# Why does enthalpy change and not internal energy help in determination of spontaneity?

Spontaneity of a process is measured by Gibbs Free Energy $$\Delta G= \Delta H-T\Delta S$$ $$\Delta H=\Delta U+\Delta (PV)$$

Is there an intuitive explanation for why Gibb's energy depends upon enthalpy change rather than internal energy change.

It seems intuitive that since systems try to attain lowest energy states, negative value $$\Delta U$$ should increase stability and therefore spontaneity.

It also seems intuitive that since particles try to achieve equilibrium, maximum disorder/maximum microstates/maximum energy distributed and therefore positive value of entropy$$\Delta S$$ would increase spontaneity.

Is there a logical explanation why/where the extra $$\Delta (PV)$$ term helps determine the spontaneity of a process or stability of products?

I do mostly get the mathematics of the equation but not logic.

• See difference between Gibbs free energy ($G=H-TS$) and Helmholtz free energy ($F=U-TS$). Commented Mar 19 at 1:38
• And the reasoning for introducing $H = U + pV$. Commented Mar 19 at 5:13