Why we needed to define enthelpy?

  • Thread starter Frigus
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In summary, the introduction of the enthalpy term was necessary because it allows for a more accurate measurement of heat under constant pressure. This is because eliminating volumetric work in the equation dU = dq - p·dV is difficult in practice, but the definition of enthalpy as H = U + p·V allows for the elimination of non-volumetric work and results in a simpler equation, dH = dq + V·dp, which is easily measured under constant pressure.
  • #1
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If we know that at constant pressure heat absorbed or gained is path independent then what was the need for introducing this new enthalpy term.
If its answer is that it is not only for a special case in which pressure is constant then how can we even use it because we cannot measure the internal energy.
Thanks
 
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  • #2
I don't understand your question. Please provide a specific problem (with actual temperatures, pressures, amounts of material, etc) that better illustrates what your issue is.
 
  • #3
Hemant said:
If we know that at constant pressure heat absorbed or gained is path independent then what was the need for introducing this new enthalpy term.

The change of internal energy is dU = dq + dw. It is easy to eliminate non-volumetric work. But than you still have dU = dq - p·dV. Eliminating the volumetric work would require keeping the volume constant which is difficult in practice. That means that an accurate calorimetric measurement would always need to be accompanied by a measurement of the change in volume.

This problem is solved by the definition of enthalpy H = U + p·V. That results in dH = dU + p·dV + V·dp, without non-volumentric work in dH = dq + V·dp and under constant pressure in dH = dq. If you manage to keep the pressure constant (which is quite easy in practice) the calorimetric measurement directly gives you the change of enthalpy.
 

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