Here's a breakdown of why:
* Ideal gas assumptions: Ideal gas theory assumes that gas molecules have no volume and don't interact with each other. This simplifies calculations but is not completely accurate in real-world scenarios.
* Real gas volume: Real gas molecules, while very small, do occupy a finite volume. This means the space available for them to move around is slightly less than the total volume of the container.
* Intermolecular forces: Real gas molecules attract each other, especially at higher pressures and lower temperatures. These attractive forces, like van der Waals forces, cause the molecules to deviate from the ideal gas behavior where they are assumed to be independent.
In summary: The finite volume and intermolecular forces of real gas molecules cause deviations from the ideal gas law, especially at high pressures and low temperatures.
