Explain: Why do real gases show deviation from ideal gas behavior?

(N/A) Real gases deviate from ideal behavior because two fundamental assumptions of the kinetic molecular theory are not strictly valid for real gases:
$1$. The assumption that there is no force of attraction between gas molecules is incorrect. If this were true,gases would never liquefy. The fact that gases liquefy upon cooling and compression indicates that intermolecular forces of attraction exist.
$2$. The assumption that the volume of gas molecules is negligible compared to the total volume of the gas is incorrect. At high pressures,the volume occupied by the molecules themselves becomes significant relative to the total volume,leading to deviations from the ideal gas law $(PV = nRT)$.