Chemistry: Disturbing Equilibria: Le Châtelier's Principle
Disturbing Equilibria: Le Châtelier's Principle
Back in the late 1800s, the French chemist Henri Le Châtelier came up with a rule about equilibria that we still use extensively today, to the chagrin of many chemistry students.
Le Châtelier's Principle states that if you change the conditions of an equilibrium, the equilibrium will shift in a way that minimizes the effects of whatever it is you did.
The Mole Says
Le Châtelier's Principle states that the equilibrium will shift to minimize the effects of whatever you did. However, it's important to note that the concentrations of the chemical species will be different in the new equilibrium than they were before. In essence, a new equilibrium will be created in order to maintain Keq for the reaction.
In other words, equilibria are like obnoxious little kids. For example, if you yell at a little kid, the kid will change his behavior to minimize your yelling. Likewise, if you change the conditions of an equilibrium, it will change in a way that partially undoes whatever it is you did to it in the first place.
Change in Concentration
Let's say that you're doing a reaction with the equation A + B ⇔ C. Le Châtelier's principle says that if we change the concentration, the position of the equilibrium will also change.
For example, if the process explained earlier is at equilibrium, we can disturb the equilibrium by adding a bunch of compound A to the reaction. To minimize the effects of the added compound, the equilibrium will shift in a way that will decrease the amount of compound A—namely, it will produce more of compound C. Likewise, by adding more of compound C, the equilibrium will be pushed toward the left, making more of A and B.
This phenomenon is often seen when an ionic compound dissolves. Let's see an example:
- CaSO4(s) ⇔ Ca+2(aq) + SO4-2(aq)
The Ksp value for calcium sulfate is 2.4 × 10-5. Doing the math we learned earlier in this section, we can easily find the equilibrium concentration of the calcium ion:
- Ksp = 2.4 × 10-5 = [x][x]
- x = 4.9 × 10-3 M
However, what would happen if we added 1.0 M Li2SO4? Because the concentration of the sulfate ion would be increased by 1.0 M, the quantity of the calcium ion would now be:
- Ksp = 2.4 × 10-5 = [x M][x + 1.0 M]
where the concentration of the sulfate ion would be (x + 0.10 M) to compensate for the added sulfate ion. Because the quantity of sulfate ion likely to dissolve is very small compared with the quantity we've added, we'll eliminate the "x" from this term to give us:
- 2.4 × 10-5 = (x M)(1.0 M)
- x = 2.4 × 10-5 M.
As you can see, the quantity of calcium ion has been greatly reduced by the addition of 1.0 M sulfate ion.
The reduced solubility of one compound caused by adding an additional quantity of one of its ions to solution is referred to as the common ion effect.
The common ion effect is when the addition of an ion affects the solubility or reactivity of a chemical compound.
Change in Pressure
Gaseous equilibria can be changed by altering the pressure of the gases. For example, let's say that we're doing the reaction A(g) + B(g) ⇔ C(g). In this reaction, two moles of gas are combining to make one mole of gas.
If we increase the pressure of this mixture of gases by squishing it into a smaller area, the pressure of each of the gases will increase. Le Châtelier's principle states that equilibria tend to want to decrease the effects of any changes, so the equilibrium will shift in a way that reduces the pressure of the system. The only way to accomplish this is to have fewer moles of gas present; as a result, the reaction will shift toward products to reduce the overall pressure.
There's another way to increase the pressure of a gaseous mixture: Add another gas that doesn't interact with any of the gases that are in equilibrium (for our example, imagine adding some of gas D to the mixture). It's important to keep in mind that although the total pressure inside the container will increase, the partial pressures of each gas will not. As a result, the addition of gas D won't change the position of the equilibrium at all.
Change in Temperature
Some reactions naturally give off energy (they're exothermic) and some reactions require energy to take place (they're endothermic). If we think of the energy in an exothermic reaction as being a product, it will have the form:
- A ⇔ B + energy.
Because endothermic reactions require energy to take place, we can think of energy as a reagent:
- A + energy ⇔ B
When we increase the temperature of a chemical reaction, what we're really doing is adding energy to it. Because we can think of energy as being either a reactant or product, the addition of extra energy will disrupt the equilibrium. For exothermic reactions, the addition of energy will push the reaction to the left toward reactants. For endothermic reactions, the addition of energy will push the reaction to the right toward the products.
Excerpted from The Complete Idiot's Guide to Chemistry © 2003 by Ian Guch. All rights reserved including the right of reproduction in whole or in part in any form. Used by arrangement with Alpha Books, a member of Penguin Group (USA) Inc.