Final answer:
Real gas molecules have finite volumes and experience intermolecular forces, which become significant under high pressure and lower the pressure compared to what is predicted by the ideal gas law.
Step-by-step explanation:
Real gases exhibit a lower pressure than ideal gases at high pressures due to two main reasons. Firstly, the molecules in a real gas have finite volumes which become significant under high pressure, causing fewer collisions with container walls.
Secondly, at high pressures, intermolecular forces of attraction within the real gas reduce the number of collisions with the container walls, thereby lowering the pressure compared to what is predicted by an ideal gas.
Boyle's law, which states that pressure is inversely proportional to volume, does not hold true for real gases at high pressures due to these factors.
Therefore, the behavior of real gases deviates from the ideal gas law (PV=nRT), where P is pressure, V is volume, n is amount of gas, R is the ideal gas constant, and T is temperature.