1. Intermolecular forces: Ideal gases are assumed to have no interactions between their molecules. However, real gases exhibit weak attractive forces, known as van der Waals forces. These forces arise from temporary fluctuations in electron distribution around molecules, leading to temporary dipoles that attract neighboring molecules. This attraction reduces the pressure exerted by the gas compared to what would be expected from the ideal gas law.
2. Finite volume of gas molecules: Ideal gases are assumed to have zero volume. In reality, molecules do occupy a finite volume. This means that the free space available for the molecules to move around in is less than the total volume of the container. This reduction in available volume increases the pressure exerted by the gas, compared to what would be expected from the ideal gas law.
These two factors, intermolecular forces and finite molecular volume, are accounted for in the van der Waals equation, which provides a more accurate description of real gas behavior than the ideal gas law.
