Internal energy $U$ is clearly an important concept; the first law of thermodynamics states that for an isolated system internal energy is constant $(\Delta U=0)$ and that for a closed system the change in internal energy is the heat absorbed by the system $Q$ and work done on the system $W$ $(\Delta U=Q+W)$.
Enthalpy $H$ is the sum of the internal energy and the product of the pressure and volume of the system $(H=U+PV)$. I was taught that enthalpy is a preferred quantity to internal energy for constant pressure systems where $\Delta H=Q$, as opposed to constant volume systems where $\Delta U=Q$. But why would anyone care about what quantity is equal to $Q$ under certain conditions instead of simply reporting $Q$? Enthalpy seems redundant in this context.
Is enthalpy a convenient concept in other contexts, such as systems with varying pressure? Is it described by any fundamental laws as internal energy is by the first law?