Under what conditions can gas be considered ideal?

The conditions under which a gas can be considered ideal are best described by the scenario where the gas is at very low pressure and at a temperature that is higher than the critical temperature. In this state, the gas particles are relatively far apart, minimizing the interactions between them, which aligns with the assumptions of the ideal gas law. At low pressure, the volume occupied by the gas molecules becomes negligible when compared to the overall volume of the gas, meaning that the size of the molecules does not significantly affect their behavior.

Additionally, being above the critical temperature ensures that the gas remains in a gaseous state even if the pressure increases, thus preventing condensation into a liquid or solid phase where intermolecular forces become significant. Under these conditions, the ideal gas assumptions hold true: molecules do not attract or repel each other significantly, and they occupy no volume themselves, allowing for predictions about their behavior based on temperature and pressure using the ideal gas law (PV=nRT).