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[subdude]

I've got a thermodynamics problem hopefully someone can explain for me.

On a vapor (non-ideal gas) enthalpy-entropy diagram, such as a Mollier diagram for steam, specific enthalpy peaks at some pressure for the saturated vapor. Beyond a certain point, the specific enthalpy (as well as internal energy) of the saturated vapor decreases with increasing pressure (and the increasing corresponding saturation temperature). Could someone explain conceptually and/or analytically why this occurs? Why does the enthalpy and internal energy of a pound of saturated vapor increase with pressure up to a point, peak, and then decrease with increasing pressure?

Looking at it another way---when saturated vapor is superheated at constant pressure, its enthalpy and internal energy increase. This is, of course, true for any pressure. However, for a superheated vapor (steam, at least), when a superheated vapor is compressed at constant temperature, its internal energy and enthalpy decrease. This happens regardless of the initial pressure and temperature of the superheated steam. The decreases with increasing pressure may be small, especially at low pressures, but at higher pressures, the decrease in internal energy and enthalpy with increasing pressure and constant temperature becomes very significant. What is going on with the vapor that makes it decrease in both internal energy and enthalpy as pressure increases and temperature remains constant? (In contrast, neither internal energy nor enthalpy change for pressure changes with an isothermal ideal gas, given that the heat and work magnitudes are equal and p * v is constant.)

Thanks.

[MrShorty]

Interesting question. I'm not quite sure how to answer it, but I wonder if it has something to do with the fact that enthalpy of vaporization has to equal zero at the critical. This means that, as T->Tc, Hvap->Hliq.