9. Liquids, Solids & Intermolecular Forces
Intermolecular forces (IMFs) — the attractions between molecules — decide boiling points, vapour pressures, and why some solids melt in your hand while others need a furnace.
Types of Intermolecular Forces
- London dispersion: present in all molecules; grows with size and surface area.
- Dipole–dipole: between polar molecules.
- Hydrogen bonding: H attached to N, O, or F attracted to a lone pair on N, O, or F — the strongest IMF, and why water is a liquid at room temperature.
Example: HF boils far higher than HCl despite being lighter — HF hydrogen-bonds; HCl does not.
Vapour Pressure and Boiling Point
Stronger IMFs mean lower vapour pressure and a higher boiling point. A liquid boils when its vapour pressure equals the external pressure.
Phase Changes and Heating Curves
Melting and boiling absorb heat without changing temperature: $q = m\Delta H_{\text{fus}}$ and $q = m\Delta H_{\text{vap}}$.
Example: Heat 50 g of ice at −10 °C to water at 20 °C ($c_{\text{ice}} = 2.09$, $c_{\text{water}} = 4.184$ J/g·°C, $\Delta H_{\text{fus}} = 334$ J/g):
$$q = (50)(2.09)(10) + (50)(334) + (50)(4.184)(20) = 1045 + 16700 + 4184 = 21.9\ \text{kJ}$$
Practice Problems
- Which has stronger intermolecular forces: CH4 or NH3? Why?
- Why does ethanol boil higher than dimethyl ether, its isomer?
- How much heat melts 100 g of ice at 0 °C?
- Which has the higher vapour pressure at 25 °C: water or diethyl ether?
Answers: 1. NH3 — it hydrogen-bonds; CH4 has only dispersion forces 2. Ethanol hydrogen-bonds; dimethyl ether cannot 3. 33.4 kJ 4. Diethyl ether — weaker IMFs mean more molecules escape