An Introduction To Reaction Kinetics (1959)

[Music] The control of chemical reactions is a subject of both scientific interest and great practical importance. Some of the factors involving the control of chemical reactions are illustrated in a relatively simple reaction that of H2 plus CL2 giving 2 HCl. This reaction can occur at varying rates. If hydrogen and chlorine are mixed together at room temperature in the dark, there is no apparent reaction. If the mixture is heated, the reaction takes place more readily. And if it is ignited by a spark, the mixture explodes. Let's now observe the steps in this reaction. Beginning with the molecules of the cold chlorine gas moving in slow motion, we see that the molecules rebound upon collision being repelled as their electron clouds begin to merge. As temperature is raised, however, their speeds increase and the more energetic collisions may break molecules into free atoms. Also, since chlorine is a colored gas, it can absorb visible light of relatively low energy and a single low energy photon thus absorbed may break a chlorine molecule into two free atoms. Now, if hydrogen is present, collisions may occur between these free chlorine atoms and hydrogen molecules to form structures called activated complexes, which have an extremely brief span of life. This activated complex can separate into a free atom of hydrogen and a molecule of hydrogen chloride. The hydrogen atoms thus produced can then strike other chlorine molecules to produce hydrogen chloride and free chlorine atoms. At increased camera speed, we observe that the reaction proceeds in this chain fashion producing more and more hydrogen chloride molecules and leaving fewer and fewer hydrogen and chlorine molecules. In addition to these chain propagating reactions, there is also a chain breaking reaction. This occurs when two free chlorine atoms strike a third body which absorbs their excess energy and allows them to combine to form a chlorine molecule. It is the energy of collision that differentiates those which initiate reactions from those which result merely in rebound. By means of this graph, we illustrate the energy involved. The vertical axis represents the potential energy. The horizontal axis called the reaction coordinate indicates the relative positions of the molecules or atoms during a reaction. This point indicates the relative energy positions of the reactants. Now using still slower motion, let us look at a reaction between a hydrogen atom and a chlorine molecule. As the reactants approach each other, the point moves to the right and we see that their relative velocity begins to decrease as the kinetic energy is changed into potential energy. Stopping our action, we observe that the reactants had sufficient energy for the system to reach a kind of changeover point, the activated complex. Here the atom and molecule are at the point of highest potential energy. In most respects, this activated complex is like any ordinary molecule except that it is in the process of coming apart. It has definite bond lengths, angles, and vibrational frequencies. Once the system has passed this point, no additional energy is needed and the reaction is quite certain to be completed. Let us observe this reaction again without interruption. Now on the other hand, if the reactants do not have sufficient energy, they collide but do not react and the system reverts to its former state. The potential energy barrier for each type of reaction is different. For the reaction between a hydrogen atom and a chlorine molecule, we have seen that this is the general shape of the barrier. While for the reaction between a hydrogen molecule and a chlorine atom, this is the general shape. In nature, relatively few reactions take place between atoms and molecules. Most collisions take place only between molecules. The reaction between hydrogen and iodine is a simple example of such a reaction. Most of these molecules do not have enough energy to react. But if they do, they first combine like this to form the activated complex and then split into two hydrogen iodide molecules. Here is the potential energy curve for the reaction between hydrogen and iodine. Recalling the potential energy curve for the reaction between the hydrogen atom and the chlorine molecule previously seen, we note that the hydrogen iodine potential curve is much higher. Again, in still slower motion, we see that the reactants must have console energy in order to surmount the potential energy barrier and react. Consequently, only a very few collisions such as this one result in a reaction. The height of the potential energy barrier depends upon the positions in which the molecules collide. In favorable collisions such as this, the reaction proceeds through a symmetrical activated complex in which this hydrogen iodine distance is equal to this. And this hydrogen iodine distance is equal to this. The height of the potential energy barrier is then the lowest possible for this reaction. For less favorable collisions, a higher barrier is encountered. Here the molecules collide at angles such as this. It is not possible to form a symmetrical activated complex and the molecules are less likely to have sufficient energy to react. A reaction and its reverse may take place at the same time. While hydrogen and iodine molecules are reacting in a forward direction to yield hydrogen iodide, hydrogen iodide molecules react in the reverse direction to yield hydrogen and iodine. The motion of the point from left to right on the potential energy surface represents the forward reaction. its motion from right to left indicates the reverse reaction. In general, the height of the potential barrier in the forward direction differs from that in the reverse direction. For this reaction between hydrogen and iodine molecules, the height is smaller for the forward reaction. While in other reactions such as that between a hydrogen molecule and a chlorine atom, the height of the barrier is smaller for the reverse reaction. Now in the reaction between a hydrogen atom and a chlorine molecule, the height of the barrier is much greater in the reverse reaction. Although the reverse reaction does occur, it requires much more energy and a considerably greater proportion of the collisions will be ineffective like this one as compared with those of the forward reaction. If we begin with a mixture of hydrogen and iodine gas, initially only the forward reaction producing hydrogen iodide occurs because as yet there is not a sufficient amount of hydrogen iodide present for the reverse reaction to be apparent. However, as more and more hydrogen iodide molecules are formed, the reverse reaction occurs more frequently until finally the rates of the forward and reverse reactions become equal and there is no further net change in the amounts of the reactants and products. At this point, equilibrium has been established. Now observe that the heights of the potential energy barrier for both forward and reverse reactions are nearly equal. Hence the energy required for going in both directions is almost the same and the amounts of the three reactants at equilibrium are nearly the same in the hydrogen and chlorine reaction. Equilibrium is also reached but the reaction goes almost completely to hydrogen chloride and only minute amounts of hydrogen and chlorine remain. The frequency of collisions between pairs of molecules increases with the product of their concentrations. Therefore, reaction rates are strongly influenced by the partial pressures of the molecular species involved. The effect of temperature can be appreciated by remembering that collisions are not effective unless they pass that halfway state called the activated complex. Whether atoms and molecules have sufficient energy to react depends on whether they are moving with sufficient relative velocity. It is possible to determine the energy needed for a reaction to take place by studying the effect of temperature on the rates of reactions. When appropriate substances are present and enter into the activated complex, the activation energy required is lowered and the reaction proceeds more rapidly. The finding of such enzymes and catalysts is a major portion of the study of reaction kinetics. The control of concentration, pressure, temperature, and catalysts is fundamental in applying the principles of reaction kinetics to the practical problems of industry. [Music]