The Van der Waals equation is an equation of state that extends the ideal gas law to include the effects of interaction between molecules of a gas, as well as accounting for the finite size of the molecules. It was derived by Johannes Diderik van der Waals in 1873 and is a modified version of the ideal gas law. The equation relates the relationship between the pressure, volume, temperature, and amount of real gases. For a real gas containing ‘n’ moles, the equation is written as:

$$\left(P+\frac{an^2}{V^2}\right)\left(V-nb\right)=nRT$$

Where P, V, T, and n are the pressure, volume, temperature, and moles of the gas, respectively, and ‘a’ and ‘b’ are constants specific to each gas. The equation can be used to approximate the behavior of real fluids above their critical temperatures and is qualitatively reasonable for their liquid and low-pressure gaseous states at low temperatures. However, near the phase transitions between gas and liquid, in the range of p, V, and T where the liquid phase and the gas phase are in equilibrium, the Van der Waals equation fails to accurately model observed experimental behavior. The constants ‘a’ and ‘b’ have positive values and are characteristic of the individual gas. The constant ‘a’ provides a correction for the intermolecular forces, while constant ‘b’ is a correction for finite molecular size and its value is the volume of one mole of the atoms or molecules.