Final answer:
True, real gases behave differently from ideal gases at very high pressures and very low temperatures due to significant volume and noticeable intermolecular forces under these conditions.
Step-by-step explanation:
True. Very high pressures and very low temperatures cause real gases to behave differently than ideal gases. At high pressures and low temperatures, the two basic assumptions behind the ideal gas law—that gas molecules have negligible volume and that intermolecular interactions are negligible—are no longer valid. Under these conditions, the volume of the gas particles becomes significant, and the intermolecular forces become noticeable. As a result, a real gas deviates most from an ideal gas under these conditions and behaves more like an ideal gas at high temperatures and low pressures.
For understanding real gas behavior, it's important to note that real gases also approach ideal gas behavior more closely at higher temperatures. Conversely, at low temperatures, the kinetic energy of the particles decreases, emphasizing the effect of intermolecular forces. This causes increased deviation from ideal gas behavior, which is not well described by the ideal gas law under such conditions. The van der Waals equation is used to correct for these deviations by incorporating the actual volume of gas molecules and intermolecular attractive forces.