# What is the reason behind the phenomenon of Joule-Thomson effect?

For an ideal gas there is no heating or cooling during an adiabatic expansion or contraction, but for real gases, an adiabatic expansion or contraction is generally accompanied by a heating or cooling effect. What is the reason behind such a phenomenon? Is it related to the property of real gases or is it something else?

In a reversible adiabatic expansion or compression, the temperature of an ideal gas does change.

In a Joule-Thompson type of irreversible adiabatic expansion (e.g., in a closed container), the internal energy of the gas does not change. For an ideal gas, its internal energy depends only on its temperature. So, for an irreversible adiabatic expansion of an ideal gas in a closed container, its temperature does not change. But, the internal energy of a real gas depends not only on its temperature but also on its specific volume (which increases in an expansion). So, for a real gas, its temperature changes. The Joule-Thompson effect is one measure of the deviation of a gas from ideal gas behavior.

Irrespective of the Joule-Thompson effect, one can show (using a combination of the first and second laws of thermodynamics) that, for a pure real gas, liquid, or solid (or one of constant chemical composition), the variation of specific internal energy with respect to temperature and specific volume is given by: $$dU=C_vdT-\left[P-T\left(\frac{\partial P}{\partial T}\right)_V\right]dV$$The first term describes the variation with respect to temperature and the second term describes the variation with respect to specific volume. For an ideal gas, the second term is equal to zero. However, for a real gas, the second term is not equal to zero, and that means that, at constant internal energy (as in the Joule-Thompson effect), the temperature will change when the specific volume changes. This is a direct result of the deviation from ideal gas behavior.