Final answer:
Real gases deviate from ideal behavior primarily due to their molecules occupying non-zero volumes and exerting intermolecular forces, which become significant under high pressures and low temperatures.
Step-by-step explanation:
The two properties of real gases that cause deviation from ideal behavior are the non-zero volume of gas molecules and the intermolecular forces between the molecules. Ideal gases assume that gas particles have no volume and experience no intermolecular attraction or repulsion, but this is not true for real gases.
At high pressures and low temperatures, these assumptions no longer hold, leading to significant deviations from ideal behavior. Real gases do, however, tend to behave more like ideal gases at higher temperatures and lower pressures.
The van der Waals equation introduces corrections for these real gas properties, addressing the volume occupied by the gas molecules and the intermolecular forces that affect pressure.