C18: Reversible Reactions and Equilibrium
Reversible reactions, dynamic equilibrium and Le Chatelier's principle
Reversible Reactions And Equilibrium
📋 Key Definitions
Reversible reaction: A reaction in which the products can react together to re-form the original reactants. Shown by the symbol ⇌ instead of →.
Dynamic equilibrium: The state reached in a closed system where the forward and backward reactions occur at the same rate, so the concentrations of reactants and products remain constant.
Le Chatelier's principle: If a change is made to the conditions of a system at equilibrium, the position of equilibrium shifts to oppose the change.
Closed system: A system where no substances can enter or leave. Dynamic equilibrium can only be reached in a closed system.
Position of equilibrium: Describes whether the equilibrium favours the reactants (lies to the left) or the products (lies to the right).
📝 Reversible Reactions
In a reversible reaction, the products can react back together to form the reactants. This means the reaction can go in both directions.
Worked Example 1: A reversible reaction
Ammonium chloride decomposes when heated, but re-forms when cooled:
NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
Heating drives the forward reaction (decomposition — endothermic). Cooling drives the backward reaction (reformation — exothermic).
Observation: white solid disappears on heating (produces two gases) and reappears on cooling.
Energy in reversible reactions: If the forward reaction is exothermic, the backward reaction is endothermic — and vice versa. The same amount of energy is transferred in each direction.
⚖️ Dynamic Equilibrium
When a reversible reaction takes place in a closed system, eventually the forward and backward reactions happen at the same rate. This is called dynamic equilibrium.
Features of dynamic equilibrium:
- The forward and backward reactions are still happening (dynamic)
- Their rates are equal, so there is no overall change (equilibrium)
- Concentrations of reactants and products stay constant — but are not necessarily equal
- It can only be reached in a closed system
Worked Example 2: Understanding equilibrium
In the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), explain what is happening at equilibrium.
Nitrogen and hydrogen are still reacting to form ammonia (forward reaction). Ammonia is still decomposing back into nitrogen and hydrogen (backward reaction). Both reactions happen at the same rate. The concentrations of N₂, H₂ and NH₃ remain constant.
📐 Le Chatelier's Principle
Le Chatelier's principle predicts how the position of equilibrium changes when conditions are altered. The equilibrium shifts to oppose the change and restore balance.
| Change | Equilibrium shifts to... | Effect on yield |
|---|---|---|
| Temperature increased | Endothermic direction (to absorb the extra heat) | Products increase if forward is endothermic |
| Temperature decreased | Exothermic direction (to release heat) | Products increase if forward is exothermic |
| Pressure increased | Side with fewer gas moles (to reduce pressure) | Products increase if fewer moles on product side |
| Pressure decreased | Side with more gas moles (to increase pressure) | Reactants increase if more moles on product side |
| Concentration of reactant increased | Right (towards products, to use up the extra reactant) | More product formed |
| Concentration of product removed | Right (towards products, to replace the product) | More product formed |
Worked Example 3: Effect of temperature on equilibrium
The reaction N₂O₄(g) ⇌ 2NO₂(g) is endothermic in the forward direction. What happens if the temperature is increased?
According to Le Chatelier's principle, the equilibrium shifts in the endothermic direction to absorb the extra heat.
The forward reaction is endothermic, so the equilibrium shifts to the right.
More NO₂ is produced. The yield of NO₂ increases.
Worked Example 4: Effect of pressure on equilibrium
In the equilibrium 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), what happens if the pressure is increased?
Left side: 2 + 1 = 3 moles of gas. Right side: 2 moles of gas.
The equilibrium shifts to the side with fewer gas moles to reduce the pressure.
Fewer moles on the right → equilibrium shifts to the right.
More SO₃ is produced. The yield of SO₃ increases.
Worked Example 5: Effect of concentration
In the equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq), what happens if more SCN⁻ is added?
Increasing the concentration of a reactant means the equilibrium shifts to the right to use up the extra reactant.
More FeSCN²⁺ is produced. The solution becomes deeper red (FeSCN²⁺ is blood-red).
