When does a gas behave like an ideal gas?

This is a concept that confuses many students, but once you understand the logic behind it, the answer becomes completely obvious. A real gas behaves like an ideal gas when the conditions allow its molecules to move freely without significantly interacting with each other. At low pressure, gas molecules are far apart from each other, so intermolecular forces of attraction become negligible. At high temperature, molecules move so fast and with such high kinetic energy that they simply do not have time to attract each other — they just zoom past. Together, these two conditions make a real gas behave as close to an ideal gas as possible.

• Low pressure: gas molecules are spread far apart, making intermolecular attractions negligible — one of the two key assumptions of ideal gas behaviour
• High temperature: molecules possess very high kinetic energy and move extremely fast, overpowering any intermolecular forces between them
• Ideal gas assumptions: an ideal gas assumes zero intermolecular forces and negligible molecular volume — both conditions are best satisfied at low pressure and high temperature
• Real gas deviation: at high pressure, molecules are forced close together and intermolecular forces dominate — behaviour deviates significantly from ideal
• Van der Waals equation: this equation was specifically developed to correct for the non-ideal behaviour of real gases at high pressure and low temperature