Reaction Kinetics: Reaction Mechanisms
By describing how atoms and molecules interact to generate products, mechanisms help us to understand how the world around us functions at a fundamental level. A mechanism is a series of elementary steps whose sum is the overall reaction. An elementary step is a reaction that is meant to represent a single collision or vibration that leads to a chemical change. For a mechanism to be considered valid, its sum must equal the overall balanced equation, its predicted rate law must agree with experimental data, and its predictions of intermediates must not be contrary to experimental observations. A mechanism may never be proven because we cannot ever see a chemical reaction--both the time scale of an elementary step and the size of atoms are too small. Furthermore, we must guess at the identity of many intermediates because they are usually so reactive that they can not be isolated. Instead, a chemist proposes reaction mechanisms and tests their validity against experimental data, ruling out mechanisms that are inconsistent with results. These experiments may be strategically designed to trap an intermediate product to prove its existence as a stepping-point in the total reaction.
To aid in our understanding of mechanisms, we will draw reaction coordinate diagrams that trace the free energy path of a reaction from reactants to products. The activation energy of a reaction represents the difference in energy between the reactants and the highest point on a reaction coordinate diagram. We will derive the Arrhenius Equation, which relates the rate constant for a reaction to its activation energy. Local minima on the reaction coordinate diagram are positions occupied by intermediates. By comparing the reaction coordinate diagram for a catalyzed and a uncatalyzed process, we can see that catalysts function by altering the route the reaction takes from reactants to products without the catalyst being altered.