I have been struggling to understand how the molecules behave during the Joule–Thomson effect. I would love to get some help on this concept. Here is what I got so far.
Under adiabatic and isenthalpic conditions, when a gas expand from high pressure to low pressure, depending on the initial pressure, initial temperature, and the molecular properties, the gas can either experience cooling or heating.
Specifically, if the gas was initially in the state where attractive intermolecular forces were dominant (typically at regions of moderate temperature and low pressure), as it expand, the kinetic energy of molecules are used to overcome the attraction reach greater separation. Also, the internal energy is overall decreased as the gas need to do work to expand. Thus the temperature of the gas decreases.
If repulsive forces were dominant (at regions of very high temperature or high pressure), as the gas expand, the repulsion kind of pushes(?) the molecules to reach greater distance. In other words, the potential energy is converted to kinetic energy here. And here although internal energy may be decreased to provide work for expansion, the molecule would have gained more kinetic energy from the repulsive forces such that the temperature of the molecules increases.
My question is, how does the system being isenthalpic play a role in this effect? I get that we need the adiabatic condition so that only the internal energy is being used to provide the work required for expansion. But what about enthalpy? How is isenthalpic + adiabatic expansion different from just adiabatic expansion?
Additionally, I would appreciate if someone could also explain how $pV$ from the enthalpy formula $H=U+pV$, changes during this expansion.