A spontaneous reaction is one that requires no external and continuous input of energy. It occurs when G<0. Very straightforward. When G>0, the reaction is not spontaneous ok. As a reaction reaches equilibrium, G trends towards while S is maximized. Ok. Very clear understandable.
My confusion begins with understanding how entropy can drive a reaction.
So lets assume a endothermic reaction: A + B –> CD + EF. Great! This reaction proceeds spontaneously because the products have a higher entropy than the reactants and the magnitude of entropy * T is higher than enthalpy.
Even though I am able to categorize this reaction as spontaneous. Why does this reaction occur spontaneously??? Obviously, the molecules involved in this reaction (specifically A and B) have no clue that the products have more microstates. Why do the reactants react in the first place than? What is the exact mechanism involved.
Why do exothermic reactions typically proceed spontaneously (with an initial energy inputted). I don’t understand this at all. I don’t want to have to just rely on the math and lack an intuitive understanding.
Edit: I wanted to add that I understand entropy to be a function of probability. Example: if I flip a coin 1000 times, it is more likely that I get 500 heads and tails. But just because the products have more arrangeable states does not explain why the reaction proceeds in the first place. For example, using the coin analogy, even though there are a lot of microstates with 500 heads and tails, why does the coin flip in the first place (if that makes sense).
Edit 2: My only working idea right now is that because products have higher entropy, the chances of the products interacting to form reactants is less likely. But this does not explain something like diamond spontaneously turning into graphite so I am reluctant to believe it.