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I haven’t really found a clear answer.

Enthalpy is measured as the sum of internal energy (kinetic and potential energy) and pressure times volume. However, I don’t really see the value of measuring enthalpy, so why is it even mentioned? Aren’t we more concerned with changes in enthalpy, like enthalpy of vaporization, enthalpy of fusion?

In other words, what are the differences of enthalpy and delta enthalpy? Is the former just a formality, and we are really concerned about measuring the energy of a thermodynamic system undergoing change (i.e. chemical reactions)?

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    $\begingroup$ If you did not know how is defined particular function, how would you know, how to calculate it's differences, differentials or derivatives ? $\endgroup$ – Poutnik Apr 18 at 3:14
  • $\begingroup$ Enthalpy is defined, not measured, as H = U + p.V. It is not measurable quantity. Measured are quantities and their changes that affect the enthalpy change. $\endgroup$ – Poutnik Apr 18 at 8:06
  • $\begingroup$ Do a few actual problems, rather than spending you valuable time speculating. You will bet the idea of how it plays out. $\endgroup$ – Chet Miller Apr 18 at 11:56
  • $\begingroup$ @ChetMiller I have done a few problems, which is why I came here. I regularly see enthalpy problems being on changes in enthalpy, rather than enthalpy itself. It seems that enthalpy, like Poutnik said, cannot be measured directly; only changes in enthalpy can. So I’m wondering what exactly its value is - for measuring enthalpy in reactions, formations, etc.; or is it more? Because if it’s more, then I wanted to clarify. $\endgroup$ – Alex Apr 18 at 12:54
  • $\begingroup$ If a quantity is not defined, it's changes are not defined either, even if only these changes really matter a/o can be measured, directly or indirectly. $\endgroup$ – Poutnik Apr 18 at 16:00
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Enthalpy vs ∆Enthalpy
We know that enthalpy represents the energy of a system. So, now think about it? Measuring the energy of a system? Consider all the changes and influences a system is under. This makes it really difficult (we could say impossible) to calculate the enthalpy of a system. But we do know one thing, that the enthalpy depends on two factors. A system will experience change in enthalpy if it's internal energy is changed and/or if work is done on/by it. So what we do is that we assume an initial state where we start our observation and a final state where our observation concludes. And the changes in state cause particular changes in enthalpy and this change in enthalpy of the system is what matters to us anyway.

Why did we define it then if we don't use it? We just know that a particular quantity depends on two particular factors. It's just that we understand the fact that it is meaningless to calculate the enthalpy of a system - because we do not know where to start from? Where will you start from? You need a point where the enthalpy was zero initially and then gained enthalpy right? But imagine - doesn't that mean there won't be any atomic activity at all if we're starting from zero enthalpy? And if there wasn't atomic activity then that chemical specie might not even be able to exist? It's complicated.

But then we realise: all that matters to us is just change in enthalpy during a thermodynamic process. And that is what thermodynamics is essentially all about - energy in motion. Thermodynamics is but a study of states (this was mentioned in NCERT Chemistry). So we study the change in Enthalpy from one state to another.

We define the enthalpy and what it depends on so that we can observe the factors that it depends on.

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  • $\begingroup$ Very excellent response! That makes a lot of sense - that because systems are constantly undergoing infinitesimal change, we cannot possibly overconcern ourselves with the enthalpy but rather the change in enthalpy. One remaining question: When a system has undergone a change in enthalpy, we say that it required or exerted energy (endothermic vs. exothermic). That energy becomes part of the system’s new enthalpy. So then, when we talk about enthalpy, why isn’t it made clear from the start that we are concerned with changes in enthalpy and not with the definition itself? $\endgroup$ – Alex Apr 18 at 13:01
  • $\begingroup$ With what I understand from your counter-question, I believe it is just a matter of exposure to resources. Sometimes we read and interpret things in different ways or sometimes overthink about it when it's basically just simple fundamental things. This especially happens in chemistry and we see learners asking about why a particular thing in chemistry is the way it is although it is a good trait to do so. My first teacher of thermodynamics did emphasise that ∆H is what we focus on and not enthalpy and maybe that's the reason I was able to address your confusion. $\endgroup$ – Desai Apr 19 at 11:46
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Enthalpy is a fundamental physical property of the material(s) comprising a system. It can often be presented for a given material in the literature in tables (e.g., the steam tables for water), relative to some specific datum state (at which it is taken to be zero). This is analogous to potential energy which is expressed relative to some specified elevation.

In the case of enthalpy, we can't only tabulate changes between states since this would require an infinite number of tables. So, in order to make and apply enthalpy tables, we need to reference everything to a specified datum state (for that material). The values in the table can be regarded as "absolute enthalpy." We can then determine changes between states by subtracting values for state 1 from state 2.

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