$word_{count}$: 220

Real gases deviate from ideal behaviour due to: (1) non-zero molecular volume, (2) intermolecular forces. Deviations increase at high P and low T. van der Waals equation: (P + $an^2$/$V^2$)(V - nb) = nRT. 'a' corrects for intermolecular attraction (units: $L^2$.atm/$mol^2$). Higher a = stronger attraction = easier liquefaction. Water: a = 5.46, CO2: a = 3.59, H2: a = 0.244. 'b' corrects for molecular volume (units: L/mol). b ≈ 4 x actual molecular volume per mole. Compressibility factor Z = $\frac{PV}{nRT}$. Ideal: Z = 1. At moderate P: Z < 1 (attraction dominates, gas occupies less volume than ideal). At very high P: Z > 1 $\frac{repulsion}{volume exclusion dominates}$. At Boyle temperature $T_B$ = $\frac{a}{Rb}$: Z ≈ 1 over wide P range (gas behaves most ideally). H2 and He have very small 'a' values: Z > 1 at almost all pressures (repulsion always dominates). All gases approach Z = 1 at very low P (ideal behaviour). The plot of Z vs P at different temperatures is frequently tested in JEE.