Which of the following correctly states the impact of pressure on the behavior of real gases compared to ideal gases?
At low pressure, real gas particles do not exert any force on the container, so they exert less pressure than ideal gases.
At high pressure, real gas particles take up more of the available space, so their volume is greater than ideal gases.
At low pressure, real gas particles expand to fill the container so their volume cannot be measured compared to ideal gases.
At high pressure, real gas particles are squeezed together, so they take up less volume than ideal gases.
The correct answer and explanation is :
The correct statement is:
At high pressure, real gas particles take up more of the available space, so their volume is greater than ideal gases.
Explanation:
The behavior of gases can be understood using the Ideal Gas Law, which assumes that gas particles do not interact with one another and occupy no volume. However, real gases deviate from this ideal behavior, especially under extreme conditions like high pressure or low temperature.
- Ideal Gas Law Assumptions:
The Ideal Gas Law (PV = nRT) assumes that:
- Gas molecules have no intermolecular forces (no attraction or repulsion).
- Gas molecules occupy no volume.
- The gas molecules are in constant random motion and perfectly elastic collisions occur. In ideal gas behavior, the volume of the gas is directly proportional to the temperature and inversely proportional to the pressure.
- Real Gas Behavior at High Pressure:
- Volume of real gas molecules: Real gas molecules do occupy a finite volume. As the pressure increases, gas molecules are forced closer together. In this situation, the volume occupied by the gas particles themselves becomes significant.
- Intermolecular forces: At high pressure, the attractive forces between gas molecules also become more pronounced. These forces lead to a reduction in the effective volume of the gas compared to the ideal gas prediction, making the gas behave differently than predicted by the ideal gas law.
- Deviations from ideal behavior: The increased volume occupied by the molecules (due to their finite size) and the attractive intermolecular forces lead to a situation where the real gas takes up more space than the ideal gas would under the same conditions. This is why real gases have a larger volume than ideal gases at high pressures.
In contrast, at low pressures, gas molecules are far apart, and the intermolecular forces are minimal, so the behavior of real gases is closer to that of ideal gases. Thus, at high pressures, the volume of real gases is typically greater than that predicted by the ideal gas law.
This is why the statement that “at high pressure, real gas particles take up more of the available space, so their volume is greater than ideal gases” is correct.