The compressibility factor $Z$ of a gas is defined as $Z = \dfrac{PV}{nRT}$. For an ideal gas, $Z$ equals:

For an ideal gas, $PV = nRT$, so $Z = \dfrac{PV}{nRT} = 1$ at all $P$ and $T$. Real gases deviate from $Z = 1$: - $Z < 1$ at intermediate pressures (attractive intermolecular forces dominate, making the gas more compressible than ideal) - $Z > 1$ at high pressures (repulsion dominates, making the gas less compressible) The departure of $Z$ from 1 is the standard measure of how non-ideal a gas is. CO₂ and NH₃ deviate strongly at lower temperatures (they liquefy easily), while He and H₂ stay close to ideal across most conditions.