Chapter 17: Reaction Kinetics
What is kinetics?
In Chemistry, it is the study of the rates of reactions. Rates are the speed at which a reaction occur, which are obviously very different for various reactions. Rusting of iron occurs over a longer period than time than that of the combustion of magnesium.
Studying kinetics requires examining the reaction mechanism, the step-by-step sequence of reactions by which the overall chemical change occurs. For example, the formation of HI has two possible reaction mechanisms. The first possible way to form HI is by the following reaction mechanism:
Note that in step one, 2I is produced that is then used in step two. These species that are produced in one step and used up in another are called intermediates.
Another possible reaction mechanism is:
The intermediates in this process is H2I and I. In reactions in which all of the reactants and products are one state of matter (gases in this gas), they are called homogenous reactions.
Molecules must collide with the right amount of energy AND orientation in order for the reaction to occur. Even if two species collide at the correct orientation, the reaction may not occur because there is not enough energy available to break the reactants bonds. If there is enough energy but not the correct orientation, an effective collision would not occur.
The energy required to transform the reactants into an activated complex is called Activation Energy. Basically, it is the energy required before the products can even be produced. Even overall exothermic reactions require activation energy in order to occur due to the energy to overcome negative forces due to electron clouds and the breaking of bonds (which in itself is endothermic) before new bonds are formed (exothermic due to heat of formation).
This represents a reaction pathway that in the forward reaction is an exothermic reaction. The activated complex is the highest point on the curve with the activation energy being the energy to go from the reactants to the activated complex. The change in energy is the energy from the reactants to the products and since this value is negative, it represents an exothermic reaction. If the reaction is reversed (starting at the right-most part of the graph heading to the left), the activation energy is much greater as an endothermic reaction occurs.
Notice that this reaction mechanism still represents an exothermic reaction since the products have a lower energy level than the reactants; however, there are THREE activated complexes with three activation energies within the mechanism.
So what is the activated complex? Well, it is a transitional structure (somewhere in between the reactants and products) that results from effective collisions as old bonds break and new bonds form. On a graph, like the one above, each peak represents a different activated complex so there could be as few as one activated complex, but also as many as necessary depending on the number of steps in the reaction mechanism.
Influences on Reaction Rates
1. Nature of Reactants- Specific species are more active than others.
2. Surface area- Area of contact increases the rate of reaction, especially for heterogeneous reactions (reactants in different phases).
3. Temperature- Increases temperature increases the average kinetic energy of the particles resulting in more effective collisions. This increases the number of activated complexes through increased collision energy and collision frequency.
4. Concentration- Increasing the concentration of a reactant will increase the overall rate of reaction because it increases the number of collisions that occur.
5. Presence of a catalyst- Adding a catalyst will increase the rate of reaction by decreasing the activation energy required for the reaction. A catalyst is a substance that changes the rate of the reaction without itself being consumed. This means that a catalyst is not actually a part of the reaction! They do not appear in the final products or reactants of a product, but are drawn above the yields arrow as the catalyst. There are two kinds of catalysts:
Homogeneous- The same phase as the reactants and products
Heterogeneous- A different phase than that of the reactants.
Homework: Page 558 #'s 1-7