The ideal gas equation has two significant drawbacks: first, it does not consider the volume of gas molecules, and second, it does not account for attractive intermolecular forces between the gas molecules. If the gas contains n number of moles and each one of them occupies volume b, then the total volume occupied by the gas molecules is nb. Thus, the volume available for gas molecules to move will be total volume minus nb. The attractive intermolecular forces reduce the pressure proportional to the square of the molar density. In the absence of attractive intermolecular forces, higher pressure is required to reach the same nRT. This is the van der Waals equation of state that tries to address the drawbacks of the ideal gas equation. Here, a and b are empirical constants that are dependent on the type of gas under study. The van der Waals equation predicts the behavior of real gas adequately, such as liquid to vapor transition and the Joule-Thomson effect.