🏭 The Haber Process
The Haber process is the industrial method for making ammonia from nitrogen and hydrogen. It is an important application of Le Chatelier's principle.
Haber process:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Forward reaction is exothermic.
Moles of gas: left = 4, right = 2
Conditions: 450°C, 200 atm, iron catalyst
| Condition | What equilibrium wants | Industrial choice | Reason (compromise) |
|---|---|---|---|
| Temperature | Low temp (forward is exothermic → more NH₃) | 450°C (moderately high) | Low temp gives high yield but slow rate. 450°C is a compromise — reasonable yield with acceptable rate. |
| Pressure | High pressure (fewer moles on right → more NH₃) | 200 atm (high but not very high) | Very high pressure gives higher yield but is expensive and dangerous. 200 atm is a compromise. |
| Catalyst | Does not affect position of equilibrium | Iron catalyst | Speeds up the rate of reaction so equilibrium is reached faster. Does not change the yield. |
Important: Unreacted nitrogen and hydrogen are recycled back into the reactor. This increases the overall yield and makes the process more efficient.
❓ Practice Questions
Q1: Foundation What is a reversible reaction? Use the ⇌ symbol in your answer.
Q2: Foundation Explain what is meant by dynamic equilibrium. Why must it occur in a closed system?
Q3: Higher The reaction 2NO₂(g) ⇌ N₂O₄(g) is exothermic in the forward direction. Predict the effect of increasing the temperature on the position of equilibrium and the yield of N₂O₄.
Q4: Higher Explain why the Haber process uses a temperature of 450°C rather than a lower temperature that would give a higher yield of ammonia.
Q5: Higher In the equilibrium H₂(g) + I₂(g) ⇌ 2HI(g), what effect does increasing the pressure have on the position of equilibrium? Explain your answer.
✅ Answers
- A reversible reaction is one where the products can react to re-form the original reactants. It is shown using the ⇌ symbol. For example: NH₄Cl(s) ⇌ NH₃(g) + HCl(g).
- Dynamic equilibrium is reached when the forward and backward reactions occur at the same rate in a closed system. The concentrations of reactants and products remain constant but are not necessarily equal. It must occur in a closed system because otherwise products or reactants could escape, preventing equilibrium from being established.
- Increasing the temperature causes the equilibrium to shift in the endothermic direction (to absorb extra heat). The forward reaction is exothermic, so the backward reaction is endothermic. The equilibrium shifts to the left, producing more NO₂ and less N₂O₄. The yield of N₂O₄ decreases.
- Lower temperature would give a higher yield of ammonia (forward reaction is exothermic). However, at low temperatures the rate of reaction is very slow, so it would take too long to produce ammonia. 450°C is a compromise between yield and rate — it gives a reasonable yield at an acceptable rate.
- Increasing pressure has no effect on the position of equilibrium. There are 2 moles of gas on both sides (1 H₂ + 1 I₂ = 2 on the left, 2 HI on the right). Le Chatelier's principle only causes a shift when there is a difference in the number of gas moles between the two sides.
🎯 Exam Tips
- When applying Le Chatelier's principle, always state the direction of shift first, then the effect on yield
- For temperature questions, first identify which direction is endothermic and which is exothermic
- For pressure questions, count the moles of GAS on each side — only gas moles matter
- A catalyst does NOT change the position of equilibrium — it only helps equilibrium be reached faster
- In Haber process questions, always explain the compromise between yield and rate
- When asked about dynamic equilibrium, mention: closed system, forward = backward rate, concentrations constant
- Remember: "shift to the right" means more products; "shift to the left" means more reactants
🔢 Maths Skills
Mathematical Skills
You need to calculate the number of moles of gas on each side of an equilibrium to determine the effect of pressure changes, and interpret data from Haber process conditions.
Maths Example
In the equilibrium 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), how many moles of gas are on each side?
Left side: 2 + 1 = 3 moles. Right side: 2 moles. Increasing pressure shifts equilibrium to the right (fewer moles), increasing SO₃ yield.
