C22: Reversible Reactions and Equilibrium
How reversible reactions reach dynamic equilibrium, and how Le Chatelier's principle lets us predict the effect of changing conditions on the position of equilibrium.
Reversible Reactions
A reversible reaction is a reaction where the products can react together to reform the original reactants. It is shown using the ⇌ symbol.
Example of a reversible reaction
Ammonium chloride ⇌ ammonia + hydrogen chloride
NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
On heating, ammonium chloride decomposes into ammonia and hydrogen chloride gases. On cooling, these gases recombine to form ammonium chloride.
In a reversible reaction:
- The forward reaction goes from left to right (reactants → products)
- The backward reaction (or reverse reaction) goes from right to left (products → reactants)
If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction. The same amount of energy is transferred in each direction.
Worked Example: Energy transfer in reversible reactions
The thermal decomposition of hydrated copper(II) sulfate:
CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)
The forward reaction (removing water) is endothermic — energy must be supplied by heating.
The reverse reaction (adding water to anhydrous copper sulfate) is exothermic — energy is released, and the sulfate gets hot.
Dynamic Equilibrium
When a reversible reaction takes place in a closed system (no substances can enter or leave), the reaction reaches a state of dynamic equilibrium.
At dynamic equilibrium:
- The forward and backward reactions are still happening, but at the same rate
- The concentrations of reactants and products remain constant (they do not change)
- The reaction has not stopped — it is dynamic, meaning both reactions continue
Dynamic equilibrium can only be reached in a closed system. If products can escape (e.g. a gas leaves an open container), the reaction will go to completion rather than reaching equilibrium.
Le Chatelier's Principle
Le Chatelier's principle states that if a system at equilibrium is subjected to a change in conditions, the system responds to counteract the change and re-establish equilibrium.
The position of equilibrium shifts to oppose the change you make.
Effect of Concentration
If the concentration of a reactant is increased, the equilibrium shifts to the right (towards the products) to use up the extra reactant. If the concentration of a product is increased, the equilibrium shifts to the left (towards the reactants).
Worked Example: Concentration change
For the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
If the concentration of nitrogen is increased, the system opposes the change by shifting to the right to reduce the nitrogen concentration. This produces more ammonia.
If ammonia is removed from the system, the equilibrium shifts to the right to replace it, producing more ammonia.
Effect of Temperature
If the temperature is increased, the equilibrium shifts in the direction of the endothermic reaction to absorb the extra heat. If the temperature is decreased, the equilibrium shifts in the direction of the exothermic reaction.
Worked Example: Temperature change
For the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (forward reaction is exothermic)
If temperature is increased, the equilibrium shifts in the endothermic direction (to the left) to absorb the extra heat. This reduces the yield of ammonia.
If temperature is decreased, the equilibrium shifts in the exothermic direction (to the right) to produce more heat. This increases the yield of ammonia.
Effect of Pressure
If the pressure is increased, the equilibrium shifts towards the side with fewer moles of gas to reduce the pressure. If the pressure is decreased, the equilibrium shifts towards the side with more moles of gas.
Worked Example: Pressure change
For the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Left side: 1 + 3 = 4 moles of gas. Right side: 2 moles of gas.
If pressure is increased, the equilibrium shifts to the side with fewer moles (the right), producing more ammonia.
If pressure is decreased, the equilibrium shifts to the side with more moles (the left), reducing ammonia yield.
