Why is there an activation energy for a chemical reaction to take place? Why don’t all collisions result in the formation of new products? To answer these questions consider the analogy illustrated in Figure 2. Imagine you are trying to roll a bowling ball up a very steep hill. On most tries the bowling ball slows down and stops before it gets to the top of the hill. The kinetic energy of the bowling ball is converted to potential energy as the ball slows down. Then it rolls back down on the same side of the hill. The hill acts as a barrier. Only occasionally does the bowler give the ball enough kinetic energy so that it gets to the top of the hill and rolls down the other side. Once one the downhill slope, the potential energy of the bowling ball gets converted back to kinetic energy, causing the ball to pick up speed.
Picture a similar situation for molecules in a chemical reaction. During molecular collisions, atoms take up new bonding arrangements that have more potential energy than either the reactants or products. These atomic arrangements (i.e. the activated complex) have high potential energy like the bowling ball at the top of the hill. There is a minimum potential energy that must be achieved by colliding reactants before they can convert to some other form. This minimum potential energy is the activation energy for a given reaction.
The relationship between the activation energy and the energy absorbed or given off in a reaction (i.e. and endothermic or exothermic reaction respectively) can be shown graphically on a potential energy diagram (Figure 3).