Inversion temperature

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The inversion temperature in thermodynamics and cryogenics is the critical temperature below which a non-ideal gas (all gases in reality) that is expanded at constant enthalpy will experience a temperature decrease, and above which will experience a temperature increase. This temperature change is known as the Joule-Thomson effect, and is exploited in the liquefaction of gases.

1 Theory
2 See also
3 References
4 External links

[edit] Theory

The Joule-Thomson effect cannot be described in the theory of ideal gases, in which interactions between particles are ignored. Instead, one must use a theory that accounts for the Van der Waals force between interacting particles that becomes much stronger as a gas becomes a liquid.

For a van der Waals gas we can calculate the enthalpy H using statistical mechanics as

$H = \frac{5}{2} N k_B T + \frac{N^2}{V} (b k_B T - 2a)$ ,

where $N$ is the number of molecules, $V$ is volume, $T$ is temperature (in the Kelvin scale), $k B$ is Boltzmann's constant, and $a$ and $b$ are constants depending on intermolecular forces and molecular volume, respectively.

From this equation, we note that if we keep enthalpy constant and increase volume, temperature must change depending on the sign of $b k B T - 2 a$ . Therefore, our inversion temperature is given where the sign flips at zero, or

$k_B T_\textrm{inv} = 2a / b = \frac{27}{4} k_B T_c$ ,

where $T c$ is the critical temperature of the substance. So for $T > T inv$ , an expansion at constant enthalpy increases temperature as the work done by the repulsive interactions of the gas is dominant, and so the change in energy is negative. But for $T < T inv$ , expansion causes temperature to decrease because the work of attractive intermolecular forces dominates, giving a positive change in energy.^[1]