In the last module, thermochemistry, we touched on the idea that chemical changes can occur at different rates. They can be fast, like an explosion, or slow, like the rusting of iron. We also mentioned that some reactions, although thermodynamically favored, are not actually observed to occur. Chemical kinetics is the study of reaction rates and the factors that determine those rates. Let’s begin our study by looking at how reaction rates are described.
The diagram on the right shows the progress of the reaction that takes place in a container of constant volume as chemical A, in red, changes into chemical B, in blue.
A → B
Looking at the left side of the diagram, we can see that we start with all A and no B molecules. Over time (looking from left to right on the diagram), the concentration of A decreases and the concentration of B increases. The reaction rate is described as the speed with which reactants disappear or products appear. The steeper the lines on the concentration vs. time graph, the faster the reaction rate is.
A → B
Before we describe reaction rates, we will use a theory to help us understand how reactions take place and how they are affected by different factors. We call this the collision theory. It states that the rate of a reaction depends upon the number of effective collisions per second among the reactant molecules. An effective collision is defined as one that actually produces a product. Only a very small percentage of collisions can really lead to a change. Why? Two conditions must be met in order for a collision to be effective.
In most reactions, molecules have to be oriented in a certain way so that the atoms that need to collide can do so. Below is a diagram of the reaction, Cl2 + H2 → 2HCl. In this image, the Cl2 molecule and H2 molecule are oriented end to end as they approach each other. This is not the proper orientation to form 2 HCl. So, the molecules collide and bounce off each other unchanged.
In the bottom image, the Cl2 molecule and H2 molecule are oriented side to side so that when they meet, they are oriented properly to form 2 HCl .
Molecular Kinetic Energy
Even if the molecules are oriented correctly, they must have enough kinetic energy to overcome the repulsion of their outer electrons and actually produce a reaction. If they do not have enough energy, the molecules will simply bounce apart. The minimum kinetic energy required for the reaction to take place is called activation energy, Ea.
Potential Energy Diagrams
To visualize the changes in energy that take place as molecules approach each other and then react, we use a potential energy diagram. The y-axis represents the changes in potential energy as the particles collide. The x-axis is called the reaction coordinate or reaction pathway. You can think of this as the passing of time. Look at the series of pictures to see how the potential energy diagram is drawn and labeled.
Did you notice the activation energy, Ea, in the diagram? You can think of this like the energy hill that the particles must be able to climb in order to successfully produce products. The activated complex is the substance that is formed before the products are formed. It is a very unstable substance, formed only for the very brief moment where the bonds of the reactants are partially broken and the bonds of the product are partially formed. The potential energy of the activated complex corresponds to the highest point on the potential-energy diagram. The activation energy is the difference in the PE of the activated complex and the PE of the reactants.
The difference between the potential energy of the reactants and the potential energy of the products is the heat of reaction, ΔH. You learned about ∆H in the thermodynamics module. Remember that if energy is released during a reaction, we call this: exothermic. If energy is required for a reaction, we call this endothermic. You can very easily tell if a reaction is endothermic or exothermic by looking at the potential energy diagram. If the potential energy of the products is less than that of the reactants, the reaction is exothermic. (You can remember the mnemonic device exo low.Exothermic PE diagrams end lower than they start.) If the potential energy of the products is more than that of the reactants, the reaction is endothermic.
For this reaction, label Ea, ΔHr, and activated complex.
Is the reaction endothermic or exothermic?