What you'll learn
This topic examines how reversible reactions respond to external changes. You must understand how to predict and explain shifts in equilibrium position when conditions change—a concept tested consistently across Paper 2 structured questions and Multiple Choice items. CXC CSEC Chemistry examiners expect you to apply Le Chatelier's Principle to industrial processes like the Haber process and Contact process.
Key terms and definitions
Reversible reaction — a chemical reaction that can proceed in both forward and backward directions, represented by the symbol ⇌.
Dynamic equilibrium — the state in a closed system where the rate of the forward reaction equals the rate of the reverse reaction, and concentrations of reactants and products remain constant.
Le Chatelier's Principle — when a system at equilibrium is subjected to an external stress (change in concentration, temperature, or pressure), the system adjusts to counteract that stress and establish a new equilibrium.
Position of equilibrium — the relative amounts of reactants and products present at equilibrium; shifts right (toward products) or left (toward reactants) in response to changes.
Closed system — a system where matter cannot enter or leave, though energy can be exchanged with surroundings.
Exothermic reaction — a reaction that releases energy to the surroundings (ΔH is negative).
Endothermic reaction — a reaction that absorbs energy from the surroundings (ΔH is positive).
Catalyst — a substance that increases the rate of both forward and reverse reactions equally without being consumed, allowing equilibrium to be reached faster without changing the equilibrium position.
Core concepts
Establishing dynamic equilibrium
A reversible reaction reaches dynamic equilibrium only in a closed system. Consider the formation of nitrogen dioxide:
2NO₂(g) ⇌ N₂O₄(g)
Initially, brown NO₂ gas converts to colourless N₂O₄. As N₂O₄ builds up, it starts decomposing back to NO₂. Eventually, both reactions occur at equal rates. The brown colour stabilizes because concentrations stop changing, even though both reactions continue.
Key characteristics of dynamic equilibrium:
- Forward and reverse reactions continue occurring
- Rates of forward and reverse reactions are equal
- Concentrations of all species remain constant
- Observable properties (colour, pressure, pH) remain constant
- Only occurs in closed systems
Effect of concentration changes
When you add or remove a reactant or product, the system shifts to counteract the change.
Adding a reactant: The equilibrium shifts right (toward products) to consume the added substance.
Example: For the iron(III) thiocyanate equilibrium used in Caribbean water quality testing:
Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) (pale yellow) (colourless) (blood red)
Adding more Fe³⁺ ions shifts equilibrium right, deepening the red colour as more FeSCN²⁺ forms.
Removing a product: The equilibrium shifts right to produce more of the removed substance.
Adding a product or removing a reactant: The equilibrium shifts left (toward reactants).
In exam questions, you must:
- Identify which substance changed
- State the direction of shift
- Explain the shift using Le Chatelier's Principle
- Describe observable changes (colour, amount of solid)
Effect of temperature changes
Temperature changes affect the equilibrium position differently depending on whether the reaction is exothermic or endothermic.
For exothermic reactions (ΔH negative):
Consider the Haber process for ammonia production, critical for fertilizer manufacturing in Caribbean agriculture:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol
Heat is a product in the forward direction. Increasing temperature shifts equilibrium left (toward reactants) because the system counteracts the added heat by favouring the endothermic reverse reaction. Less ammonia forms at higher temperatures.
Decreasing temperature shifts equilibrium right (toward products), producing more ammonia.
For endothermic reactions (ΔH positive):
Thermal decomposition of calcium carbonate (limestone, abundant in Jamaica's Cockpit Country):
CaCO₃(s) ⇌ CaO(s) + CO₂(g) ΔH = +178 kJ/mol
Heat is a reactant in the forward direction. Increasing temperature shifts equilibrium right, producing more calcium oxide and carbon dioxide.
Temperature summary:
- Increasing temperature favours the endothermic direction
- Decreasing temperature favours the exothermic direction
- Temperature is the only factor that changes the equilibrium constant (K)
Effect of pressure changes (for gaseous equilibria)
Pressure changes only affect equilibria involving gases. The system responds by shifting toward the side with fewer gas molecules.
