B. Disturbing Chemical Equilibrium: Le Chatelier's Principle

This is what you'll find in most textbooks:
In 1888 Henri-Lewis Le Chatelier (1850 - 1936) a French industrial chemist made the observation:
"Any change in one of the variables that determines the state of a system in equilibrium causes a shift in the position of equilibrium in a direction that tends to counteract the change in the variable under consideration."
Perhaps it would be better to have a less direct quote. Simply put, Le Chatelier's Principle states that a system in equilibrium responds to any stress by restoring the equilibrium. But even this can be confusing. If you think of what is needed to help out one side of the equilibrium over the other, you can easily predict what will happen.

In Class Examples

1. Each of the following reactions has reached equilibrium. What will be the effect on the equilibrium concentration of each substance when the change described is made?

a. 2 H₂ + 2 NO = N₂ + 2 H₂O

reacting hydrogen with a metal

Remember that at equilibrium, both reactions are still going on, but at the same rate. But if we suddenly remove hydrogen by getting it to react with a metal, we are hindering the forward reaction. Meanwhile the reverse reaction keeps proceeding, so that the amounts of nitrogen and water get depleted, and the amount of NO accumulates.

b. 2 SO₂₍_g)+ O_2(g)= 2 SO_3(g)

increasing the pressure on the system

The forward reaction involves 3 gaseous molecules at a time, whereas the reverse reaction involves only two. Increasing pressure will increase both the forward and reverse rates, but it will be a greater help to the reaction that involves more collisions. (the one involving more gaseous molecules). As a result, we will get more SO₃ forming and less SO₂ and O₂ will be present at a higher pressure.

Here’s another and better explanation:

2 SO₂₍_g) + O_2(g)= 2 SO_3(g)

Increasing pressure by compressing volume will increase the concentration of all gases.

Let’s say that concentration doubles.

If we look at the forward reaction the rate is given by

Rate forward= k[SO₂]²[O₂]. If each gas’ concentration doubles, then we will increase the rate of the forward reaction by a factor of 2²*2 = 8.

The reverse rate is given by:

Rate forward= k[SO₃]² If each gas’ concentration doubles, then we will increase the rate of the reverse reaction by only a factor of 2² = 4.

So clearly the forward rate is helped out more by an increase in pressure.

c. H₂O₍_g) = H₂O_(l) + heat

(1) cooling the system

The reverse reaction needs heat. Cooling will interfere with the reverse reaction. The forward reaction will continue, unhindered, depleting gaseous H₂O and causing more liquid H₂O to be formed.

(2) decreasing the pressure

Decreasing the pressure will make it more difficult for the gas H2O molecules to collide, bond and condense. Pressure will have no effect on the liquid, so it will continue to evaporate, and we'll end up with less liquid and more steam.

2. Consider the following reaction:

CaCO₃_(s) = CaO_{(s) +} CO₂_(g) DH = (+)

limestone lime

How would you maximize the amount of CaO produced ?

decrease pressure to discourage CO₂ from reacting
better still, remove the CO₂ as it forms
increase the temperature to encourage the forward endothermic reaction
continuously add limestone, assuming you have a way of separating it from the lime