Equilibria

Reversible Reactions

Most chemical reactions go to completion — the reactants are used up and only products remain. However, some reactions are reversible. In a reversible reaction, the products can react to re-form the original reactants. This is shown by the equilibrium symbol $\rightleftharpoons$ instead of a single arrow. Many important industrial and biological reactions are reversible, and understanding how to control them is essential for maximising product yield in chemical manufacturing.

For example, the decomposition of ammonium chloride: $NH_4Cl(s) \rightleftharpoons NH_3(g) + HCl(g)$. When heated, ammonium chloride breaks down into ammonia and hydrogen chloride gases. When cooled, these gases recombine to form solid ammonium chloride again. The forward reaction is endothermic (absorbs energy); the reverse reaction is exothermic (releases energy). Another example is the reaction of hydrated copper sulfate: blue $CuSO_4 \cdot 5H_2O$ turns white (anhydrous $CuSO_4$) when heated, and turns blue again when water is added.

Key Fact

In a reversible reaction, if the forward reaction is exothermic, the reverse reaction is endothermic (and vice versa). The energy change in each direction has the same magnitude but opposite sign.

Dynamic Equilibrium

When a reversible reaction takes place in a closed system (where nothing can enter or leave), a dynamic equilibrium is eventually reached. At equilibrium, three conditions are met:

• The rate of the forward reaction equals the rate of the reverse reaction.

• The concentrations of reactants and products remain constant (but not necessarily equal to each other).

• Both reactions are still occurring at the molecular level — this is why it is called “dynamic.” Molecules are continuously reacting in both directions, but the overall proportions do not change.

Equilibrium does not mean the amounts of reactants and products are equal — it means they are constant. The equilibrium position can favour the products (more products than reactants) or the reactants (more reactants than products). Also, equilibrium is not a static state — both reactions continue at equal rates, like two escalators running in opposite directions at the same speed. You can only achieve equilibrium in a closed system; in an open system, products can escape and equilibrium is never reached.

Le Chatelier’s Principle (Higher)

Le Chatelier’s principle states that if the conditions of a system at equilibrium are changed, the equilibrium position shifts to oppose the change and partially counteract it. This principle allows us to predict which direction the equilibrium will move when we alter temperature, concentration or pressure.

Temperature: Increasing temperature favours the endothermic direction (to absorb the extra heat and reduce the temperature). Decreasing temperature favours the exothermic direction. This changes both the rate and the equilibrium position. Concentration: Increasing the concentration of a reactant shifts equilibrium towards the products (to use up the extra reactant). Increasing the concentration of a product shifts equilibrium towards the reactants. Removing a product also shifts the equilibrium to the right. Pressure (for reactions involving gases): Increasing pressure favours the side with fewer gas molecules (to reduce the pressure). Decreasing pressure favours the side with more gas molecules. Count the total moles of gas on each side to predict the shift.

A catalyst does not change the equilibrium position — it speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but the final proportions of reactants and products are unchanged. This is an important distinction that is frequently tested.

The Haber Process

The Haber process is the most important industrial example of Le Chatelier’s principle in action. It produces ammonia, which is essential for making fertilisers, explosives and cleaning products. The equation is:

$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$

The forward reaction is exothermic ($\Delta H = -92$ kJ/mol). There are 4 moles of gas on the left and 2 moles on the right. Applying Le Chatelier’s principle: a low temperature would favour the exothermic (forward) direction and increase yield, but the rate would be unacceptably slow. A high pressure would favour the side with fewer moles of gas (the products), increasing yield, but very high pressures are dangerous and expensive. Industrial conditions are therefore a compromise: temperature of about 450°C (high enough for a reasonable rate), pressure of about 200 atm (high enough for a good yield without excessive cost), and an iron catalyst (speeds up the reaction so equilibrium is reached faster without changing the position). Unreacted nitrogen and hydrogen are recycled back through the reactor to maximise efficiency, meaning that even though the percentage conversion per pass is only about 15%, the overall yield is much higher over time. Ammonia is a crucial chemical: it is used to make ammonium nitrate fertiliser, nitric acid, and various cleaning products.