Revision Snapshot

PV = nRT with R = 0.0821 L·atm/(mol·K) or 8.314 J/(mol·K). T always in Kelvin.
Dalton's Law: p_i = x_i × P_total; each gas acts independently.
Graham's Law: r_{1}/r_{2} = √( $M_{2}$ / $M_{1}$ ); lighter gas diffuses FASTER.
Molecular speeds: v_rms > v_avg > v_mp (ratio 1.224 : 1.128 : 1). Mnemonic: RAM.
Formulas: v_rms = √(3RT/M); v_avg = √(8RT/πM); v_mp = √(2RT/M).
KE per mole = (3/2)RT — depends only on T, not on type of gas.
Van der Waals: (P + $an^{2}$ / $V^{2}$ )(V − nb) = nRT. 'a' = attraction, 'b' = size.
Z = PV/nRT: Z = 1 (ideal), Z < 1 (attraction), Z > 1 (repulsion/size). $H_{2}$ and He: always Z > 1.
Critical constants: T_c = 8a/(27Rb), P_c = a/(27 $b^{2}$ ), V_c = 3b. Ratio: P_cV_c/T_c = 3R/8.
Boyle temperature: T_B = a/(Rb); gas shows near-ideal behavior over wide P range.
Joule-Thomson effect: cooling on expansion below inversion temperature → basis of liquefaction.
Liquid state: vapour pressure ↑ with T; surface tension ↓ with T; viscosity ↓ with T.
Boiling point = T where vapour pressure = external pressure; lower at high altitude.
Intermolecular forces: London < dipole-dipole < H-bonding (weakest to strongest).
STP molar volume: 22.4 L/mol for any ideal gas at 0°C, 1 atm.