Understanding why enthalpy is a state function

[AnkitaNair1E]

In lecture today, we defined enthalpy as the amount of heat absorbed/released in a system where pressure of constant. Given this definition, I'm still having trouble understanding conceptually why enthalpy is a state function. Could someone clarify this for me?

[Catherine_L_1C]

Hi,
Enthalpy is considered a state function because its current value will only depend upon the final and initial values of heat in a reaction, but not the path or process that occurred for it to reach that value. In other words, theoretically a reaction could have the same enthalpy value whether it started as a solid to a liquid or if it turned a liquid into a gas. This is in contrast to a path function, with processes like heat and work falling in this category.
Hope this helped!

[Tycho_Meimban_2B]

Enthalpy depends only on two thermodynamic properties of the state the substance is at the moment, not on the path that it took to get there. When calculating enthalpy, we use only two values (the temperature at the beginning, and the temperature at the end) to calculate our value, unlike heat, which depends on the pressure during heating/cooling and not just the initial and final temperatures and pressures.

If the calculation for heat did not include using pressure in its calculation, then we would be able to classify heat as a state function since we would only consider its two states (beginning and end) for us to calculate.

Hope this helps!

[Alex Uy 2D]

Could someone explain why enthalpy (which I understand to be heat at a constant pressure) is a state function while just heat is not a state function? Why does having a constant pressure suddenly make heat become enthalpy and therefore a state function?

[Chem_Mod]

If pressure is constant, then the PV work is equal to zero. If we begin with the equation q = ∆U - w, we will eventually get to the equation ∆H = q_p (since there is no PV work being done).