Chemical Kinetics Simulation

What's happening?

If you saw nothing on the previous page, then your browser does not support Java applets.

Otherwise, you found yourself at the controls of a simple simulation of a binary chemical reaction. There are four types of molecules in this simulation, red, yellow, green and blue. A pair of red and yellow molecules if they collide may react to form a pair of green and blue molecules, and vice versa. This chemical reaction would be written like so:

R + Y G + B
The central double arrow in this reaction indicates that both the forward (R + Y --> B + G) and reverse (B + G --> R + Y) reactions are in principle possible. This is true of all chemical reactions. However, as is also true of all chemical reactions, the probability of reaction during each collision, abbreviated "k", is not necessarily the same for the forward and reverse reactions.

Take command!

You are given controls at the top of the simulation, using which you can set the initial concentrations (number per unit box) of molecules of R, Y, G and B and the reaction rate constants (probabilities of reaction per collision) kf (R+Y --> G+B) and kr (G+B --> R+Y). To use the controls, just click your mouse over the appropriate box, type in the new numbers you want, and press the "Reset" button. The simulation will immediately begin over again with your new values. Experiment!

Some notes on performance

  1. The simulation is running on your computer, so the complexity of the simulation is limited by your computing power. If you find the simulation running too slowly or jerkily, you can try quitting any other applications or windows you have open. You can also reduce the workload on your computer by reducing the total number of molecules in the simulation. The default setting is a total of 20 molecules, but to get the best appearance you'll want to increase that if you can.

  2. To stop the Applet from running you have to leave the Web page on which it sits, or turn off Java on your browser. If the applet doesn't run at all, you can try reloading the page by clicking the RELOAD button on your browser, sometimes while holding the SHIFT key down.

  3. You should see three things on the page: the controls at the top of the simulation, the simulation itself, and at the bottom a constantly updated list of the number of molecules in the box. If you don't see all three components, or they are weird or distorted, or you have other problems, please contact the author of this Applet so he can try to fix things.

Some note about chemistry

  1. If you set the reverse probability to zero, the reaction goes "to completion." What are the final concentrations (number of molecules per unit box) of reactants and products in terms of the initial concentrations? In this case the initial concentrations and the "stoichiometry" (the ratio of products to reactants set by the reaction) control the final concentrations.

  2. If, on the other hand, neither probability is zero, then the final concentrations will on average after a while ("at equilibrium") be some constant numbers different from zero for all species. The equilibrium ratio of products to reactants is a constant K, 
             [ C ] [ D ]
  K  =   -----------
         [ A ] [ B ]
    called, with great originality, "the equilibrium constant." Can you figure out what K should be in terms of the two reaction rate constants kf and kr? The answer is rather simple and does not depend on the absolute size of either kf or kr, or on the initial concentrations of the reactants and products.

  3. At equilibrium the average concentrations are not changing, indeed. But you will see that the instantaneous (moment by moment) concentrations certainly do. These are the "concentration fluctuations" at equilibrium, and they are very important because they are (1) properties of the equilibrium system, and so can be calculated by relatively simple theory from measurements of the system at equilibrium, but (2) these fluctuations tell you a lot about how the system behaves dynamically when it is not at equilibrium. Perhaps by thinking about how fluctuations away from equilibrium happen in this simulation and how the simulation approaches equilibrium you will see intuitively the truth of this fundamental and important theorem -- called the "Fluctuation-Dissipation Theorem", and a bit of a trick to prove mathematically.

  4. Note that at equilibrium the chemical reactions is still going on just as fast as during the approach to equilibrium. Equilibrium does not mean nothing is happening! What distinguishes the approach to equilibrium in terms of the two reaction rates from the system at equilibrium itself? (The reaction rate is the overall rate at which reactants react -- this is not the same as the reaction rate constant which is the probability that each collision will result in a reaction. You have to take into account the frequency of collisions.)

  5. If you reduce both reaction rate constants equally (e.g. divide both by two), what happens to the approach to equilibrium? And to the size of the fluctuations at equilibrium? To the equilibrium concentrations? (The last answer is "nothing"!)

  6. Along the same lines, can you distinguish the kinetics of this reaction from the thermodynamics? The "thermodynamics" is expressed by whether at equilibrium there are more products or reactants (i.e. whether K is large or small). What determines this? The "kinetics" is whether equilibrium is reached quickly or slowly. What determines this?

    How could you set the reaction rate constants so that this reaction is thermodynamically favorable (K is big) but kinetically essentially impossible (equilibrium is never reached)? How could you set up the system so that product was often formed (kinetically accessible) but the equilibrium concentration of product was small (thermodynamically unfavorable)? Would this be useful? Yes, if you had another chemical reaction which started with a product of this reaction. The overall thermodynamics of the two coupled reactions could be controlled by this second reaction; that is, the net reaction A + B --> C + D + E --> F + G could be quite favorable because of the second reaction C + D + E --> F + G. Then the overall reaction might go quite quickly. This illustrates the concept of a "reactive intermediate", a species which is thermodynamically disfavored, but which allows a reaction which is favored overall to proceed. Notice that it is critical that the intermediate be kinetically accessible.

  7. Suppose you change the density of the reactants. How does the reaction rate change? You can see why the reaction rate constant is a more useful quantity, in a sense, then the reaction rate. But the rate is easier to measure, of course.

    References

    Chemistry 1C, the first undergraduate UCI course that discusses chemical kinetics.

    Physical Chemistry Faculty at UCI, folks who do research on chemical kinetics.

    Return to the Chemical Kinetics Simulation.

    Return to the Instructional Applets page.

    Return to the Chemistry Instruction page.

    Return to the UCI Chemistry home page.

    cgrayce@limantour.ps.uci.edu