How does the pressure of a real gas compare to the pressure of an ideal gas at low temperatures?

• A. The pressure of a real gas is lower than that predicted by the ideal gas law.
• B. The pressure of a real gas cannot be compared to that of an ideal gas at low temperatures.
• C. The pressure of a real gas is equal to that predicted by the ideal gas law.
• D. The pressure of a real gas is higher than that predicted by the ideal gas law.

Answer: (A)

Verified step by step guidance

1. Recall that the ideal gas law assumes no intermolecular forces and that gas particles occupy no volume, which is a good approximation at high temperatures and low pressures.
2. At low temperatures, gas particles move more slowly, so intermolecular attractive forces become significant, causing the particles to pull closer together.
3. Because of these attractive forces, the gas particles collide with the container walls with less force than predicted by the ideal gas law, leading to a lower observed pressure.
4. Therefore, the pressure of a real gas at low temperatures is generally lower than the pressure predicted by the ideal gas law due to these intermolecular attractions.
5. This behavior is often described and corrected by the van der Waals equation, which adjusts for particle volume and intermolecular forces.