It's somewhat "incorrect" use of words in physics, because heat is not a state variable, there is no unique number we can assign to any system that would tell us how much heat it contains in some objective universal sense. This is because heat is a mode of energy transfer, and the same energy can go in via heat, or work. After the energy is in the system, there is no way to find out, from state variables of the system, whether it came in through heat, or work.
"Heat content" refers to and applies to a restricted set of situations, where we put in some energy only via heat, not via work, during this transfer the system changes its state (typically it expands and does work on the atmosphere), while the external pressure $p$ remains constant (usually, due to atmosphere). Later we can reverse all of this and extract that same amount of heat and put it elsewhere. During heat absorption/heat release, the system is allowed to do work or accept work from atmosphere.
Enthalpy comes up because in this kind of process with constant external pressure, its change equals the heat accepted:
$$
\Delta H = Q.\tag{*}
$$
So if 2.2MJ of heat is accepted by the system, then its enthalpy has increased by 2.2MJ. In this context, telling the value of enthalpy (and the minimum enthalpy the system had before we put heat in it) tells us how much heat we can recover back from the system.
Why does enthalpy behave this way? Internal energy $U$ doesn't; this is because part of the energy income is going out via work on the outside atmosphere. But enthalpy is defined as
$$
H = U + pV
$$
to get a quantity that does not decrease due to volume work. We can see this easily: we have
$$
\Delta H =\Delta U + \Delta(p V) = \Delta U + p \Delta V
$$
and since the 1st law of thermodynamics for volume work implies
$$
\Delta U = Q - p\Delta V
$$
combining these two equations, we obtain (*).