# Why are processes where $\Delta G \le 0$ considered spontaneous?

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.
• 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. – sangstar May 8 at 14:44
• 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? – Chet Miller May 8 at 15:01