I've been trying to get my head around the differences between enthalpy change and internal energy change for a system. Here's my understanding of it so far. I'd appreciate any feedback on my understanding so far on the concepts - especially if there are misconceptions (and there probably are)
- Internal energy is the 'intrinsic' energy of some system, can be measured with reference to something within the system, e.g. chemical bonds, vibrational energy states, etc.
- Enthalpy change is the heat (i.e. energy transferred due to temperature difference) change at constant pressure due to some chemical process.
$\Delta H = \Delta U + P \Delta V$
$\Delta H$ is characteristic of a given chemical process - i.e. this value is the same regardless of what conditions it is run under (e.g. varying pressures, volumes, starting temperatures). As long as it's the same reaction, $\Delta H$ will be the same. It's the overall energy input required to make a process work, and is split between influencing $U$ and performing $PV$ work.
$\Delta U$ is a function of the specific conditions (i.e. volume, pressure) in which a specific instance of a chemical process is run.
So it's theoretically possible to set up a chemical process (say vaporisation) to run in such a way so that $\Delta U$ is zero, and all the enthalpy change is accounted for in the $P\Delta V$ term?
And what about cases where P is not constant? What happens then? Can we calculate the value of $\Delta H$? And does $\Delta H$ even have a meaning in this case?