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Reversible Reactions

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Reversible Reactions

Summary: Reversible reactions reach dynamic equilibrium when forward and backward rates are equal in a closed system. Le Chatelier’s Principle predicts that the position of equilibrium shifts to oppose any change in temperature, pressure, or concentration. Catalysts speed up both directions equally without affecting equilibrium position. Tags: igcse chemistry reactions Created: 2026-07-14 Last Updated: 2026-07-16

Reversible Reactions

Definition: A reversible reaction is one that can proceed in both the forward and the reverse direction.

Notation: The symbol ⇌ (two half-arrows pointing in opposite directions) is used instead of a single arrow.

Example:

  • Forward reaction: A + B → C + D
  • Reverse reaction: C + D → A + B
  • Reversible: A + B ⇌ C + D

Key idea: The forward reaction converts reactants to products; the reverse reaction converts products back to reactants. Both reactions occur simultaneously.

Conditions for reversibility: For a reversible reaction to be observed, the system must usually be closed — no substances can enter or leave. If products escape (e.g., a gas produced in an open container), the reverse reaction cannot occur.

Dynamic Equilibrium

Definition: Dynamic equilibrium is the state reached in a reversible reaction in a closed system when:

  1. The rate of the forward reaction = the rate of the backward reaction.
  2. The concentrations of reactants and products remain constant (but NOT necessarily equal).

Why it is called “dynamic”: Both the forward and reverse reactions continue to happen — they do NOT stop. The equilibrium is dynamic (active, ongoing) rather than static. Reactants are continuously being converted to products, and products are continuously being converted back to reactants, at exactly the same rate. There is no net change.

Requirements for dynamic equilibrium:

  • The system must be closed (no exchange of matter with surroundings).
  • The temperature must remain constant.
  • The reaction must be reversible.

Example — the Haber process at equilibrium: N2(g) + 3H2(g) ⇌ 2NH3(g)

  • At equilibrium: rate of N2 + 3H2 → 2NH3 = rate of 2NH3 → N2 + 3H2
  • Concentrations of N2, H2, and NH3 remain constant (but not necessarily equal).
  • Both reactions are happening continuously.

Le Chatelier’s Principle

Definition: If a system at equilibrium is subjected to a change in conditions (temperature, pressure, or concentration), the position of equilibrium will shift to oppose/counteract the change.

What “shift” means: The equilibrium moves either to the right (favouring the forward reaction, producing more products) or to the left (favouring the reverse reaction, producing more reactants) until a new equilibrium is established.

Effect of Temperature Changes

When temperature changes, the equilibrium shifts to oppose the temperature change.

  • Increasing temperature: The equilibrium shifts in the endothermic direction to absorb the added heat (cooling the system down).
  • Decreasing temperature: The equilibrium shifts in the exothermic direction to release heat (warming the system up).

Whether the forward reaction is exothermic or endothermic must be known — this is usually stated in the question or can be deduced.

Example — the Haber process: N2(g) + 3H2(g) ⇌ 2NH3(g) (forward reaction is exothermic, ΔH = -92 kJ/mol)

  • Increase temperature: equilibrium shifts to the left (endothermic direction), decreasing NH3 yield. The system absorbs the added heat by favouring the reverse (endothermic) reaction.
  • Decrease temperature: equilibrium shifts to the right (exothermic direction), increasing NH3 yield. The system releases heat to oppose the cooling.

Important industrial compromise: Lower temperatures give a higher yield of NH3, but the rate of reaction is too slow. The Haber process is run at about 450 degrees C as a compromise between yield and rate (with an iron catalyst to increase the rate).

Effect of Pressure Changes (Gaseous Systems Only)

Pressure changes only affect equilibria involving gases where the number of moles of gas differs between the two sides.

  • Increasing pressure: The equilibrium shifts towards the side with FEWER gas molecules (fewer moles of gas), reducing the pressure.
  • Decreasing pressure: The equilibrium shifts towards the side with MORE gas molecules (more moles of gas), increasing the pressure.

The number of moles of GAS on each side must be counted (ignore solids and liquids — they are virtually incompressible).

