At what conditions do gases typically deviate from ideal behavior?

Gases typically deviate from ideal behavior at low temperatures and high pressures due to the interactions between gas molecules and the physical limitations imposed by their volume. Ideal gas behavior is based on the assumption that gas molecules do not interact with one another and occupy no volume. However, at low temperatures, the kinetic energy of gas molecules decreases, allowing intermolecular forces (such as van der Waals forces) to become significant. These forces can cause attraction or repulsion between particles, affecting the gas's pressure and volume.

Additionally, at high pressures, gas molecules are forced closer together, leading to an increased likelihood of intermolecular interactions and a significant effect of the volume that the gas particles occupy. Under these conditions, the assumptions of the ideal gas law, which considers gas molecules to be point-like with negligible volume, can no longer hold true, resulting in deviations from ideal behavior. Understanding these conditions helps in predicting the behavior of real gases in various situations, which is essential for applications in chemistry and engineering.