| Change | Equilibrium Shifts To | Effect on Yield of Product |
|---|---|---|
| Increase concentration of reactant | Right (towards products) | Increases |
| Increase concentration of product | Left (towards reactants) | Decreases |
| Increase temperature | Endothermic direction | Depends on which direction is endothermic |
| Decrease temperature | Exothermic direction | Depends on which direction is exothermic |
| Increase pressure | Side with fewer gas moles | Depends on which side has fewer moles |
| Decrease pressure | Side with more gas moles | Depends on which side has more moles |
The Haber Process
The Haber process makes ammonia from nitrogen and hydrogen. It is an important application of equilibrium principles in industry.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol (exothermic)
The nitrogen is obtained from the air and the hydrogen from natural gas (methane). The conditions used are a compromise:
Temperature Compromise
- Low temperature favours the exothermic forward reaction — higher yield of ammonia
- But low temperature means a very slow rate of reaction
- Compromise: 450 °C — gives a reasonable yield at an acceptable rate
Pressure Compromise
- High pressure favours the side with fewer gas moles (4 → 2) — shifts equilibrium right, higher yield
- But very high pressure is expensive and dangerous (stronger equipment needed)
- Compromise: 200 atm — gives a reasonable yield without excessive costs
Catalyst
- Iron catalyst is used to increase the rate of reaction
- It does not change the position of equilibrium or the yield
- It helps the system reach equilibrium faster
Only about 15% of nitrogen and hydrogen convert to ammonia per pass. Unreacted gases are recycled back through the reactor, giving an overall conversion of about 97%.
Worked Example: Predicting equilibrium shifts in the Haber process
Question: What happens to the yield of ammonia if the pressure is increased to 300 atm?
Answer: Increasing pressure shifts equilibrium towards the side with fewer moles of gas. There are 4 moles on the left and 2 moles on the right, so equilibrium shifts right. The yield of ammonia increases.
Question: What happens if the temperature is increased to 600 °C?
Answer: The forward reaction is exothermic. Increasing temperature shifts equilibrium in the endothermic direction (to the left). The yield of ammonia decreases. However, the rate of reaching equilibrium would be faster.
Predicting Equilibrium Shifts — Step by Step
When asked to predict the effect of a change on equilibrium:
- Identify the change (concentration, temperature, or pressure)
- Apply Le Chatelier's principle — the system opposes the change
- State the direction of the shift (left or right)
- State the effect on the yield of the desired product
Worked Example: Predicting equilibrium shift
For the equilibrium: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (forward reaction exothermic)
Question: What is the effect of increasing the pressure?
Step 1: Count moles — left: 2 + 1 = 3 moles, right: 2 moles
Step 2: Increasing pressure shifts equilibrium towards fewer moles (right)
Step 3: Equilibrium shifts right
Step 4: Yield of SO₃ increases
Worked Example: Predicting equilibrium shift
For the equilibrium: CH₃COOH(l) + C₂H₅OH(l) ⇌ CH₃COOC₂H₅(l) + H₂O(l)
Question: What is the effect of removing water as it forms?
Step 1: The concentration of a product (water) is being decreased
Step 2: The system opposes this by shifting right to replace the water
Step 3: Equilibrium shifts right
Step 4: More ester is produced
When answering equilibrium questions, always state: (1) which direction the equilibrium shifts, (2) why it shifts in that direction, and (3) what happens to the yield. Do not just say "it changes" — be specific about the direction.
Practice Questions
1. Explain what is meant by dynamic equilibrium. (3 marks)
In a closed system, the forward and backward reactions occur at the same rate. The concentrations of reactants and products remain constant. The reactions are still happening (dynamic) but there is no overall change.
2. The equilibrium below exists in a closed system: 2NO₂(g) ⇌ N₂O₄(g) (forward reaction exothermic). Predict the effect of increasing the temperature on the position of equilibrium and the yield of N₂O₄. (3 marks)
Increasing temperature shifts equilibrium in the endothermic direction (to the left) to absorb the extra heat. This means the yield of N₂O₄ decreases and more NO₂ is formed.
3. In the Haber process, explain why a temperature of 450 °C is used rather than a lower temperature. (3 marks)
A lower temperature would give a higher yield of ammonia because the forward reaction is exothermic. However, at low temperatures the rate of reaction is too slow. A compromise temperature of 450 °C gives a reasonable yield at an acceptable rate.
