Some reactions are not 100% complete, even after an infinite period of time - they remain as a mixture of reactants and products after all visible reaction has ceased. A competition exists between forward and reverse reactions, and equilibrium is a condition of balance between these opposing processes.
We now come to the question: Why do some reactions, which by their free energy values should be spontaneous and far from equilibrium, sit inert and unreactive for years, whereas other reactions go with explosive rapidity? The decomposition of NO to nitrogen and oxygen is thermodynamically spontaneous, so why do we have photochemical smog from oxides of nitrogen?
If all combustions with oxygen liberate free energy, and the atmosphere is full of oxygen, then why doesn't everything that is potentially flammable burn at once, including ourselves? The answer is that these decompositions and combustions, although thermodynamically spontaneous, occur at miniscule rates at room temperature. The rates of chemical reactions and the factors that affect them are the subjects of this chapter.
The central theme to be developed in this chapter is that the rate of a chemical reaction depends on its reaction mechanism . Two molecules coming together must collide and rearrange their atoms to make product molecules. The intermediate arrangements of atoms may have a high energy, and if so, the reaction will be slow because not all colliding molecules will have enough energy to rearrange properly.