What you'll learn
This revision guide covers reversible reactions and dynamic equilibrium as required for AQA GCSE Chemistry (specification 4.6.2). You'll learn how to identify reversible reactions, understand what happens at equilibrium, and predict how changing conditions affects the position of equilibrium using Le Chatelier's principle. These concepts are essential for understanding industrial processes like the Haber process and hydration of ethene.
Key terms and definitions
Reversible reaction — a chemical reaction where the products can react to form the original reactants, indicated by the ⇌ symbol
Dynamic equilibrium — the state reached in a closed system when the forward and reverse reactions occur at equal rates, so the concentrations of reactants and products remain constant
Closed system — a system where no substances can enter or leave, though energy can be transferred to or from the surroundings
Le Chatelier's principle — the idea that if conditions change in a system at equilibrium, the position of equilibrium shifts to counteract that change
Position of equilibrium — the relative amounts of reactants and products present at equilibrium; can lie to the left (favouring reactants) or right (favouring products)
Forward reaction — the reaction proceeding from left to right in the equation (reactants → products)
Reverse reaction — the reaction proceeding from right to left in the equation (products → reactants)
Exothermic reaction — a reaction that transfers energy to the surroundings, usually as heat (temperature increases)
Core concepts
Understanding reversible reactions
In most reactions studied at GCSE, the reaction goes to completion and all reactants are converted to products. However, some reactions are reversible — the products can react together to re-form the original reactants.
The symbol ⇌ indicates a reversible reaction. For example:
Ammonium chloride decomposition: NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
When solid ammonium chloride is heated, it decomposes into ammonia and hydrogen chloride gases (forward reaction). When these gases cool, they recombine to form solid ammonium chloride (reverse reaction).
Hydration of copper sulfate: CuSO₄(s) + 5H₂O(l) ⇌ CuSO₄·5H₂O(s) (white) (blue)
Anhydrous copper sulfate (white) reacts with water to form hydrated copper sulfate (blue). The blue crystals can be heated to remove water and reform the white powder.
Energy changes in reversible reactions:
- If the forward reaction is exothermic, the reverse reaction is endothermic (and vice versa)
- The energy change is the same magnitude but opposite direction
- Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
- Forward reaction (exothermic): ΔH = -92 kJ/mol
- Reverse reaction (endothermic): ΔH = +92 kJ/mol
Dynamic equilibrium in closed systems
For equilibrium to be established, the reversible reaction must occur in a closed system where no reactants or products can escape.
How equilibrium is reached:
- Initially, only reactants are present, so only the forward reaction occurs
- As products form, the reverse reaction begins
- The forward reaction rate decreases (less reactants available)
- The reverse reaction rate increases (more products available)
- Eventually, both reactions occur at exactly the same rate
- The concentrations of all substances remain constant (but not necessarily equal)
Key characteristics of dynamic equilibrium:
- Both forward and reverse reactions continue happening (it's dynamic, not static)
- The rates of forward and reverse reactions are equal
- Macroscopic properties (concentration, colour, pressure) remain constant
- The equilibrium can be approached from either direction
- The same equilibrium position is reached regardless of starting quantities (in a given set of conditions)
Important: Equal rates does NOT mean equal concentrations. The amounts of reactants and products at equilibrium depend on the specific reaction.
Le Chatelier's principle
When the conditions of a system at equilibrium change, the system responds by shifting the position of equilibrium to counteract that change. This is Le Chatelier's principle.
The position of equilibrium can shift:
- To the right (favouring products) — more products are formed
- To the left (favouring reactants) — more reactants are formed
This principle allows us to predict how changing temperature, pressure, or concentration affects equilibrium.
Effect of temperature on equilibrium
Temperature changes affect the position of equilibrium. The direction of shift depends on whether the forward reaction is exothermic or endothermic.
