Reversible Reactions & Dynamic Equilibrium
Reversible Reactions
Most chemical reactions go to completion — reactants are fully converted into products and the reaction stops. However, some reactions are reversible. In a reversible reaction, the products can react together to re-form the original reactants. The reaction can go in both directions.
Reversible reactions are shown using the equilibrium symbol \(\rightleftharpoons\) instead of a normal arrow. For example, when ammonium chloride is heated it decomposes into ammonia and hydrogen chloride, but when the gases cool they recombine:
$$\text{NH}_4\text{Cl}(s) \rightleftharpoons \text{NH}_3(g) + \text{HCl}(g)$$
A reversible reaction can proceed in both the forward direction (reactants \(\rightarrow\) products) and the backward direction (products \(\rightarrow\) reactants). The \(\rightleftharpoons\) symbol shows this. The direction of some reversible reactions can be altered by changing the reaction conditions.
Hydrated copper sulfate is another example you may see in practicals. Blue hydrated copper sulfate turns white when heated (losing water), but adding water to the white anhydrous powder turns it blue again:
$$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}(s) \rightleftharpoons \text{CuSO}_4(s) + 5\text{H}_2\text{O}(l)$$
Dynamic Equilibrium
When a reversible reaction takes place in a closed system (where no substances can enter or leave), it can reach a state of dynamic equilibrium.
At dynamic equilibrium:
- The forward reaction and the backward reaction are both still happening.
- The rate of the forward reaction equals the rate of the backward reaction.
- The concentrations of reactants and products remain constant (but are not necessarily equal).
At dynamic equilibrium, both reactions are still taking place — it is not that the reactions have stopped. The word “dynamic” means the particles are still reacting in both directions. The word “equilibrium” means the rates are balanced, so there is no overall change in the amounts of reactants and products.
The Haber Process
The Haber process is an important industrial example of a reversible reaction. It produces ammonia from nitrogen (extracted from the air) and hydrogen (obtained from natural gas):
$$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$$
This reaction can reach dynamic equilibrium. The conditions used in the Haber process are a compromise chosen to give a reasonable yield at a reasonable rate:
The three conditions you need to recall are:
- Temperature: 450 °C
- Pressure: 200 atmospheres
- Catalyst: Iron
The iron catalyst does not change the position of equilibrium — it speeds up both the forward and backward reactions equally, so equilibrium is reached faster.
Changing the Position of Equilibrium
The position of equilibrium describes the relative amounts of reactants and products at equilibrium. If conditions change, the system shifts to counteract the change and establish a new equilibrium position.
If the forward reaction is exothermic, increasing the temperature shifts the equilibrium in the endothermic direction (backwards), producing more reactants. Decreasing the temperature shifts it in the exothermic direction (forwards), producing more products. For the Haber process, the forward reaction is exothermic, so a lower temperature would give a higher yield of ammonia — but the rate would be too slow, hence the compromise of 450 °C.
Increasing the pressure shifts the equilibrium towards the side with fewer moles of gas. In the Haber process, there are 4 moles of gas on the left (1 \(\text{N}_2\) + 3 \(\text{H}_2\)) and 2 moles on the right (2 \(\text{NH}_3\)). Higher pressure favours the forward reaction, increasing the yield of ammonia. This is why 200 atmospheres is used.
Increasing the concentration of a reactant shifts the equilibrium towards the products (to the right), because the system acts to reduce the increased concentration by using it up. Increasing the concentration of a product shifts the equilibrium towards the reactants (to the left).
When conditions change, the equilibrium shifts to oppose the change. Increasing temperature favours the endothermic direction. Increasing pressure favours the side with fewer gas molecules. Increasing concentration of a substance shifts equilibrium away from that substance.