When does a real gas deviate from ideal behavior?

A real gas deviates from ideal behavior primarily at low temperatures and high pressures. Under these conditions, the assumptions that characterize an ideal gas—such as the gas particles having no volume and experiencing no intermolecular forces—begin to break down.

At high pressures, the gas particles are forced closer together, making their actual volume significant compared to the volume of the container. This leads to deviations from ideal behavior since the space occupied by gas molecules cannot be ignored.

At low temperatures, the kinetic energy of the gas particles decreases, which allows intermolecular forces (like Van der Waals forces) to become more pronounced. These forces cause the gas to condense into a liquid state or just not behave like an ideal gas, which assumes that these forces are negligible.

In contrast, at high temperatures and low pressures, gases behave more ideally because the increased kinetic energy of the particles overcomes any intermolecular forces, and the distances between particles are large enough that their volumes become less significant. Hence, real gases are observed to follow the ideal gas laws more closely under such conditions.