Internal Energy vs Enthalpy I struggle to see the difference between internal energy and enthalpy, or at least the need for differentiation between the two. I understand that enthalpy, supposedly unlike internal energy, accounts for the expansion work required for the system to exist, but it seems that that would only be necessarily stated if one had failed to account for the negative contribution to work in the first place. 

Could enthalpy, as $$H = U + pV$$ and thus, $$H = q + w + pV$$ not be written as equally validly as, $$H = q + w_1 + w_2$$ i.e. $$H = U \ ?$$

It would seem that this definition of enthalpy implies that internal energy accounts only for work done to the system, disregarding work done by it.
Trying to figure out a good intuitive physical interpretation of enthalpy has caused thermodynamics student over the last couple of centuries to waste huge amount of their valuable time (in my judgment).  The really fundamental physical quantity is internal energy U, not enthalpy, and its physical interpretation is simple and straightforward.  I regard enthalpy as just a convenient function to work with in solving many types of physical problems.  If the function enthalpy had never been invented, we would still be able to solve these problems, but the results would not be as concise.