Dynamic equilibrium in reversible reactions, the effect of concentration, pressure, temperature and catalysts on the position of equilibrium, Le Chatelier's principle, and how industry uses these to maximise yield.

What this key area is asking

The SQA wants you to describe dynamic equilibrium in a reversible reaction, use Le Chatelier's principle to predict the effect of changing concentration, pressure and temperature, explain the role of a catalyst, and apply these ideas to industrial processes such as the Haber process. The "predict and explain" style means a bare prediction earns little; the marks are in the reasoning.

Dynamic equilibrium

Le Chatelier's principle

• Concentration: adding more reactant shifts the equilibrium towards the products; removing product also shifts it towards the products.
• Pressure (gases): increasing the pressure shifts the equilibrium towards the side with fewer moles of gas.
• Temperature: raising the temperature shifts the equilibrium in the endothermic direction; lowering it favours the exothermic direction.

Catalysts at equilibrium

A catalyst increases the rate of both the forward and reverse reactions equally. It therefore lets the system reach equilibrium more quickly but does not change the position of equilibrium or the final yield.

Worked example: predicting a shift quantitatively

Industrial application

Industry chooses conditions to balance yield, rate and cost. For an exothermic reaction, a lower temperature gives a higher yield but a slower rate, so a compromise (moderate) temperature is used with a catalyst to keep the rate economical. High pressure can raise the yield where fewer gas moles are formed, but it is expensive, so a compromise pressure is chosen.

Examples in context

The Haber process for ammonia ($N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$) is the textbook case: it runs at about $450 \,^{\circ}\text{C}$ and $200 \text{ atm}$ with an iron catalyst, a deliberate compromise that sacrifices some equilibrium yield for a workable rate. The ammonia made this way is the basis of nitrogen fertilisers that feed roughly half the world's population. The Contact process for sulfuric acid uses the same Le Chatelier reasoning, removing $SO_3$ and using a vanadium(V) oxide catalyst at a moderate temperature. Even fizzy drinks illustrate equilibrium: $CO_2(g) \rightleftharpoons CO_2(aq)$ is bottled under high pressure to push the equilibrium towards dissolved gas, and opening the can drops the pressure so the equilibrium shifts back and the drink fizzes.

Try this

Q1. Predict the effect of increasing the pressure on $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$. [2 marks]

• Cue. It shifts the equilibrium to the right (4 moles of gas to 2 moles), increasing the yield of ammonia.

Q2. Explain why a catalyst does not change the yield at equilibrium. [2 marks]

• Cue. It speeds up the forward and reverse reactions equally, so equilibrium is reached faster but the position is unchanged.

Q3. For the exothermic reaction $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$, predict and explain the effect of lowering the temperature on the yield of $SO_3$. [2 marks]

• Cue. The equilibrium shifts in the exothermic (forward) direction to oppose the cooling, so the yield of $SO_3$ increases.