Under which conditions are gases considered ideal?

Gases are considered ideal when they behave according to the ideal gas law, which assumes that the gas molecules do not interact with each other and occupy no volume. This ideal behavior is most closely approximated under conditions of high temperature and low pressure.

At high temperatures, the kinetic energy of the gas particles increases, which causes them to move more rapidly and reduces the effect of intermolecular forces. As temperature increases, gas molecules are less likely to condense into a liquid or experience attraction or repulsion from one another, aligning with the assumption of ideal gas behavior that their interactions are negligible.

Additionally, at low pressures, gas particles are spaced far enough apart that their volumes contribute very little to the overall behavior of the gas. The particles are less likely to collide with each other, and the effects of their physical size can be ignored, further supporting ideal gas assumptions.

In contrast, at low temperatures and high pressures, gases deviate from ideal behavior due to significant intermolecular attractions and the volume of the gas particles themselves becoming significant in comparison to the space they occupy. Thus, the combination of high temperature and low pressure creates the conditions under which gases behave most ideally.