Equilibrium, Chemical

Equilibrium, Chemical

How chemical equilibrium works

Upsetting our equilibrium

Resources

Chemical equilibrium is the final condition of a chemical reaction after the reacting substances have completely reacted. Depending on the reaction, they may reach this condition quickly or slowly, but eventually they will come to a condition in which there are definite, unchanging amounts of all the relevant substances.

Chemical equilibrium is one of the most important features of chemical processes, involving everything from the dissolving of sugar in a cup of coffee to the vital reactions that carry essential oxygen to cells and poisonous carbon dioxide away from them.

How chemical equilibrium works

A chemical reaction occurs between substance A and substance B. The reaction can be written as:

A + B → C + D.

If the reaction is not one way, as indicated by the arrow, and if C and D can also react with one another to form A and B—the reverse of the process above—this process is indicated with the arrow going the other way:

A + B ← C + D.

This is far from an unusual situation. In fact, almost all chemical reactions are reversible, proceeding in either direction, as long as A, B, C, and D molecules are available. After all, a chemical reaction is a process in which atoms or electrons are rearranged, reshuffled, or transferred from one molecule to another. Chemists call a two-way reaction a reversible reaction, and they write the equation with a two way arrow:

In adverse reactions, A and B molecules produce C and D molecules, just as we wrote in the first equation above. As they react, the number of remaining As and Bs will decrease. But as more and more Cs and Ds are produced, they will begin to collide and reproduxce some of the As and Bs. The reaction is now going in both directions—to the right and to the left.

As more and more Cs and Ds build up, they will bump into each other more and more often, thereby reproducing As and Bs at a faster and faster clip. Pretty soon, the C + D (leftward) reaction will be going as fast as the A + B (rightward) reaction; As and Bs will reappear just as quickly as they disappear. The result is that there will be no further change in the numbers of As and Bs. The system is at equilibrium.

A and B molecules continue to collide and produce C and D molecules. Likewise, we see that Cs and Ds are still madly bumping into each other to produce As and Bs. In other words, both reactions are still vigorous. But they’re going on at exactly the same rate. If somebody is giving you money and taking it away at the same rate, the amount you have isn’t changing. The total amounts of A, B, C and D aren’t changing at all. It looks as if the reaction is over, and for all practical purposes, it is.

When the rate of leftward reaction is the same as the rate of rightward reaction, the system stays at equilibrium—no net movement in either direction.

Upsetting our equilibrium

Can we influence an equilibrium reaction in any way? Are we stuck with whatever mixture is left after a reaction has come to its final equilibrium condition? That is, are we stuck with so much of each substance, A, B, C and D? Or can we tamper with those ‘final’ equilibrium amounts? Can we force reactions to give us more of what we want (for example, C and D) and less of what we don’t want (for example,A and B)?

Suppose we have a reaction that has already come to equilibrium:

There are certain equilibrium amounts of all four chemicals. But if the number of B molecules is increased, the As and Bs will continue to react, because there are now more Bs for the As to collide into, extra Cs and Ds are produced, and fewer As remain. The same thing would happen if we added some extra As instead of Bs: we’d still get an extra yield of Cs and Ds.

Why would we want to do this? Well, maybe we’re a manufacturer of C or D, and the further we can push the reaction to produce more of them, the more we can sell. We have succeeded in pushing the equilibrium point in the direction we want—toward more C and D—by adding Bs alone; no extra As were needed. If we’d wanted to, we could have gotten the same result by adding some extra As instead of Bs, whichever one is cheaper.

We could also have accomplished the desired result—more C and D—not by speeding up the collision rate of As and Bs, but by slowing down the collision rate of Cs and Ds. That is, instead of adding extra As or Bs, we could keep removing some of the Cs or Ds from the pot as they are formed. Then the A and B will keep colliding and producing more, unhindered by the building up of the backward process. We’d get a 100% reaction between A and B.

In both cases, it appears that when we changed the amount of any one of the chemicals (by adding or removing some) we imposed an unnatural stress upon the equilibrium system. The system then tried to relieve that stress by using up some of the added substance or by producing more of the removed substance. In 1884, French chemist Henry Le Châtelier (1850-1936) realized that these shifts in the equilibrium amounts of chemicals were part of a bigger principle, which now goes by his name. In general terms, Le Châtelier’s principle can be stated this way: When any kind of stress is imposed upon achemical system that has already reached its equilibrium condition, the system will automatically adjust the amounts of the various substances in such a way as to relieve that stress as much as possible.

Le Châtelier’s principle is being applied every day in manufacturing industries to maximumize efficiency of chemical reactions.

Resources

BOOKS

Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002.

Parker, Sybil, ed. McGraw-Hill Encyclopedia of Chemistry.2nd ed. New York: McGraw Hill, 1999.

Umland, Jean B. General Chemistry. St. Paul: West, 1993.

OTHER

Oracle ThinkQuest. “Chemical Equilibrium” <http://library. thinkquest.org/10429/low/equil/equil.htm > (accessed November 24, 2006).

Steve Lower’s Website. “All about Chemical Equilibrium”<http://www.chem1.com/acad/webtext/chemeq/> (accessed November 24, 2006).

Robert L. Wolke