The pV term may be understood by the following example of an isobaric process. Consider gas changing its volume (by, for example, a chemical reaction) in a cylinder, pushing a piston, maintaining constant pressure p. The force is calculated from the area A of the piston and definition of pressure p = F/A: the force is F = pA. By definition, work W done is W = Fx, where x is the distance traversed. Combining gives W = pAx, and the product Ax is the volume traversed by the piston: Ax = V. Thus, the work done by the gas is W = pV, where p is a constant pressure and V the expansion of volume. Including this pV term means that during constant pressure expansion, any internal energy forfeited as work on the environment does not affect the value of enthalpy. The enthalpy change can be defined ΔH = ΔU + W = ΔU + Δ(pV), where ΔU is the thermal energy lost to expansion, and W the energy gained due to work done on the piston.
Difference between enthalpy and internal energy
Chemists routinely use H as the energy of the system, but the pV term is not stored in the system, but rather in the surroundings, such as the atmosphere. When a system, for example, n moles of a gas of volume V at pressure P and temperature T, is created or brought to its present state from absolute zero, energy must be supplied equal to its internal energy U plus pV, where pV is the work done in pushing against the ambient (atmospheric) pressure. This additional energy is therefore stored in the surroundings and can be recovered when the system collapses back to its initial state. In basic chemistry scientists are typically interested in experiments conducted at atmospheric pressure, and for reaction energy calculations they care about the total energy in such conditions, and therefore typically need to use H. In basic physics and thermodynamics it may be more interesting to study the internal properties of the system and therefore the internal energy is used.
