Gases, Liquids, Solids and Phase Changes
Gas laws
Ideal gas equation: PV = nRT with R = 0.0821 L atm mol-1 K-1 (or 8.314 J mol-1 K-1). Temperature must be in kelvin. Combined law: P1V1/T1 = P2V2/T2 for a fixed amount. Dalton: the total pressure of a mixture equals the sum of partial pressures, and the mole fraction of a gas equals its partial pressure divided by the total. Kinetic molecular theory explains pressure as collisions of particles with the walls; real gases deviate at high pressure and low temperature, where particle volume and attractions matter.
Phase changes
Vapour pressure rises with temperature and boiling occurs when it equals the external pressure - which is why water boils below 100 °C on a mountain. A heating curve has sloping sections (sensible heat, q = mcΔT) and flat sections at melting and boiling (latent heat, q = nΔH). On a phase diagram the triple point is where all three phases coexist and the critical point ends the liquid-vapour boundary.
Solids
Ionic (high melting, brittle, conducts when molten), molecular (soft, low melting, held by intermolecular forces), covalent network such as diamond or quartz (very high melting, hard, nonconducting), metallic (conducts, malleable).
Worked example
0.50 mol of gas occupies 12.0 L at 27 °C. Convert: T = 300 K. Then P = nRT/V = (0.50 x 0.0821 x 300)/12.0 = 12.32/12.0 = 1.03 atm. Note that using 27 instead of 300 would give a meaningless answer.
Traps
Using degrees Celsius in gas calculations. Treating the flat parts of a heating curve as if the temperature rises. Assuming that vapour pressure depends on the amount of liquid - it depends only on temperature and on the substance.