Example 1 — the Haber process: N2(g) + 3H2(g) ⇌ 2NH3(g)

  • Left side: 1 + 3 = 4 moles of gas
  • Right side: 2 moles of gas
  • Increasing pressure: shifts to the right (fewer moles, 4 → 2), increasing NH3 yield.
  • Decreasing pressure: shifts to the left (more moles, 2 → 4).

Example 2 — no effect: H2(g) + I2(g) ⇌ 2HI(g)

  • Left side: 1 + 1 = 2 moles of gas
  • Right side: 2 moles of gas
  • Changing pressure has NO effect on equilibrium position (equal number of moles on both sides).

Effect of Concentration Changes

  • Adding more of a reactant: Equilibrium shifts to the right (forward reaction favoured) to use up the added reactant.
  • Adding more of a product: Equilibrium shifts to the left (reverse reaction favoured) to use up the added product.
  • Removing a product (e.g., by letting a gas escape in an open system, or continuously removing product in an industrial process): Equilibrium shifts to the right (forward reaction favoured) to replace the removed product.

Example — esterification: RCOOH + R’OH ⇌ RCOOR’ + H2O

  • Adding more alcohol (reactant) shifts equilibrium to the right, producing more ester.
  • Removing water (product) as it forms shifts equilibrium to the right, producing more ester (this is done practically using a drying agent or by distillation).

Effect of a Catalyst

A catalyst:

  • Speeds up BOTH the forward and reverse reactions equally.
  • Does NOT affect the position of equilibrium.
  • Does NOT increase the yield of products.
  • Equilibrium is reached faster, but the equilibrium mixture has the same composition as without a catalyst.

Why? A catalyst provides an alternative reaction pathway with a lower activation energy. This lowers the activation energy for BOTH the forward and reverse reactions by the same amount, so both rates increase equally. The equilibrium composition depends on thermodynamics (ΔG), not kinetics.

Key Reversible Reactions (IGCSE Syllabus)

Hydrated Copper(II) Sulfate ⇌ Anhydrous Copper(II) Sulfate

CuSO4.5H2O(s) ⇌ CuSO4(s) + 5H2O(l) (blue hydrated crystals) ⇌ (white anhydrous powder) + water

  • Heat the blue hydrated crystals: Water of crystallisation is driven off. The solid turns white (anhydrous CuSO4). This is the forward (decomposition) reaction — endothermic.
  • Add a few drops of water to the white powder: The anhydrous solid rehydrates. The solid turns blue again and heat is released (exothermic). This is the reverse reaction.

Uses:

  • Test for water: anhydrous CuSO4 (white) turns blue in the presence of water. This is a chemical test for water (as opposed to physical tests like boiling point measurement).
  • Note: anhydrous CuSO4 detects LIQUID water — it will not detect water vapour unless condensation occurs.

Hydrated Cobalt(II) Chloride ⇌ Anhydrous Cobalt(II) Chloride

CoCl2.6H2O(s) ⇌ CoCl2(s) + 6H2O(l) (pink hydrated crystals) ⇌ (blue anhydrous solid) + water

  • Heat: Pink → Blue (water of crystallisation driven off)
  • Add water or expose to moist air: Blue → Pink (rehydration)

Use: Cobalt(II) chloride paper (blue) turns pink in the presence of water. This is used as a test for water (often in combination with anhydrous CuSO4).

Thermal Dissociation of Ammonium Chloride

NH4Cl(s) ⇌ NH3(g) + HCl(g) (white solid) ⇌ (ammonia gas) + (hydrogen chloride gas)

  • On heating: NH4Cl appears to sublime, but actually thermally dissociates into NH3 and HCl gases.
  • On cooling: The gases recombine to form white NH4Cl solid (often seen as a white deposit further up the tube).
  • This is an example of a reversible reaction involving dissociation/recombination.

The IGCSE often asks about “sublimation” of NH4Cl. Strictly, it is thermal dissociation followed by recombination — but for IGCSE purposes, describing the appearance (white solid → colourless gases → white solid again) is what earns marks.

