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Enthalpy's physical interpretation

  1. Oct 9, 2014 #1
    Hi I'm new to this site, so I'm not sure where to post this question but chemistry seemed fitting.

    I'm current taking 2 thermo classes, engineering and chemistry, and I'm having trouble with the concept of enthalpy.

    This is what I believe I understand. I Hope the following is all correct if not please correct me.
    H= U + PV
    The heat exchanged in a closed system under constant pressure is equal to (delta)H= q(for constant pressure). Which would make (delta)H's physical interpretation the heat added or lost to the system under constant pressure.

    The heat exchanged in closed system under constant volume is equal to (delta)U=q(for constant volume) meaning the heat added or lost to the system under constant volume is equal to the change in internal energy

    Now to the actual question does enthalpy have a physical meaning in this particular scenario?

    Under a closed system in which pressure and volume are not constant

    the change in enthalpy would be (delta)H= (delta)U + delta(PV)

    which would make (delta)H= (delta)U + V(delta)(P) +P(delta)(V)

    However V(delta)P seems to not have a physical meaning either. But if you subtract that term from both sides you get

    (delta)H-V(delta)P=deltaU+P(delta)V which I believe would equal this


    so is V(delta)P just a correction term to obtain the heat from the change in enthalpy? or do both of them actually have a physical meaning?

    Thanks in advance for any help
  2. jcsd
  3. Oct 9, 2014 #2


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    Staff: Mentor

    This post (and the thread itself) may be helpful.
  4. Oct 9, 2014 #3
    Thanks the thread was helpful
  5. Oct 10, 2014 #4


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  6. Oct 29, 2014 #5
    I have always thought of enthalpy as being a "thermodynamic potential energy" function which only carries meaning at constant pressure. Much like potential energy functions in Classical Mechanics, which can only be defined for conservative forces (such as gravity). I pair up "Heat" with "Enthalpy", and "Work done" with "Gravitational/Elastic/Electrostatic Potential Energy". I don't know whether or not this analogy is 100% correct though.
  7. Oct 30, 2014 #6
    I have an entirely different perspective for you to consider. Think of enthalpy as a physical property of a material, independent of any particular process applied to that material. The enthalpy per unit mass of a material is a unique function of its temperature and pressure. For any given material (e.g., water, methane, ammonia, etc.), it is possible to prepare a table of enthalpy values (per unit mass) as a function of temperature and pressure. The numbers in this table will never change, no matter what process the material is subjected to. For the case of water, the name we give to such a table is the "steam tables."

    Where do the numbers in these tables come from? Some of the numbers can be obtained by performing experiments on the material at constant pressure. In such experiments, the change in enthalpy from the initial thermodynamic equilibrium state to the final thermodynamic equilibrium state can be determined by measuring the amount of heat added to the material during the transition.

    Physical properties like enthalpy that are functions of temperature and pressure at thermodynamic equilibrium are called functions of state. Other properties that fall into this category are internal energy U, entropy S, Helmholtz free energy A, and Gibbs free energy G. The internal energy U can be interpreted physically as that total amount of kinetic energy of the molecules in the material plus the total potential energy of interaction between the molecules of the material.

    The enthalpy H, Helmholtz free energy A, and Gibbs free energy G are very convenient functions to work with when considering certain kinds of practical processes.

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