Reversible Reactions and Le Chatelier's Principle
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 using the equilibrium symbol \(\rightleftharpoons\) instead of a single arrow.
A well-known example is the dehydration of hydrated copper sulfate: blue hydrated copper sulfate (\(CuSO_4 \cdot 5H_2O\)) loses water when heated to form white anhydrous copper sulfate (\(CuSO_4\)). Adding water to the white powder restores the blue colour. Another example is the decomposition of ammonium chloride: \(NH_4Cl(s) \rightleftharpoons NH_3(g) + HCl(g)\). When heated, it decomposes into gases; when cooled, the gases recombine to form solid ammonium chloride again.
In a reversible reaction, if the forward reaction is exothermic, the reverse reaction is endothermic by the same amount of energy. The two energy changes are equal in magnitude but opposite in sign.
Reversible reactions are significant in many biological and industrial contexts. The reaction between haemoglobin and oxygen in the blood is reversible: at the lungs, where oxygen concentration is high, haemoglobin combines with oxygen to form oxyhaemoglobin; at the tissues, where oxygen concentration is low, the reverse reaction releases oxygen to the cells. This is a natural example of Le Chatelier’s principle operating in a biological system — the equilibrium position shifts depending on the local conditions. Understanding reversibility is essential for designing efficient chemical manufacturing processes, because it helps chemists choose conditions that maximise the yield of the desired product while minimising energy use and waste.
Dynamic Equilibrium
When a reversible reaction takes place in a closed system, a dynamic equilibrium is eventually reached. At equilibrium:
- The rate of the forward reaction equals the rate of the reverse reaction.
- The concentrations of reactants and products remain constant (but are not necessarily equal to each other).
- Both reactions are still occurring at the molecular level — this is why the equilibrium is described as “dynamic.”
Equilibrium does not mean the amounts of reactants and products are equal — it means they are constant. The equilibrium position can strongly favour the products or the reactants. Equilibrium is not a static state — both forward and reverse reactions continue, but at equal rates, so the net composition does not change. Equilibrium can only be reached in a closed system; if products escape, the reaction continues in the forward direction until completion.
The position of equilibrium — how far the reaction proceeds towards products — is described qualitatively. We say the equilibrium “lies to the right” if products predominate, or “lies to the left” if reactants predominate. In many industrial reactions, chemists deliberately choose conditions that push the equilibrium position as far to the right as possible while keeping the rate acceptable. This involves an understanding of both kinetics (rate) and thermodynamics (equilibrium) working together, which is why industrial chemistry is a careful balancing act between speed and yield.
Le Chatelier’s Principle (Higher Tier)
Le Chatelier’s principle states that if the conditions of a system at equilibrium are changed, the equilibrium position shifts in the direction that opposes the change. This allows us to predict how temperature, concentration, and pressure affect the yield of a reaction.
Temperature: Increasing temperature shifts the equilibrium in the endothermic direction (to absorb the extra heat). Decreasing temperature shifts it in the exothermic direction. This changes the yield of the reaction.
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 it towards the reactants. Removing a product also shifts equilibrium to the right (towards products).
Pressure (reactions involving gases): Increasing pressure shifts equilibrium towards the side with fewer moles of gas (to reduce the pressure). Decreasing pressure shifts it towards the side with more moles of gas. Count the total moles of gas on each side to predict the direction of shift.
Catalyst: A catalyst does not change the equilibrium position — it speeds up both forward and reverse reactions equally, so equilibrium is reached faster but the final proportions of reactants and products are unchanged. This is a frequently tested distinction.
The Haber Process: A Le Chatelier’s Example
The Haber process is the most important industrial application of Le Chatelier’s principle. It produces ammonia, which is essential for making fertilisers, explosives and cleaning products:
$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad \Delta H = -92 \text{ kJ/mol}$$
The forward reaction is exothermic and there are 4 moles of gas on the left and 2 moles on the right. Applying Le Chatelier’s principle: a lower temperature favours the exothermic direction (more ammonia), but the rate would be unacceptably slow. A higher pressure favours the side with fewer moles of gas (the products), increasing yield, but very high pressures are dangerous and expensive to maintain.
Industrial conditions are a compromise: approximately 450°C (high enough for acceptable rate, low enough for reasonable yield), approximately 200 atm (high yield without excessive cost and safety risks), and an iron catalyst (increases rate without changing yield). Unreacted nitrogen and hydrogen are recycled through the reactor to maximise efficiency.