Other reversible reaction examples for IGCSE:

  • N2(g) + 3H2(g) ⇌ 2NH3(g) (Haber process)
  • 2SO2(g) + O2(g) ⇌ 2SO3(g) (Contact process)
  • CaCO3(s) ⇌ CaO(s) + CO2(g) (in a closed system; in an open system, CO2 escapes so the reaction goes to completion)

Key Points

  • Reversible reaction: can proceed in both directions (⇌)
  • Dynamic equilibrium: rate of forward = rate of backward; concentrations constant; system must be closed
  • Le Chatelier’s Principle: if a change is imposed on a system at equilibrium, the position of equilibrium shifts to oppose the change
  • Increase temperature: favours ENDOTHERMIC direction
  • Decrease temperature: favours EXOTHERMIC direction
  • Increase pressure: favours side with FEWER moles of gas
  • Decrease pressure: favours side with MORE moles of gas
  • Catalyst: speeds up BOTH directions equally; equilibrium reached faster but position UNCHANGED; NO effect on yield
  • Concentration: add reactant → shift right; add product → shift left; remove product → shift right
  • CuSO4.5H2O (blue) ⇌ CuSO4 (white) + 5H2O — test for water
  • CoCl2.6H2O (pink) ⇌ CoCl2 (blue) + 6H2O — test for water (cobalt chloride paper)
  • NH4Cl(s) ⇌ NH3(g) + HCl(g) — thermal dissociation (appears to sublime)

Key Concepts from Past Papers

  • Reversible reaction: a reaction that can proceed in both the forward and reverse directions
  • Dynamic equilibrium: the state in a reversible reaction when the rate of the forward reaction equals the rate of the backward reaction and the concentrations of reactants and products remain constant
  • Le Chatelier’s Principle: if a system at equilibrium is subjected to a change, the position of equilibrium shifts to oppose the change
  • Copper(II) sulfate turns from blue to white on heating; water is given off
  • Adding water to anhydrous copper(II) sulfate turns it from white to blue
  • Anhydrous copper(II) sulfate can be used as a test for water
  • At equilibrium, the rate of the forward reaction equals the rate of the backward reaction
  • A catalyst does not affect the position of equilibrium / does not increase the yield
  • Increasing temperature favours the endothermic reaction
  • Increasing pressure favours the side with fewer moles of gas

Keywords from Past Papers

reaction, equations, reversible, description, including, attempts, removing, adding, waters, crystallisation, under, closed, conditions, haber, process

Sources

  • OpenStax Chemistry 2e — [Chapter 13: Fundamental Equilibrium Concepts], Rice University (free, CC BY 4.0)
  • BBC Bitesize GCSE Chemistry — [Reversible Reactions and Equilibrium], BBC (free educational resource)
  • Cambridge IGCSE Chemistry 0620 — Syllabus Section 6: Chemical Reactions (Reversible Reactions and Equilibrium), Cambridge Assessment International Education
  • CK-12 Chemistry for High School — [Chapter 19: Equilibrium], CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

  • 0620/31 May/June 2015: Q55(c)(i) (0m)

Common Misconceptions

MisconceptionReality
”At equilibrium, the amounts/concentrations of reactants and products are equal”They are CONSTANT but not necessarily EQUAL. At equilibrium, the RATES are equal, not the amounts.
”At equilibrium, the reactions have stopped”Equilibrium is DYNAMIC — both forward and reverse reactions continue at the same rate. There is no net change, but reactions are still happening.
”A catalyst increases the yield of products”A catalyst ONLY speeds up the rate at which equilibrium is reached. It has NO effect on the equilibrium position or the yield.
”Le Chatelier’s Principle says the equilibrium shifts to ‘help’ the reaction”The equilibrium shifts to OPPOSE/COUNTERACT the imposed change. It does not “help” or “assist” — it opposes.
”Pressure changes always affect equilibrium”Pressure changes only affect equilibria involving GASES where there is a DIFFERENCE in the number of moles of gas on each side.
”Increasing temperature always increases the rate of the forward reaction more”Both forward and reverse rates increase with temperature, but the endothermic reaction’s rate increases MORE (higher activation energy), shifting the equilibrium position.

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