Which of the following describes real gases more accurately compared to ideal gases?

Real gases differ from ideal gases in several significant ways that stem from the assumptions of the ideal gas law, which states that gas particles have no volume and do not experience any interactions or forces between them.

The correct choice, which notes that real gases experience intermolecular forces, highlights a crucial distinction. In reality, gas molecules are attracted or repelled by one another due to various intermolecular forces, such as van der Waals forces or dipole-dipole interactions. These attractions, especially noticeable under conditions of high pressure or low temperature, lead to deviations from ideal gas behavior. For instance, as gas particles come closer together, the effects of these intermolecular forces become significant, resulting in behavior that cannot be accurately described by the ideal gas equation.

Other considerations involve the volume of gas particles and their motion. Real gases do not have negligible volume; they occupy space and the volume of the gas particles themselves can contribute to overall behavior, especially at high pressures. Additionally, while ideal gases are thought to move in perfect motion with no interactions, real gas particles may collide with one another, causing energy loss and alterations in direction due to these interactions.

Thus, the presence of intermolecular forces is a defining characteristic that allows for a more accurate description of