Chemists often write chemical equations for reactions as a single step that shows only the net result of a reaction. However, most chemical reactions occur over a series of elementary reactions. The complete sequence of these elementary steps is called a reaction mechanism. The reaction mechanism is the step-by-step process by which reactants actually become products. It is the "how" of the reaction, whereas the overall balanced equation shows only the "what" of the reaction. In kinetics, the rate of a reaction with several steps is determined by the slowest step, which is known as the rate-determining, or rate-limiting, step. Take the following example of a gas phase reaction:
If this reaction occurred in a single step, its rate law would be:
However, experiments show that the rate equation is:
You will learn more about experiments to determine rate constants shortly. That the rate law determined experimentally does not match the rate law from the equilibrium reaction suggests that the reaction occurs over multiple steps. Further, the experimental rate law is second-order, suggesting that the reaction rate is determined by a step in which two NO2 molecules react, with the CO molecule entering at another, faster step. A possible mechanism that explains the rate equation is:
The overall reaction rate depends almost entirely on the rate of the slowest step. Here, the first reaction produces a new molecule, NO3, which is neither a reactant nor a product. The second step then consumes that molecule, and NO3 therefore does not appear in the overall reaction. Because of this, NO3 is called a "reaction intermediate." Intermediates play important roles in the rates of many reactions. Since the first step is the slowest, and the entire reaction must wait for it, it is the rate-determining step. The other steps are fast enough that their rates are irrelevant, as they are always waiting for the slower step to complete. If the first step in a mechanism is rate-determining, it is easy to find the rate law for the overall expression from the mechanism. If the second or a later step is rate-determining, determining the rate law is slightly more complicated. We will explore how to write that rate law later.
Why do reactions have different rates? In a reaction mechanism, the reaction that requires the most energy is the rate-determining step. The minimum energy needed for a reaction to occur is called its activation energy. The activation energy is controlled by a few factors: how often the reactants collide, with how much energy they collide, the angle at which they collide, and how easily the electrons bond. The particles must be moving fast enough for their collision to satisfy the activation energy. Without the necessary energy, the particles will bounce off of each other with no reaction. The particles must also collide in the proper orientation so that the bonding orbitals overlap, enabling a bond to form. For a chemical reaction to proceed at a reasonable rate, there should be an appreciable number of molecules with energy equal to or greater than the activation energy. The reactions can be accelerated by increasing the concentration of the reactants, changing the temperature of the system to increase the energy of the collisions, or adding an catalyst to help orient the bonding orbitals or make the reactants more reactive (Figure 1). Understanding the rate-determining step and the activation energy is very important to the optimization and understanding of many chemical processes, such as catalysis and combustion.