Otherwise, you found yourself looking at a computer simulation of a binary chemical reaction. There are four types of molecules in this simulation, red, yellow, green and blue. Each wanders around at random in the box you see. Every time a pair of red and yellow molecules collide they may react to form a pair of green and blue molecules, and vice versa. This chemical reaction can be written like so:
Note: You are not watching an animation of a pre-computed simulation. This simulation is running in real time, on your computer, as you watch. Every time you restart the simulation, the molecules are distributed randomly in the box and given random velocities.
From the control panel you can also turn on and off a "stripchart" that records the instantaneous concentrations of the four species of molecule as the simulation proceeds. Just click on the button labeled "stripchart ON".
Experiment!
[ G ] [ B ] K = ----------- [ R ] [ Y ]
Remarkably, deducing the existence of thermodynamics by directly considering the limit of a simulation like this as the number of molecules is increased to infinity is profoundly difficult, and I am not actually sure it has ever been done. We must accept thermodynamics at present on the basis of experience and intuition alone. (If you would like learn more about thermodynamics, consider visiting The Second Law.)
The fluctuations have many other important roles and meanings. I will mention just one more: the fluctuations 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) they tell you how the system behaves dynamically when it is not at equilibrium (but fairly close). Perhaps by comparing (in this simulation) how fluctuations away from equilibrium happen, and how the system approaches equilibrium -- watch the strip chart -- you will be able to see intuitively the truth of this fundamental and important observation, which is called the Fluctuation-Dissipation Theorem and is a bit of a trick to prove mathematically.
How could you set up this simulation so that the forward reaction would be thermodynamically favorable (K is big, so there'll be a lot of products at equilibrium) but kinetically nearly impossible (equilibrium will take a very long time to be reached)? Try it and see!
How could you set up the system so that product was often formed (kinetically accessible) but the equilibrium concentration of the product was small (thermodynamically unfavorable)? Would this be useful? Yes! If you had another chemical reaction which started with a product of this reaction, then the overall thermodynamics of the two coupled reactions could be controlled by the second reaction. That is, the net reaction R + Y --> G + B + X --> Y + Z could be favorable because of the second reaction G + B + X --> Y + Z. Since the intermediate products G and B are formed quickly it does not actually matter that very few of them are around at any given moment; quite a lot of the product Y + Z will nevertheless be formed. This illustrates the concept of a reactive intermediate. Notice that it is critical that this intermediate be kinetically accessible.
Physical Chemistry Faculty at UCI, folks who do research on chemical kinetics, among other things.
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