I understand how, for any process where starting pressure and temperature start and end points are the same, that

$$\Delta G \le 0$$

But I don't see how, from looking at this, that when $\Delta G$ is negative, this implies processes are spontaneous. Why is this?


It's just a rule of thumb, and is not exact. More precisely, $\Delta G < 0$ means that the equilibrium constant is large, thus favoring products over reactants. $\Delta G > 0$ means that the equilibrium constant is small, thus favoring reactants over products.

  • $\begingroup$ What is the equilibrium constant? Also, I've never seen an example where $\Delta G$ is used for anything. Is there one you could think of? It all sounds incredibly abstract to me looking at this for the first time. $\endgroup$ – sangstar May 8 at 14:44
  • $\begingroup$ Are you saying that you are not familiar with the equation $$RT\ln{K}=-\Delta G^0$$where K is the equilibrium constant, expressed in terms of the partial pressures of products and reactants at equilibrium, and $\Delta G^0$ is the standard Gibbs free energy change for the reaction? $\endgroup$ – Chet Miller May 8 at 15:01
  • $\begingroup$ I suppose I am saying that, as I'm still in the process of learning thermodynamic potentials and I am not familiar with this. $\endgroup$ – sangstar May 8 at 15:02
  • $\begingroup$ If your course includes chemical thermodynamics, you will learn about this soon. In my opinion, this material is a little complicated (i.e., hard to get used to), so take a deep breath. $\endgroup$ – Chet Miller May 8 at 15:43

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