The thermodynamic property enthalpy is defined by

$$\boxed{H=U+pV},$$

where $U$ is internal energy, $p$ is pressure and $V$ is volume.

Because $U$, $p$ and $V$ are properties of the state, enthalpy is also a state property. Its change

$$\Delta H=H_{\mathrm f}-H_{\mathrm i}$$

depends only on the initial and final states, not on the sequence of intermediate steps used to connect them.

Enthalpy is not “heat stored in a system.” Heat is energy transferred because of a temperature difference; enthalpy belongs to the state.

Why enthalpy is useful at constant pressure

For a closed system at constant external pressure with pressure-volume work as its only mechanical work mode, the first law gives

$$\Delta U=q_p-p\Delta V.$$

At constant pressure,

$$\Delta H=\Delta(U+pV)=\Delta U+p\Delta V,$$

so

$$\boxed{q_p=\Delta H}.$$

This relation is why enthalpy changes are especially convenient for reactions carried out in open vessels at approximately constant atmospheric pressure.

If $\Delta H<0$, the system releases energy as heat to the surroundings under these conditions; if $\Delta H>0$, it absorbs heat.

Molar and mass-specific forms

Enthalpy can be reported per amount of substance or per mass. The molar enthalpy is $H/n$. For engineering calculations, defining

$$h=\frac{H}{m},\qquad u=\frac{U}{m},\qquad v=\frac{V}{m}$$

gives

$$h=u+pv.$$

These normalized forms change the scale used to report enthalpy, not the underlying state-property concept.