In terms of temperature and pressure, when do gases behave ideally?

Gases behave ideally under conditions of high temperatures and low pressures. At high temperatures, the kinetic energy of gas molecules increases, which allows them to move more freely and minimizes interactions between particles. This is crucial because ideal gas behavior assumes that gas molecules do not attract or repel each other, which is more plausible at higher energies.

Additionally, low pressures reduce the density of the gas, which means that the volume of the gas particles in relation to the empty space increases. Under such conditions, the assumptions of the ideal gas law—like negligible volume of gas particles and no intermolecular forces—are more likely to hold true. This combination effectively allows the gas molecules to behave like an ideal gas, where the behavior can be accurately described by the ideal gas law, PV=nRT.

In contrast, low temperatures tend to slow down the molecules and may cause them to cluster together due to intermolecular forces, while high pressures compress the gas, making molecular volume significant and increasing the influence of these intermolecular forces. This is why gases do not behave ideally under those conditions.