonded to N, O or F.
  • Ion–dipole forces operate in solutions of ionic compounds.
  • Stronger intermolecular forces give higher melting and boiling points, larger enthalpy of vaporisation and lower vapour pressure.

Phase behaviour

  • Vapour pressure rises with temperature; a liquid boils when its vapour pressure equals the external pressure.
  • A phase diagram shows the triple point and critical point; water's solid–liquid line slopes backwards, so ice melts under pressure.

Worked example

0.575 g of a volatile liquid gives 200 cm3 of vapour at 100 °C and 101 kPa. n = PV/RT = (101000 × 2.00 × 10−4)/(8.314 × 373) = 20.2/3101 = 6.51 × 10−3 mol. M = 0.575/6.51 × 10−3 = 88.3 g mol−1, consistent with a small ester such as C4H8O2.

Traps that cost marks

  • Substituting degrees Celsius or cubic centimetres directly into PV = nRT.
  • Claiming hydrogen bonding whenever hydrogen is present; it needs an N–H, O–H or F–H bond.
  • Forgetting that HF boils above HCl despite being lighter, because of hydrogen bonding.
  • Assuming ice is denser than water; the open hydrogen-bonded lattice makes it less dense.
  • Ignoring that dispersion forces increase with chain length and branching reduces them.