1. Attractive and Repulsive Forces:
* Assumption: The kinetic-molecular theory assumes that gas molecules have negligible intermolecular forces.
* Reality: Real gas molecules do experience attractive forces (like van der Waals forces) at close distances, and repulsive forces at very short distances. These forces become significant at high pressures or low temperatures when molecules are closer together.
2. Non-Zero Molecular Volume:
* Assumption: The kinetic-molecular theory assumes that gas molecules have negligible volume compared to the volume of the container.
* Reality: Real gas molecules do have a finite volume. This volume becomes significant at high pressures when molecules are packed more closely together.
3. Non-Ideal Collisions:
* Assumption: The kinetic-molecular theory assumes that collisions between gas molecules are perfectly elastic, with no loss of energy.
* Reality: Real gas collisions can involve some energy loss due to the intermolecular forces. These forces can cause molecules to "stick" together for brief periods, affecting the energy transfer during collisions.
4. Non-Uniform Velocity Distribution:
* Assumption: The kinetic-molecular theory assumes that gas molecules have a uniform distribution of velocities at a given temperature.
* Reality: In real gases, the distribution of velocities deviates from the ideal Maxwell-Boltzmann distribution, especially at high pressures and low temperatures.
Consequences of these deviations:
* Real gases are more compressible than ideal gases: This is due to the attractive forces between molecules, which allow them to be packed more closely together.
* Real gases have different boiling points than ideal gases: Attractive forces between molecules affect the energy required to overcome these forces and enter the gas phase.
* Real gas behavior can deviate significantly from ideal gas laws: The ideal gas law (PV=nRT) is only an approximation for real gases, especially at high pressures and low temperatures.
When do these deviations become significant?
* High pressure: At high pressures, molecules are closer together, making intermolecular forces and molecular volume more significant.
* Low temperature: At low temperatures, molecules have less kinetic energy, making intermolecular forces more significant.
How to account for real gas behavior:
* Equations of state: Equations like the van der Waals equation and the Redlich-Kwong equation attempt to account for the deviations of real gases from ideal gas behavior by introducing correction factors for intermolecular forces and molecular volume.
In summary, while the kinetic-molecular theory provides a useful foundation for understanding gas behavior, real gases exhibit deviations from these ideal assumptions, especially at high pressures and low temperatures. These deviations are important to consider for accurate predictions of gas behavior in various applications.
