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I know it is pretty hard to measure the change in volume in open systems. Accordingly, calculating the change in internal energy using the first law of thermodynamics is pretty hard also. Is this the only reason for introducing enthalpy ($H$) as a new thermodynamic potential to make use of the constant atmospheric pressure, and to calculate the change in enthalpy ($\Delta{H}$) as it is equal to the in/out quantity of heat ($Q$) in this case? If no, does the absolute value of enthalpy have any physical significance?

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  • $\begingroup$ You are right. Of course, the enthalpy has been introduced to take care of the change of volume. $\Delta$ H is by definition equal to the quantity of heat Q, without taking the work into account. It is sort of "apparent internal energy". It is a little bit similar to the difference between the weight on an object in a vacuum, and the weight measured in air, which a little bit smaller than in a vacuum, due to Archimede's correction. $\endgroup$ – Maurice Apr 19 at 9:33
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The absolute value of enthalpy has direct relation to the mass of the system:

$$H = U + pV = mc^2 + pV$$

with the second term mostly negligible. If the system accepts thermal energy Q, its mass change is $\Delta m = \frac {Q}{c^2}$

For most of enthalpy applications, the absolute $H$ value is irrelevant.

The absolute value interpretation is the total energy stored in the system at the constant pressure, consisting of

  • the equivalent energy to the mass of contained matter
  • work done by reaching the system volume against the external pressure.

This energy is formally extractable from the system by annihilation and shrinking to zero volume.

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