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Under normal conditions, water vapor *deviates* from the perfect gas model. This deviation is more pronounced when: the pressure is high and the temperature is low. That's because, under high pressures, the water molecules are forced closer together, and their size matters more. At low temperatures, the molecules slow down, and the attractive forces between them become more significant. Think of it like a crowd. When everyone is spread out (low pressure), you can treat each person as an individual without worrying too much about interactions. But when they're packed tight (high pressure), their individual sizes and how they interact become more important. Under these conditions, the ideal gas law starts to become less accurate. We often need to use more complex equations, like the *Van der Waals equation*, to accurately model the behavior of water vapor. This equation takes into account both the volume of the molecules and the attractive forces between them. So, the closer you get to the conditions where water might condense into a liquid, the more you have to account for these real-world effects.
