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.

The concept of an "activation-energy " barrier to reaction is illustrated with the mountain analogy on the previous page. The boulder cannot roll off the edge of the mountain without first surmounting the activation barrier crowned by a double dagger symbol (the conventional indication of an activated intermediate state).

A catalyst makes a chemical reaction go faster by providing an alternate path with a lower activation-energy barrier. This is symbolized by the winding path down the side of the mountain.