Why Do Real Gases Deviate From Ideal Behavior?

Real gases differ from ideal behavior because, when set in low temperatures and high pressures, real gases defy the two assumptions of the kinetic molecular theory. Dutch physicist Johannes van der Waals was the first to develop an explanation for the real gas deviations.

The kinetic molecular theory has two assumptions for real gases that cause problems in low temperatures and high pressures (as in real gases deviate from this idea behavior). The kinetic molecular theory assumes that gas particles will only take up a tiny fraction of the total volume of the gas. Secondly, the theory assumes that the gas molecules will have no attraction.

The first assumption is only valid at pressures that are roughly one atm. However, when gas compression increases the pressure, that assumption no longer works; the real gas volume becomes larger than the ideal gas equation anticipates.

The second assumption is invalid because if there were no attraction between gas particles, that gas could never become a liquid, which would require it to condense. In reality, a minuscule force of attraction exists, keeping the molecules together. When temperatures fall, real gases become liquids, defying the assumptions of ideal behavior.