7.7: Activation Energy
Activation energy is the minimum amount of energy necessary for a chemical reaction to move forward. The higher the activation energy, the slower the rate of the reaction. However, adding heat to the reaction will increase the rate, since it causes molecules to move faster and increase the likelihood that molecules will collide. The collision and breaking of bonds represents the uphill phase of a reaction and generates the transition state. The transition state is an unstable high-energy state of the reactants. The formation of new chemical bonds and end products, and the release of free energy is the the downhill phase of the reaction. Catalysts increase the rate of a reaction by lowering the activation energy. For example, in biology, enzymes that metabolize sugar and fats increase the rate at which their breakdown happens and at the same time, prevent the overproduction of free energy that would otherwise denature proteins in the cell.
A catalyst is a substance that increases the rate of a reaction by lowering the activation energy, and in the process, regenerates itself. A catalyst provides an alternative pathway or mechanism for the reaction to take place and it accelerates both the forward and reverse reactions. In biology, enzymes are examples of catalysts because they lower the activation energy required for reactions in cellular metabolism.
For example, humans metabolize sugar and fat for energy. Enzymes are vital to humans for breaking down these molecules, because if thermal energy alone were to be used, the free energy released in the form of heat would cause proteins in the cell to denature. Furthermore, thermal energy would non-specifically catalyze all reactions. However, enzymes only bind to specific chemical reactants, called substrates, and lower their activation energy to catalyze selective cellular reactions.