For a van der Waals gas, there are London dispersion forces causing the gas particles to be attracted to each other while the law of excluded volume provides repulsive forces. How do each of these characteristics effect the pressure, volume, and temperature of the gas? I assumed that for the attractive intermolecular forces, pressure and volume would decrease because the particles would clump together (smaller volume) and hit the sides of their container less (less pressure). At the same time, I think that temperature would increase because the molecules would have a greater kinetic energy from bumping into each other. Thus, I believe the opposite would be the case for the repulsive forces. Is this actually the case?