Under what conditions of temperature and pressure is the behavior of real gases differ most from that of ideal gases?
Do real gases liquify? Why?
Do real gases have attractive forces between molecules? How does that affect low pressures?
The behavior of real gases differs most from that of ideal gases under conditions of high pressure and low temperature.
To understand why, we need to know the assumptions behind the behavior of ideal gases. Ideal gases are modeled based on several assumptions, including that the gas particles are point masses with no volume, that they do not attract or repel each other, and that there are no intermolecular forces between the particles.
In reality, real gases do have volume and experience intermolecular forces. At high pressures, the particles are forced closer together, and their volumes become significant compared to the overall volume of the gas. This means that the particles occupy a larger volume than predicted by the ideal gas law.
Additionally, at low temperatures, the kinetic energy of the gas particles decreases. As a result, intermolecular forces, such as Van der Waals forces, become more significant. These attractive forces cause the gas particles to be closer together than predicted by the ideal gas law.
Therefore, under conditions of high pressure and low temperature, the volume occupied by gas particles and the intermolecular forces between them have a significant impact on the behavior of real gases, making them differ most from ideal gases.