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:
To understand why this is the correct statement, we need to compare the behavior of real gases with ideal gases and consider how pressure influences each.
Ideal Gases vs. Real Gases:
- Ideal gases are theoretical gases that follow the ideal gas law perfectly, represented by the equation:
[
PV = nRT
]
where: - ( P ) is pressure,
- ( V ) is volume,
- ( n ) is the number of moles,
- ( R ) is the gas constant,
- ( T ) is temperature. Ideal gases are assumed to have no interactions between molecules and their volume is considered negligible.
- Real gases, on the other hand, deviate from the ideal gas law due to intermolecular forces and the finite volume of gas molecules. These deviations become more pronounced under certain conditions, like high pressures and low temperatures.
Pressure and Real Gas Behavior:
- At low pressure, the volume of a real gas is close to that predicted by the ideal gas law. This is because, at low pressure, gas molecules are far apart, and intermolecular forces (such as attraction or repulsion) have a minimal effect. Therefore, a real gas behaves similarly to an ideal gas.
- At high pressure, gas molecules are forced closer together, and their volume becomes significant compared to the volume of the container. The molecules experience intermolecular attractions, which cause real gases to behave differently from ideal gases. Real gases tend to have a larger volume than ideal gases at high pressure because:
- The volume occupied by gas molecules becomes noticeable.
- Attractive forces between molecules cause them to occupy more space than predicted by the ideal gas law.
In contrast, at extremely high pressures, repulsive forces between molecules may also start to dominate, leading to further deviations from ideal behavior. But generally, the volume of real gases is greater than that of ideal gases at high pressures.
Therefore, the correct statement emphasizes that real gases, at high pressures, will take up more volume due to intermolecular forces and the finite size of the molecules.