- #1
JeweliaHeart
- 68
- 0
Hello. I am a thermodynamics novice trying to gain a better understanding of state functions, particularly enthalpy.
I understand that enthalpy is defined as
"A measure of the total energy of a thermodynamic system, including internal energy, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure."
The equation:
ΔH=ΔU(internal energy) + ΔPV
confuses me b/c
ΔU= q(heat added) - w(work done by system on environment)
so
ΔH really means:
ΔH=q - w + ΔPV
There are two terms of work (w and ΔPV) and b/c of the opposite sign, they cancel out, leaving only q. This means ΔH= q which is at odds with the accepted definition of enthalpy. Where did I mess up?
I understand that enthalpy is defined as
"A measure of the total energy of a thermodynamic system, including internal energy, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure."
The equation:
ΔH=ΔU(internal energy) + ΔPV
confuses me b/c
ΔU= q(heat added) - w(work done by system on environment)
so
ΔH really means:
ΔH=q - w + ΔPV
There are two terms of work (w and ΔPV) and b/c of the opposite sign, they cancel out, leaving only q. This means ΔH= q which is at odds with the accepted definition of enthalpy. Where did I mess up?