4. For the equilibrium H₂(g) + I₂(g) ⇌ 2HI(g), explain why changing the pressure has no effect on the position of equilibrium. (2 marks)
There are 2 moles of gas on the left and 2 moles of gas on the right. Since both sides have the same number of moles, changing pressure has no effect on the position of equilibrium.
5. The equilibrium: CoCl₄²⁻(aq) + 6H₂O(l) ⇌ [Co(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) is endothermic in the forward direction. A blue solution turns pink when cooled. Explain this observation using Le Chatelier's principle. (3 marks)
Decreasing the temperature causes the equilibrium to shift in the exothermic direction (to the right) to release heat. This produces more pink [Co(H₂O)₆]²⁺ ions, so the solution turns pink.
Maths Skills
Reading Equilibrium Graphs
Equilibrium graphs typically show the concentration of reactants and products over time. At the start of a reversible reaction, reactant concentrations decrease and product concentrations increase. When the concentrations become constant (the lines flatten), dynamic equilibrium has been reached.
Key features to identify:
- The point where the lines level off — this is when equilibrium is reached.
- The relative heights of the lines — these show the relative concentrations at equilibrium and indicate the position of equilibrium.
- If a condition changes (e.g. temperature increases), the concentrations will shift and then level off again at new values.
Example calculation: If at equilibrium, [N₂] = 2.0 mol/dm³, [H₂] = 6.0 mol/dm³ and [NH₃] = 4.0 mol/dm³, the ratio of product to reactant concentrations is 4.0 : 8.0, showing the equilibrium lies to the left (more reactants).
Common Misconceptions
Misconceptions About Equilibrium
At equilibrium, the amounts of reactants and products are equal.
At equilibrium, the rates of the forward and reverse reactions are equal, not the amounts. The concentrations of reactants and products can be very different at equilibrium — it depends on the position of equilibrium.
Le Chatelier's principle predicts what you WANT to happen.
Le Chatelier's principle predicts that the equilibrium will shift to OPPOSE the change you made. If you increase temperature, the system shifts in the endothermic direction to reduce the temperature — it tries to counteract the change, not to help it.
Equilibrium means the reaction has stopped.
At dynamic equilibrium, both the forward and reverse reactions are still happening — they are just happening at the same rate, so there is no overall change in concentrations.
6-Mark Question
Explain Le Chatelier's principle and apply it to the Haber process.
Le Chatelier's principle states that if a change is made to the conditions of a system at equilibrium, the system responds to oppose the change and restore equilibrium. In the Haber process, N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (exothermic forward direction). Increasing pressure shifts equilibrium towards the side with fewer gas molecules — the right (4 moles → 2 moles), so the yield of ammonia increases. Increasing temperature shifts equilibrium in the endothermic direction — the left — so the yield of ammonia decreases. However, a lower temperature gives a very slow rate, so a compromise temperature of 450 °C is used to give a reasonable yield at a reasonable rate. A high pressure of 200 atm increases yield but is expensive and has safety risks, so a compromise is used. The unreacted N₂ and H₂ are recycled to improve the overall yield.
AO3: Analysis and Evaluation
Evaluating Haber Process Conditions
A company wants to maximise ammonia production from the Haber process: N₂ + 3H₂ ⇌ 2NH₃ (exothermic forward direction).
Evaluate why they cannot simply use the highest possible pressure and lowest possible temperature to maximise yield, and explain why compromise conditions of 450 °C and 200 atm are used instead.
Answer: Theoretically, low temperature would maximise yield (equilibrium shifts right) but the rate would be unacceptably slow. High pressure would maximise yield (fewer moles on product side) but very high pressures require expensive equipment, thick-walled vessels, and more energy for compression, increasing costs and safety risks. At 200 atm, a reasonable yield is obtained without excessive costs. At 450 °C, the rate is fast enough to make the process economically viable even though the equilibrium yield at this temperature is lower than at lower temperatures. The use of an iron catalyst further increases the rate without affecting the equilibrium position, and recycling unreacted gases ensures high overall conversion.
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