Enthalpy and Heat: What is the Relationship Between Them and How Do They Differ?

[kathyt.25]

Concerning the first law of thermodynamics:

dE = q - w
dE: change in internal energy of a system
q: heat added/removed from system
w: work on/by the system

What is the difference between q (heat added/removed from the system) and H (enthalpy, or heat flow)? Are they analogous to each other?

 

[JaWiB]

First off, enthalpy is a state quantity and heat is not. For instance, you can talk about how much enthalpy a system has, but you can't talk about how much heat it has.

You could, however, compare a change in enthalpy to the heat added to a system. Enthalpy is defined as:
[tex]H = U + PV[/tex]
Where P is the pressure of the environment and V is the volume of the system. If you have a change in enthalpy at constant pressure, then:
[tex]{\Delta}H = {\Delta}U + P{\Delta}V= Q-W+P{\Delta}V[/tex] (using the convention you chose where W is the work done by the system.

As long as the only work you do on the gas is compression/expansion, [tex]W=P{\Delta}V[/tex] so the change in enthalpy is equal to the heat added to the system. So in a sense enthalpy is a way of letting you ignore compression-expansion work during a given process.