We have established that in order for a reaction to occur at all, ΔG < 0, however, we still have not concluded anything about the rate of reaction.  This is the realm of kinetics.

For example, petrol at room temperature is stable, although volatile.  Cars do not spontaneously explode; the petrol sits in the tank essentially for any length of time without burning.  This is despite the fact that the combustion of petrol is highly exothermic (there is a large enthalpy change (ΔH), and accompanied by a favourable entropy increase (ΔS).

Petrol at room temperature is kinetically stable, because although the Gibb's energy change will be very large, and K will be enormous (i.e. essentially no product left), this reaction at room temperature takes a very long time.  At elevated temperatures (i.e. that of a match, or spark-plug spark), this reaction proceeds much faster (hence cars work!).

What we have stumbled across is that the rate of reaction is not determined by ΔG, but by something else; the energy of activation:

ΔG‡ is this energy of activation; in simple terms, it is the energy hump that the system must overcome before the reaction can proceed.  This basic scheme shown above is a very useful representation that we shall refer to in the future.

The peak of the graph shown above is known variously as; the transition state, or activated complex.  It is definitely not an intermediate, which we will discuss later.  The transition state is an unstable state which is passed through momentarily by a reacting species:

Transition State

The species above in brackets is the transition state of that particular reaction.  The more unstable this transition state is, the higher the energy of activation, and hence the slower the reaction is.