For an exothermic forward reaction:
- Increasing temperature shifts equilibrium to the LEFT (favouring reverse endothermic reaction)
- Decreasing temperature shifts equilibrium to the RIGHT (favouring forward exothermic reaction)
For an endothermic forward reaction:
- Increasing temperature shifts equilibrium to the RIGHT (favouring forward endothermic reaction)
- Decreasing temperature shifts equilibrium to the LEFT (favouring reverse exothermic reaction)
Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (forward reaction is exothermic)
- At higher temperature: equilibrium shifts left, less ammonia produced
- At lower temperature: equilibrium shifts right, more ammonia produced
Important: Temperature is the only factor that changes the value of the equilibrium constant (not required for GCSE calculation, but the concept matters).
Effect of pressure on equilibrium (gases only)
Pressure changes only affect equilibrium when gases are involved and there are different numbers of gas molecules on each side of the equation.
Count the gas molecules by looking at the large numbers (coefficients) in the balanced equation.
Rules:
- Increasing pressure shifts equilibrium towards the side with fewer gas molecules
- Decreasing pressure shifts equilibrium towards the side with more gas molecules
Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
- Left side: 1 + 3 = 4 gas molecules
- Right side: 2 gas molecules
- Increasing pressure shifts equilibrium RIGHT (towards fewer molecules)
- More ammonia is produced
When pressure has no effect:
- Equal numbers of gas molecules on both sides
- Example: H₂(g) + I₂(g) ⇌ 2HI(g) (2 molecules on each side)
- No solids or liquids involved (pressure doesn't affect their volumes significantly)
Effect of concentration on equilibrium
Changing the concentration of a reactant or product causes the equilibrium to shift to counteract that change.
Rules:
- Increasing the concentration of a substance shifts equilibrium away from that substance
- Decreasing the concentration of a substance shifts equilibrium towards that substance
Example: Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) (pale yellow) (colourless) (blood red)
- Adding more Fe³⁺ ions: equilibrium shifts RIGHT, solution becomes darker red
- Removing Fe³⁺ ions: equilibrium shifts LEFT, solution becomes paler
- Adding more FeSCN²⁺: equilibrium shifts LEFT, using up the added product
Catalyst effects: Catalysts are NOT mentioned in Le Chatelier's principle because they:
- Speed up both forward and reverse reactions equally
- Help equilibrium be reached faster
- Do NOT change the position of equilibrium
- Do NOT change the amounts of products or reactants at equilibrium
Industrial applications: The Haber process
The Haber process manufactures ammonia from nitrogen and hydrogen:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol (exothermic)
Raw materials:
- Nitrogen from air (fractional distillation of liquefied air)
- Hydrogen from natural gas (methane) or crude oil
Conditions used:
- Temperature: 450°C (moderate temperature)
- Pressure: 200 atmospheres (high pressure)
- Catalyst: Iron
Why these conditions?
Temperature compromise:
- Lower temperatures favour products (exothermic forward reaction) — better yield
- However, lower temperatures mean slower reaction rate
- 450°C is a compromise between reasonable yield and economically viable rate
High pressure:
- Shifts equilibrium right (4 gas molecules → 2 gas molecules)
- Increases yield of ammonia
- Higher pressures would be better but:
- Equipment costs increase significantly
- Safety concerns
- Energy costs for compression
- 200 atmospheres is an economic compromise
Iron catalyst:
- Speeds up rate at which equilibrium is reached
- Makes the process economically viable at moderate temperature
- Doesn't affect equilibrium position
Process details:
- Unreacted N₂ and H₂ are recycled
- Ammonia is liquefied and removed (cooling)
- Typical yield per pass: 15-20% (but recycling improves overall efficiency)
Uses of ammonia:
- Manufacture of fertilisers (ammonium nitrate, ammonium sulfate)
- Production of nitric acid
- Manufacture of explosives
- Cleaning products
Industrial applications: Hydration of ethene
Ethene reacts reversibly with steam to produce ethanol:
C₂H₄(g) + H₂O(g) ⇌ C₂H₅OH(g) ΔH = -46 kJ/mol (exothermic)
Conditions used:
- Temperature: 300°C
- Pressure: 60-70 atmospheres
- Catalyst: Phosphoric acid (H₃PO₄) on a solid support
Applying Le Chatelier's principle:
Temperature:
- Forward reaction is exothermic
- Lower temperature would increase yield
- 300°C used as compromise for reasonable rate
Pressure:
- 2 gas molecules on left → 1 gas molecule on right
- High pressure shifts equilibrium right
- Increases ethanol yield
- 60-70 atmospheres is economic compromise
Unreacted ethene is recycled to improve overall conversion efficiency.
