How does the behavior of ideal gases differ from that of real gases at high pressures?

The behavior of ideal gases is defined by the Ideal Gas Law, which assumes that gas particles do not interact with each other and occupy no volume. This model holds true under low-pressure and high-temperature conditions where gas particles are far apart and interactions can be neglected. However, at high pressures, the assumptions of the Ideal Gas Law begin to break down.

As pressure increases, the volume occupied by gas molecules themselves becomes significant because they are forced closer together. Real gases experience intermolecular forces that affect their behavior, such as attraction or repulsion between particles. At high pressures, these forces become pronounced, leading to deviations from the predictions made by the Ideal Gas Law. Consequently, real gases do not behave as ideal gases under these conditions, resulting in phenomena such as condensation or increased pressure not corresponding simply to increased volume.

Therefore, the observation that ideal gases follow the gas laws precisely while real gases begin to deviate significantly at high pressures accurately captures the differences in behavior. It highlights the limitations of the ideal gas model in predicting the characteristics of real gases under such conditions.