Maths Example 2
In the Haber process, the yield of NH₃ at 200 atm is about 15%. If the pressure is increased to 1000 atm, the yield would be about 50%. Explain why 200 atm is used instead of 1000 atm.
Although 1000 atm gives a higher yield, very high pressures are expensive (stronger pipes and vessels needed) and dangerous (risk of explosions). 200 atm is a compromise — it gives an acceptable yield at reasonable cost and safety. The iron catalyst also helps reach equilibrium faster.
⚠️ Common Misconceptions
Watch Out!
1. Wrong: At equilibrium, the reaction has stopped Correct: At dynamic equilibrium, both forward and backward reactions are STILL happening — they just occur at the SAME RATE, so concentrations stay constant
2. Wrong: At equilibrium, the concentrations of reactants and products are equal Correct: Concentrations are CONSTANT but not necessarily EQUAL — the position of equilibrium may favour one side
3. Wrong: A catalyst changes the position of equilibrium Correct: A catalyst does NOT change the position of equilibrium — it speeds up both forward and backward reactions equally, so equilibrium is reached faster but the yield is unchanged
4. Wrong: Le Chatelier's principle means the system completely reverses the change Correct: The equilibrium shifts to PARTIALLY oppose the change — it reduces the effect but does not completely reverse it
✍️ 6-Mark Question
Extended Answer
6 marks: The Haber process uses the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g). The forward reaction is exothermic. Explain why the industrial conditions of 450°C and 200 atm are a compromise, and explain the role of the iron catalyst.
The forward reaction is exothermic, so Le Chatelier's principle predicts that a lower temperature would shift equilibrium to the right, increasing the yield of ammonia. However, at low temperatures the rate of reaction is very slow, so it would take too long to produce ammonia. 450°C is a compromise — it gives a reasonable yield at an acceptable rate. Higher pressure favours the forward reaction because there are 4 moles of gas on the left and only 2 on the right, so increasing pressure shifts equilibrium to the side with fewer moles. Very high pressure would give a higher yield but is expensive (requires stronger, thicker vessels) and dangerous (risk of explosions). 200 atm is a compromise between yield, cost and safety. The iron catalyst does not change the position of equilibrium or the yield — it speeds up both the forward and backward reactions equally, so equilibrium is reached faster. This makes the process more economically viable because ammonia is produced more quickly. Unreacted N₂ and H₂ are recycled to improve the overall yield.
Mark scheme: 1 mark for explaining low temperature gives high yield; 1 mark for explaining why 450°C is a compromise (rate vs yield); 1 mark for explaining high pressure gives high yield; 1 mark for explaining why 200 atm is a compromise (cost/safety vs yield); 1 mark for correct explanation of catalyst (does not change position, speeds up reaching equilibrium); 1 mark for mentioning recycling of unreacted gases
📊 AO3: Analyse & Evaluate
Analysis and Evaluation
A chemist investigated the equilibrium: N₂O₄(g) ⇌ 2NO₂(g). The forward reaction is endothermic. The equilibrium mixture was a pale yellow colour at 25°C and a dark brown colour at 100°C.
(a) Explain the colour change using Le Chatelier's principle.
(b) At 25°C, the equilibrium mixture contains 20% NO₂. At 100°C, it contains 60% NO₂. A student claims that "increasing temperature always increases the yield of products." Evaluate this claim.
(c) If the pressure is increased, what would happen to the colour? Explain your answer.
Answers: (a) Increasing temperature shifts equilibrium in the endothermic direction (forward) to absorb the extra heat. This produces more NO₂, which is dark brown, so the colour darkens. (b) The student's claim is incorrect. Increasing temperature only increases the yield of products if the forward reaction is endothermic. If the forward reaction is exothermic, increasing temperature would decrease the product yield by shifting equilibrium backwards. The direction of shift depends on the enthalpy change of the forward reaction, not on a general rule. (c) There is 1 mole of gas on the left and 2 moles on the right. Increasing pressure shifts equilibrium to the side with fewer moles (left, towards N₂O₄). This would reduce the amount of dark brown NO₂, so the mixture would become paler.
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