Ammonia synthesis:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Left side: 4 gas molecules (1 + 3) Right side: 2 gas molecules
Increasing pressure shifts equilibrium right because the system reduces pressure by favouring the side with fewer molecules. More ammonia forms.
Decreasing pressure shifts equilibrium left, producing less ammonia.
When pressure has no effect:
If both sides have equal numbers of gas molecules, pressure changes don't shift equilibrium.
H₂(g) + I₂(g) ⇌ 2HI(g)
Left side: 2 molecules Right side: 2 molecules
Changing pressure has no effect on this equilibrium position.
Important notes:
- Only total gas molecules matter, not stoichiometric coefficients alone
- Solids and liquids are incompressible and don't contribute to pressure considerations
- Pressure changes affect position of equilibrium but not the equilibrium constant
Effect of catalysts
Catalysts speed up both forward and reverse reactions equally by lowering activation energy for both directions.
Effects of adding a catalyst:
- Equilibrium is reached faster
- Same equilibrium position
- Same equilibrium concentrations
- No shift in equilibrium
In the Contact process for sulfuric acid production (manufactured at petroleum refineries in Trinidad):
2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
Vanadium(V) oxide (V₂O₅) catalyst allows equilibrium to be reached quickly at moderate temperatures, making the process economically viable. The catalyst doesn't change how much SO₃ forms at equilibrium, only how quickly that equilibrium is achieved.
Industrial applications
Haber Process (Ammonia production):
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol
Optimal conditions balance yield and rate:
- Temperature: 450°C (compromise—lower would give better yield but too slow)
- Pressure: 200 atm (high pressure increases yield)
- Catalyst: Iron catalyst (speeds reaction)
- Unreacted N₂ and H₂ are recycled
Contact Process (Sulfuric acid production):
2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -197 kJ/mol
Industrial conditions:
- Temperature: 450°C (compromise temperature)
- Pressure: 1-2 atm (sufficient, as forward reaction is highly favoured)
- Catalyst: Vanadium(V) oxide
Compromise conditions: Industrial processes rarely use the conditions that give maximum yield because:
- Very high pressures are expensive and dangerous
- Very low temperatures make reactions too slow
- Economic considerations require balancing yield, rate, and cost
Worked examples
Example 1: Concentration change in an aqueous equilibrium
The chromate-dichromate equilibrium is established in solution:
2CrO₄²⁻(aq) + 2H⁺(aq) ⇌ Cr₂O₇²⁻(aq) + H₂O(l) (yellow) (orange)
(a) State and explain the colour change observed when dilute sulfuric acid is added to a yellow solution containing chromate ions. [3 marks]
(b) Describe how sodium hydroxide solution could be used to reverse this colour change. [2 marks]
Solution:
(a)
- The solution turns orange. [1 mark]
- Adding H₂SO₄ increases [H⁺]. [1 mark]
- By Le Chatelier's Principle, equilibrium shifts right to consume the added H⁺, forming more orange Cr₂O₇²⁻. [1 mark]
(b)
- NaOH neutralizes H⁺ ions: OH⁻ + H⁺ → H₂O [1 mark]
- Removing H⁺ shifts equilibrium left, reforming yellow CrO₄²⁻. [1 mark]
Example 2: Temperature effects on equilibrium yield
The reaction for the production of methanol is:
CO(g) + 2H₂(g) ⇌ CH₃OH(g) ΔH = -91 kJ/mol
(a) State whether this reaction is exothermic or endothermic in the forward direction. [1 mark]
(b) Explain, using Le Chatelier's Principle, what happens to the yield of methanol when the temperature is increased. [3 marks]
(c) Suggest why industrial plants still operate this reaction at moderately high temperatures despite your answer to part (b). [2 marks]
Solution:
(a) Exothermic (because ΔH is negative). [1 mark]
(b)
- Increasing temperature favours the endothermic direction. [1 mark]
- For this exothermic reaction, the reverse (endothermic) reaction is favoured. [1 mark]
- Equilibrium shifts left, decreasing the yield of methanol. [1 mark]
(c)
- At low temperatures, the reaction rate is too slow [1 mark]
- Higher temperatures provide a compromise between acceptable rate and reasonable yield. [1 mark]
Example 3: Pressure effects on gaseous equilibrium
Consider the equilibrium:
PCl₅(g) ⇌ PCl₃(g) + Cl₂(g)
(a) State the total number of gas molecules on each side of the equation. [1 mark]
(b) Predict and explain the effect of increasing pressure on the amount of PCl₅ present at equilibrium. [3 marks]
Solution:
(a) Left side: 1 molecule; Right side: 2 molecules [1 mark]
(b)
- Increasing pressure shifts equilibrium toward the side with fewer gas molecules. [1 mark]
- Equilibrium shifts left (toward reactants). [1 mark]
- The amount of PCl₅ increases. [1 mark]
Common mistakes and how to avoid them
Mistake: Stating that catalysts shift the equilibrium position or increase yield. Correction: Catalysts only affect the rate at which equilibrium is reached. They don't change equilibrium concentrations, yield, or position. Always state: "Catalyst speeds up forward and reverse reactions equally without affecting equilibrium position."
