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Le Chatelier's Principle
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This states that when a change to the conditions of an equilibrium occurs, then the equilibrium will shift in such a way as to resist that change. |
Put in a non-scientific way, this is really an application of Sod's Law - it says equilibria are awkward things! We can use the well-known equilibrium for the Haber process to illustrate this principle:
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N2(g) + 3H2(g) |
DH = -92.2 kJ mol-1 |
This reaction is exothermic, or to be more precise the reaction from left to right is exothermic. As it is a reversible reaction, the reverse reaction (from right to left) must be endothermic. Various changes can be applied to this chemical equilibrium. If more of one of the reactants (nitrogen or hydrogen) is added, the equilibrium will shift to the right, so removing some of the added reactant:
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N2(g) + 3H2(g) ⇒ 2NH3(g) |
If some more product (ammonia) is added, the equilibrium will shift to the left in order to remove some of the added ammonia:
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N2(g) + 3H2(g) ¬ 2NH3(g) |
Removing one of the reactants causes the equilibrium to shift to the left, in order to replace some of the removed reactant. In the same way, removing some of the product causes the reaction to move to the right to make some more product.
Increasing the temperature of the equilibrium will cause the endothermic reaction to take place, which will cool down the equilibrium. In this case, the equilibrium will shift to the left:
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N2(g) + 3H2(g) ¬ 2NH3(g) |
When dealing with pressure, it is important to realize that the appearance of gases will cause a significant increase in pressure. So if we increase the pressure, the equilibrium will shift in a direction which leads to a reduction in the total number of gas molecules. In this case, an increase in pressure causes the equilibrium to shift to the right, as the equation shows that this leads to a reduction from four moles of gas to two:
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N2(g) + 3H2(g) ⇒ 2NH3(g) |
In the industrial production of ammonia, we use a pressure of around 200 atmospheres, as this leads to more ammonia being produced without the expense of extremely high-pressure equipment. It may, at first, seem odd that a temperature of around 500°C is used, as this will lead to a smaller yield of ammonia than would be achieved at room temperature. However, the rate of reaction is around 300 million, million times faster at this temperature! There may be more ammonia produced at room temperature, but it takes an awful long time! Removing the ammonia by liquefaction also drives the equilibrium to the right.
2NH3(g)