1. Why Kinetic Theory Has Limitations
Kinetic theory provides a simple model of gas molecules, assuming they are point-sized particles with no forces between them. This works well for ideal gases but not always for real gases. When real conditions differ, the theory cannot fully explain gas behaviour.
2. Ideal Assumptions vs Real Behaviour
Kinetic theory is built on assumptions that are not always true. These assumptions break down under several conditions.
2.1. 1. Assumption of Zero Intermolecular Forces
Kinetic theory assumes molecules do not attract or repel each other. But real gases experience significant attractive forces at low temperature and high pressure.
2.2. 2. Assumption of Negligible Molecular Volume
The theory treats molecules as point masses with no size. But real molecules occupy space, especially noticeable at high pressure when they are squeezed close together.
2.3. 3. Perfectly Elastic Collisions
Kinetic theory assumes no energy is lost during collisions. In real gases, some energy can be transferred to rotational or vibrational modes, slightly deviating from perfect elasticity.
3. Where Kinetic Theory Cannot Predict Behaviour Accurately
Certain situations make the ideal assumptions invalid, revealing limitations of the theory.
3.1. 1. High Pressure
At high pressure, molecules come very close together. Their finite size becomes important, increasing the actual volume compared to the ideal prediction.
3.2. 2. Low Temperature
At low temperature, molecules move slowly. Attractive forces become significant, causing pressure to drop below ideal values.
3.3. 3. Near Liquefaction
As a gas approaches liquefaction, intermolecular forces dominate. Molecules form clusters, which cannot be explained using kinetic theory.
4. Failure to Explain Real Gas Heat Capacities
Kinetic theory assumes that heat capacity depends only on translational motion and degrees of freedom. But experimental results show that:
- heat capacities vary with temperature
- vibrational modes become active at higher temperature
The theory oversimplifies energy distribution.
4.1. Example
Diatomic gases have increasing heat capacities at high temperatures due to activation of vibrational modes, which kinetic theory cannot predict alone.
5. Failure to Explain Real Gas Deviations
Kinetic theory cannot explain:
- why real gases have a compressibility factor Z ≠ 1
- why real gases condense into liquids
- why intermolecular attraction affects pressure
5.1. Correction Needed
Modified equations such as the van der Waals equation are needed to describe real gas behaviour accurately.
6. Why Kinetic Theory Still Matters
Even though the theory has limitations, it is extremely important because it provides simple, intuitive explanations for gas pressure, temperature, and molecular motion. Most gases behave almost ideally under normal conditions, making kinetic theory very useful.
6.1. Simple Understanding
Kinetic theory is a good approximation for everyday conditions but not for extreme situations like very high pressure or very low temperature.