Mistake: Confusing "shift right" with "reaction speeds up" or "equilibrium constant increases." Correction: Shifting right means more products form and fewer reactants remain at the new equilibrium. The equilibrium constant (K) only changes with temperature. Use precise language: "equilibrium shifts right" not "reaction goes right."
Mistake: Ignoring the state symbols when applying pressure changes—including solids and liquids in molecule counts. Correction: Only count gaseous species when predicting pressure effects. For CaCO₃(s) ⇌ CaO(s) + CO₂(g), only one gas molecule exists, so you cannot meaningfully compare sides. Pressure changes have minimal effect when solids are involved.
Mistake: Thinking that adding more catalyst shifts equilibrium to produce more product. Correction: Adding catalyst or increasing catalyst amount speeds the approach to equilibrium but doesn't change final concentrations. Both forward and reverse rates increase proportionally.
Mistake: Incorrectly identifying the endothermic/exothermic direction from ΔH sign. Correction: Remember: negative ΔH means forward reaction is exothermic (releases heat), positive ΔH means forward reaction is endothermic (absorbs heat). Heat can be treated as a product (exothermic) or reactant (endothermic) when applying Le Chatelier's Principle.
Mistake: Writing "equilibrium moves" instead of "equilibrium shifts." Correction: Use standard terminology. Write "equilibrium shifts right/left" or "position of equilibrium shifts toward products/reactants." Examiners expect precise terminology for full marks.
Exam technique for Le Chatelier's Principle and Factors Affecting Equilibrium
Command word "Explain": You must state what happens AND use Le Chatelier's Principle to justify why. Example: "Adding NH₃ shifts equilibrium left [what] because the system counteracts the increase in product concentration [why]." This typically earns 2-3 marks.
Command word "Predict": State the direction of shift and the observable result. "Equilibrium shifts right; more SO₃ forms" is complete. You may not need full explanation unless marks allocated suggest otherwise. Check mark allocation.
Observable changes: CXC examiners frequently ask about colour changes, gas volume changes, or precipitate formation. Link equilibrium shifts to physical observations: "More FeSCN²⁺ forms, deepening the red colour."
Industrial processes: When asked about compromise conditions, address multiple factors. Template: "High pressure increases yield [Le Chatelier] but is expensive and dangerous [economic/safety]. Moderate pressure of X atm provides acceptable yield at lower cost." This structured approach secures full marks.
Quick revision summary
Dynamic equilibrium occurs when forward and reverse reaction rates equal in a closed system. Le Chatelier's Principle states that a system at equilibrium counteracts any imposed stress. Increasing concentration of reactants shifts equilibrium right; increasing products shifts left. Raising temperature favours the endothermic direction; lowering temperature favours exothermic. Increasing pressure shifts equilibrium toward fewer gas molecules. Catalysts speed both reactions equally without changing equilibrium position. Industrial processes use compromise conditions balancing yield, rate, and cost. Always justify predictions using Le Chatelier's Principle and link shifts to observable changes for full marks on CXC CSEC Chemistry examinations.