Van der Waals

The deviation of real gas behavior from ideal gas is explained by Van der Waals equation. This equation takes into account the volume occupied by gas molecules and the attractive forces between them, which are ignored in the ideal gas law.

Van der Waals Equation

The Van der Waals equation corrects the ideal gas law by adding two correction factors:
1. The first correction factor accounts for the volume occupied by gas molecules. In an ideal gas, the volume of gas molecules is considered negligible, but in reality, gas molecules have a finite volume which reduces the available volume for the gas to occupy.
2. The second correction factor adjusts for the attractive forces between gas molecules. In ideal gases, it is assumed that there are no intermolecular forces present, but in reality, gas molecules do attract each other, causing the gas to deviate from ideal behavior.

Significance

The Van der Waals equation is crucial in understanding real gas behavior, especially at high pressures and low temperatures where gas molecules are closer together and the effects of molecular volume and intermolecular forces are more pronounced. This equation helps to predict the behavior of gases more accurately under these conditions compared to the ideal gas law.

In conclusion, Van der Waals equation plays a significant role in explaining the deviations of real gases from ideal gases by considering factors like molecular volume and intermolecular forces, which are neglected in the ideal gas law.