GCSE — chemistryHigher7 min read

Reversible Reactions and Equilibrium

Understand reversible reactions, dynamic equilibrium, and Le Chatelier's principle at GCSE level.

Tutor AI Team2026-02-25T12:36:25.123Z

# Reversible Reactions and Equilibrium

Not all chemical reactions go to completion. Some reactions can go both ways — products can react to re-form the reactants. These are called reversible reactions. When the rates of the forward and backward reactions become equal, the system reaches a state called dynamic equilibrium. Understanding this concept is important for Higher-tier GCSE Chemistry.

1. Reversible Reactions

A reversible reaction is one that can proceed in both the forward and backward directions.

We use the symbol $\rightleftharpoons$ instead of $\rightarrow$ :

$\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$

Key Example: Ammonium Chloride

$\text{NH}_4\text{Cl}(s) \rightleftharpoons \text{NH}_3(g) + \text{HCl}(g)$

When heated, solid ammonium chloride decomposes into ammonia and hydrogen chloride gases (endothermic forward reaction). When cooled, these gases recombine to form ammonium chloride (exothermic reverse reaction).

Energy in Reversible Reactions

If the forward reaction is exothermic (releases energy), then the reverse reaction is endothermic (absorbs energy)
The energy change is the same magnitude but opposite sign

For example:

Forward: exothermic, releases 100 kJ
Reverse: endothermic, absorbs 100 kJ

Other Examples

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O} \rightleftharpoons \text{CuSO}_4 + 5\text{H}_2\text{O}$

Forward (heating): blue hydrated copper sulfate → white anhydrous copper sulfate + water (endothermic)
Reverse (adding water): white anhydrous copper sulfate → blue hydrated copper sulfate (exothermic)

This is used as a test for water — adding water to white anhydrous copper sulfate turns it blue.

2. Dynamic Equilibrium

In a closed system (nothing can enter or leave), a reversible reaction reaches dynamic equilibrium.

At dynamic equilibrium:

The rate of the forward reaction equals the rate of the backward reaction
The concentrations of reactants and products remain constant (but are NOT necessarily equal)
Both reactions are still happening — it's "dynamic" (moving), not static

Important: Equilibrium does not mean equal amounts of reactants and products. It means the rates are equal, so the amounts stay constant.

Conditions for Equilibrium

The reaction must be reversible
The system must be closed (no substances enter or leave)
The forward and backward rates must be equal

3. Le Chatelier's Principle

Le Chatelier's Principle states:

If a system at equilibrium is subjected to a change, the system will shift to oppose the change and partially restore the original conditions.

The system acts like it's trying to "undo" whatever you did to it.

3.1 Effect of Temperature

Consider: $\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$ (forward is exothermic)

Change	System shifts	Effect
Increase temperature	Shifts towards endothermic direction (backward)	More reactants, less products
Decrease temperature	Shifts towards exothermic direction (forward)	More products, less reactants

3.2 Effect of Concentration

Change	System shifts	Effect
Increase concentration of a reactant	Shifts forward (to use up the extra reactant)	More products formed
Decrease concentration of a reactant	Shifts backward (to replace the reactant)	Less products
Increase concentration of a product	Shifts backward (to use up the extra product)	More reactants

3.3 Effect of Pressure (for gaseous reactions)

Change	System shifts	Effect
Increase pressure	Shifts towards the side with fewer gas moles	Reduces the number of gas molecules
Decrease pressure	Shifts towards the side with more gas moles	Increases gas molecules

Example: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$

Left side: 1 + 3 = 4 moles of gas Right side: 2 moles of gas

Increasing pressure → shifts right (towards fewer moles) → more NH₃ produced.

3.4 Effect of Catalysts

A catalyst does NOT change the position of equilibrium.

It speeds up both the forward AND backward reactions equally
Equilibrium is reached faster, but the proportions of reactants and products are unchanged

4. The Haber Process

The Haber process is the industrial production of ammonia:

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

The forward reaction is exothermic ( $\Delta H = -92$ kJ/mol).

Conditions Used

Condition	Value	Reason
Temperature	450°C	Compromise: higher yield at low temp, but too slow
Pressure	200 atm	Higher pressure = more NH₃ (fewer gas moles on right)
Catalyst	Iron	Speeds up reaction; doesn't change yield

Why 450°C and Not Lower?

