Le Chatelier's Principle

It seems that by raising the pressure of the system and at the same time reducing its temperature we can maximise the yield of ammonia at equilibrium. This is not without its problems. Maintaining a higher pressure is expensive; more robust plant (thicker pipes and stronger joints, for example) is required, and more electricity would have to be used to drive pumps to maintain this pressure. Lowering the temperature slows the rate at which chemical reactions take place resulting in a longer wait for equilibrium to be achieved. A 'compromise' has to be settled upon. The conditions selected approximate to those given above.

Does a catalyst affect the equilibrium position?

No. The addition or removal of a catalyst does not cause a shift in equilibrium. A catalyst is a substance that affects the rate of a reaction, but is not consumed in the reaction. It does this by providing an alternative mechanism for the reaction but of lower activation enthalpy. It therefore changes the rate of approach of equilibrium and so affects the forward and reverse reaction rates in the same way. The composition of the equilibrium mixture is unchanged. 

Heterogeous Equilibrium


An equilibrium involving more than one phase - gas and solid, for example, or liquid and solid - is said to be heterogeneous. For example:

H₂(g) + I₂(s)        2HI(g)

In the above example, the equilibrium expressions for this reaction are:
Notice that I₂ is missing from the equilibrium expressions. This is because, at room temperature, iodine is a solid. The composition of the equilibrium mixture of H₂(g) and HI(g) is independent of the amount of solid iodine present, as long as some of it is always present. [The partial pressure of I₂(g) (which is its vapour pressure) in equilibrium with I₂(s) is constant at a constant temperature, and is independent of the volume of the container and of the quantity of solid.] 

Equilibrium in Solution

Here we will consider homogeneous and heterogeneous chemical equilibria in which the solvent is much more abundant than all the other components put together. The solvent may also be a reactant or product itself.
First of all, let's look at an equilibrium that is not in solution: 
 

CH₃COOH(l) + CH₃CH₂OH(l)        CH₃COOCH₂CH₃(l) + H₂O(l)

 In this case, water is not a solvent but a product only.
Now consider the same equilibrium in very dilute aqueous solution. The chemical equation is: 

CH₃COOH(aq) + CH₃CH₂OH(aq)        CH₃COOCH₂CH₃(aq) + H₂O(l)

And the equilibrium expression is:
Water is now both a product and a solvent. In reaching a state of equilibrium the concentration of the water changes so little that it is effectively constant. For this reason its value is taken along with the equilibrium constant, K_c.

Here are some more examples of chemical equilibria in aqueous solution:

CH₃COOH_(aq) + H₂O_(l)        CH₃COO^-_(aq) + H₃O⁺_(aq)

K_a is called the acid dissociation constant. Again, the concentration of water changes negligibly on reaching equilibrium and so its value is taken along with the equilibrium constant (K_c) to form the 'modified' equilibrium constant, K_a.

NH_3(aq) + H₂O_(l)        NH₄⁺_(aq) + OH^-_(aq)

K_b is called the base dissociation constant. Again, the concentration of water changes negligibly on reaching equilibrium and so its value is taken along with the equilibrium constant (K_c) to form the 'modified' equilibrium constant, K_b.

AgCl_(s)        Ag⁺_(aq) + Cl^-_(aq)

K_s is the Solubility Product. It is the equilibrium constant for the equilibrium that exists between a slightly soluble salt and its ions in saturated solution. The solid, AgCl, is omitted from the equilibrium expression. As long as there is some solid present the concentration of the saturated solution of ions remains constant at a given temperature. 

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