Worked examples
Example 1: Predicting equilibrium shifts (3 marks)
Question: Sulfur dioxide reacts with oxygen to form sulfur trioxide in a reversible reaction:
2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -197 kJ/mol
Explain the effect of increasing the temperature on the yield of sulfur trioxide.
Answer:
- The forward reaction is exothermic ✓
- Increasing temperature shifts the equilibrium to the left / favours the reverse (endothermic) reaction ✓
- This decreases the yield of sulfur trioxide ✓
Mark scheme notes: Must identify the forward reaction as exothermic, state the direction of shift correctly, and link to the effect on yield.
Example 2: Applying Le Chatelier's principle to the Haber process (4 marks)
Question: The Haber process produces ammonia: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
The forward reaction is exothermic. Explain why a temperature of 450°C is used rather than a lower temperature.
Answer:
- A lower temperature would increase the yield of ammonia / shift equilibrium to the right ✓
- However, a lower temperature would make the rate of reaction too slow ✓
- The process would not be economically viable at lower temperatures ✓
- 450°C is a compromise between yield and rate ✓
Mark scheme notes: Must explain both the equilibrium effect and the kinetic (rate) consideration, plus the economic aspect.
Example 3: Pressure effects (3 marks)
Question: Explain why increasing pressure increases the yield of ammonia in the Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Answer:
- There are 4 gas molecules on the left / reactant side ✓
- There are 2 gas molecules on the right / product side ✓
- Increasing pressure shifts equilibrium towards the side with fewer molecules / to the right, increasing ammonia yield ✓
Mark scheme notes: Must count molecules correctly and link pressure to the side with fewer gas molecules.
Common mistakes and how to avoid them
Confusing "equilibrium" with "equal amounts": Dynamic equilibrium means equal rates of forward and reverse reactions, NOT equal concentrations. The amounts at equilibrium can be very different.
Forgetting to count gas molecules correctly for pressure effects: Count the large numbers (coefficients) in the equation. In N₂ + 3H₂ ⇌ 2NH₃, there are 4 molecules on the left (1 + 3), not 2.
Thinking catalysts shift equilibrium position: Catalysts speed up both forward and reverse reactions equally. They help reach equilibrium faster but don't change the position or the final amounts present.
Applying pressure effects to non-gas equilibria: Pressure changes only significantly affect gases. Changes in pressure have negligible effects on equilibria involving only solids and liquids.
Reversing temperature effects: For exothermic forward reactions, increasing temperature shifts equilibrium LEFT (not right). The system counteracts the temperature increase by favouring the endothermic (reverse) reaction.
Not explaining compromises in industrial processes: In exam questions about the Haber process, you must explain that conditions are a compromise between yield (equilibrium position) and rate, considering economic factors.
Exam technique for "The rate and extent of chemical change: reversible reactions and dynamic equilibrium"
"Explain" questions (3-4 marks): State which direction equilibrium shifts, explain why using Le Chatelier's principle, and describe the effect on yield/product. Reference the specific reaction type (exothermic/endothermic) or number of molecules.
Command word "predict": State what will happen to the position of equilibrium or yield. You may need to give a brief reason (1-2 marks typically).
Industrial process questions: Always mention compromise conditions balancing yield, rate, and cost. State that unreacted materials are recycled. The Haber process appears frequently—learn the conditions and justifications thoroughly.
Symbol recognition: ⇌ indicates reversible; → indicates one-way. Using the wrong symbol loses marks in equations.
Quick revision summary
Reversible reactions can proceed in both directions, shown by ⇌. In a closed system, dynamic equilibrium occurs when forward and reverse reaction rates are equal, keeping concentrations constant. Le Chatelier's principle states that changing conditions shifts equilibrium to counteract the change. Increasing temperature favours the endothermic direction; increasing pressure favours the side with fewer gas molecules. The Haber process uses 450°C, 200 atmospheres, and an iron catalyst as a compromise between yield and rate. Catalysts speed up reaching equilibrium but don't change the equilibrium position.