By Le Chatelier's principle: Since the forward reaction is exothermic, a lower temperature would give a higher yield of ammonia. However, at low temperatures, the rate would be too slow to be economically viable. 450°C is a compromise — reasonable rate AND reasonable yield.

Why Not Higher Pressure?

Higher pressure would give more NH₃, but very high pressures are dangerous and expensive (stronger equipment needed). 200 atm is a compromise.

Worked Example: Applying Le Chatelier's Principle

Problem

Question: The reaction $\text{2SO}_2 + \text{O}_2 \rightleftharpoons \text{2SO}_3$ is exothermic. What happens to the yield of SO₃ if the temperature is increased?

Solution

The forward reaction is exothermic. Increasing temperature shifts equilibrium towards the endothermic direction (backward). So the yield of SO₃ decreases.

Worked Example: Pressure Changes

Problem

Question: For the reaction $\text{N}_2\text{O}_4 \rightleftharpoons 2\text{NO}_2$ , predict the effect of increasing pressure.

Solution

Left: 1 mole of gas; Right: 2 moles of gas
Increasing pressure shifts towards fewer moles → shifts left
More N₂O₄ and less NO₂

Worked Example: Haber Process

Problem

Question: Explain why an iron catalyst is used in the Haber process.

Solution

The iron catalyst increases the rate of reaction by providing an alternative pathway with lower activation energy, so equilibrium is reached faster. This makes the process more economical. However, the catalyst does not change the yield of ammonia — it only speeds up the process.

6. Practice Questions

1. What is meant by a reversible reaction?
1. Define dynamic equilibrium. Why must the system be closed?
1. State Le Chatelier's Principle.
1. For: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ $N_{2} + 3 H_{2} ⇌ 2 NH_{3}$ (exothermic forward):
  - (a) What happens to yield if temperature increases?
  - (b) What happens to yield if pressure increases?
  - (c) Why is 450°C used as a compromise?
1. Explain why a catalyst doesn't affect the position of equilibrium.

Want to check your answers and get step-by-step solutions?

7. Common Misconceptions

Misconception	Reality
At equilibrium, nothing is happening	Both reactions are still occurring — rates are equal (dynamic)
Equilibrium means 50:50 mix	The ratio depends on conditions; it's rarely exactly equal
Catalysts change the equilibrium position	Catalysts speed up both directions equally — no change in position
Higher temperature always gives more product	Only if the forward reaction is endothermic

8. Exam Tips

Use the phrase "the equilibrium shifts to the right/left" clearly
When discussing Haber process conditions, always mention compromise
Remember: catalysts don't change yield — they change the rate of reaching equilibrium
For pressure questions, count the moles of gas on each side
In Le Chatelier's principle questions, state what happens, then explain why

Summary

Reversible reactions can go forward and backward ( $\rightleftharpoons$ )
Dynamic equilibrium: rates of forward and backward reactions are equal; concentrations constant
Le Chatelier's principle: the system opposes changes
Temperature: shifts towards endothermic side when heated
Pressure: shifts towards side with fewer gas moles when increased
Concentration: shifts to use up added substance
Catalysts: no effect on equilibrium position — only speed up reaching equilibrium
Haber process: 450°C, 200 atm, iron catalyst — compromise conditions

Tags:

#gcse#chemistry#rate-of-reaction#equilibrium#reversible

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# Reversible Reactions and Equilibrium

1. Reversible Reactions

Key Example: Ammonium Chloride

Energy in Reversible Reactions

Other Examples

2. Dynamic Equilibrium

Conditions for Equilibrium

3. Le Chatelier's Principle

3.1 Effect of Temperature

3.2 Effect of Concentration

3.3 Effect of Pressure (for gaseous reactions)

3.4 Effect of Catalysts

4. The Haber Process

Conditions Used

Why 450°C and Not Lower?

Why Not Higher Pressure?

Worked Example: Applying Le Chatelier's Principle

Worked Example: Pressure Changes

Worked Example: Haber Process

6. Practice Questions

7. Common Misconceptions

8. Exam Tips

Summary

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