- #1
Oliver321
- 59
- 5
Hello everyone!
I have a course in thermodynamics this year, and there is a question about enthalpy that I cannot answer: given the definition of enthalpy H=U+PV and the integral form of the internal energy U=TS-PV we conclude that H=TS.
We normally say that enthalpy equals the heat exchanged in a isobaric processes. But where does the pressure appear in this equation? Related to this: the differential form of H is said to be dH=TdS+Vdp arising from dH=d(U+PV)=dU+d(PV). But if I do the same with the formula above dH=d(TS)=TdS+SdT I get a different result. How can this be?
And at least: We know that U=Q+W and U=TS-PV. Does this (in this case) mean, that Q=TS and W=-PV?
Thanks for every helping answer!
I have a course in thermodynamics this year, and there is a question about enthalpy that I cannot answer: given the definition of enthalpy H=U+PV and the integral form of the internal energy U=TS-PV we conclude that H=TS.
We normally say that enthalpy equals the heat exchanged in a isobaric processes. But where does the pressure appear in this equation? Related to this: the differential form of H is said to be dH=TdS+Vdp arising from dH=d(U+PV)=dU+d(PV). But if I do the same with the formula above dH=d(TS)=TdS+SdT I get a different result. How can this be?
And at least: We know that U=Q+W and U=TS-PV. Does this (in this case) mean, that Q=TS and W=-PV?
Thanks for every helping answer!