What is the primary difference between an ideal gas and a real gas?

The primary distinction between an ideal gas and a real gas lies in their adherence to gas laws under varying conditions. An ideal gas is a theoretical construct that perfectly follows the gas laws—such as Boyle's Law, Charles's Law, and Avogadro's Law—regardless of temperature or pressure. In contrast, real gases do not maintain this behavior under all conditions due to the influence of intermolecular forces and the finite volume that gas particles occupy.

At high pressures or low temperatures, real gases exhibit deviations from ideal behavior because the attractive or repulsive forces between gas molecules become significant. These forces can lead to reduced pressure or volume, as real gas molecules can clump together or compress more than what would be predicted by an ideal gas model. Therefore, the statement accurately reflects the fundamental distinction in the behavior of ideal versus real gases and highlights the limitations of the ideal gas law in practical applications.