Le Chatelier's Principle: Getting More (or Less) Product - Chemical Reactions - Chemistry Essentials for Dummies

Chemistry Essentials for Dummies

Chapter 7. Chemical Reactions

Le Chatelier's Principle: Getting More (or Less) Product

If you’re, say, an industrial chemist, you want as much of the reactants as possible to be converted to product. You’d like the reaction to go to completion (meaning you’d like the reactants to keep creating the product until they’re all used up), but that doesn’t happen for an equilibrium reaction. But it would be nice if you could, in some way, manipulate the system to get a little bit more product formed. There is such a way — through Le Chatelier’s Principle.

A French chemist, Henri Le Chatelier, discovered that if you apply a change of condition (called stress) to a chemical system that’s at equilibrium, the reaction will return to equilibrium by shifting in such a way as to counteract the change (the stress). This is called Le Chatelier’s Principle.

You can stress an equilibrium system in three ways:

Change the concentration of a reactant or product.

Change the temperature.

Change the pressure on a system that contains gases.

In this section, you see how this applies to the Haber process:

Changing the concentration

REMEMBER. In general, if you add more of a reactant or product to an equilibrium system, the reaction will shift to the other side to use it up. If you remove some reactant or product, the reaction shifts to that side in order to replace it.

Suppose that you have the ammonia system at equilibrium, and you then put in some more nitrogen gas. To reestablish the equilibrium, the reaction shifts from the left to the right, using up some nitrogen and hydrogen, and forming more ammonia and heat.

The equilibrium has been reestablished. You have less hydrogen and more nitrogen, ammonia, and heat than you had before you added the additional nitrogen. The same thing would happen if you had a way of removing ammonia as it was formed. The right-hand side of the teeter-totter would again be lighter, and weight would be shifted to the right in order to reestablish the equilibrium. Again, more ammonia would be formed.

Changing the temperature

REMEMBER. In general, heating a reaction causes it to shift to the endothermic (heat-absorbing) side. (If you have an exothermic reaction where heat is produced on the right side, then the left side is the endothermic side.) Cooling a reaction mixture causes the equilibrium to shift to the exothermic (heat-releasing) side.

Suppose that you heat the reaction mixture of nitrogen and hydrogen. You know that the reaction is exothermic — heat is given off, showing up on the right-hand side of the equation. So if you heat the reaction mixture, the reaction shifts to the left to use up the extra heat and reestablish the equilibrium. This shift uses up ammonia and produces more nitrogen and hydrogen. And as the reaction shifts, the amount of heat also decreases, lowering the temperature of the reaction mixture.

Changing the pressure

REMEMBER. Changing the pressure affects the equilibrium only if there are reactants and/or products that are gases. In general, increasing the pressure on an equilibrium mixture causes the reaction to shift to the side containing the fewest number of gas molecules.

In the Haber process, all species are gases, so you do see a pressure effect. Think about the sealed container where your ammonia reaction is occurring. (The reaction has to occur in a sealed container because everything is a gas.) You have nitrogen, hydrogen, and ammonia gases inside. There is pressure in the sealed container, and that pressure is due to the gas molecules hitting the inside walls of the container.

Now suppose that the system is at equilibrium, and you want to increase the pressure. You can do so by making the container smaller (with a piston type of arrangement) or by putting in nonreactive gas, such as neon. You get more collisions on the inside walls of the container, and therefore, you have more pressure. Increasing the pressure stresses the equilibrium; to remove that stress and reestablish the equilibrium, the pressure must be reduced.

Take another look at the Haber reaction and look for some clues on how this may happen.

Every time the forward (left-to-right) reaction takes place, four molecules of gas (one nitrogen and three hydrogen) form two molecules of ammonia gas. This reaction reduces the number of molecules of gas in the container. The reverse reaction (right-to-left) takes two ammonia gas molecules and makes four gas molecules (nitrogen and hydrogen). This reaction increases the number of gas molecules in the container.

The equilibrium has been stressed by an increase in pressure; reducing the pressure will relieve the stress. Reducing the number of gas molecules in the container will reduce the pressure (fewer collisions on the inside walls of the container), so the forward (left-to-right) reaction is favored because four gas molecules are consumed and only two are formed. As a result of the forward reaction, more ammonia is produced!