Reversible reactions

Key Vocabulary

  • Reversible reaction – a reaction where the products react back into the reactants and vice versa.

  • Dynamic equilibrium – a reversible reaction where the forward and backward reactions occur at the same time and at the same rate. Concentrations of reactants and products remain constant.

  • Position of equilibrium – describes the relative amounts of products and reactants at equilibrium. Left = mostly reactants, Right = mostly products.

  • Shift – describes how equilibrium responds to changes in conditions. Shift left = towards reactants; shift right = towards products.

  • Catalyst – speeds up a reaction without being chemically changed. Provides an alternative route with lower activation energy.

Reversible Reactions and Chemical Equilibria

  • Reversible reaction: a chemical change in which the products can be converted back to the reactants under suitable conditions.

Example: Thermal decomposition of ammonium chloride

  • When solid ammonium chloride is heated, it decomposes to form ammonia (NH₃) gas and hydrogen chloride (HCl) gas:

    • NH₄Cl(s) → NH₃(g) + HCl(g)

  • Type of reaction: decomposition

  • Further up the tube, the gases cool and react together to reform solid ammonium chloride:

    • NH₃(g) + HCl(g) → NH₄Cl(s)

  • Type of reaction: synthesis

  • Overall reaction: reversible decomposition of ammonium chloride

Thermal Decomposition of Blue Hydrated Copper(II) Sulphate Crystals

  • On heating, the blue crystals turn into a white anhydrous solid.

  • Reaction: CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(l)

  • Type of reaction: decomposition

  • The crystals lose their water of crystallisation when heated.

  • Adding water to the anhydrous copper(II) sulphate turns the white solid blue again – this is used as a test for water. Notice the temperature of the test tube.

  • The forward reaction (heating) is endothermic.

  • The reverse reaction (hydration) is exothermic.

Key point:

  • If a reversible reaction is endothermic in one direction, it is exothermic in the opposite direction. The same amount of energy is transferred each time.

Overall reaction: reversible thermal decomposition of hydrated copper(II) sulphate

Energy profile diagrams:

  • Forward reaction (endothermic) – energy absorbed from surroundings

  • Reverse reaction (exothermic) – energy released to surroundings

Dynamic Equilibrium

  • When a reversible reaction occurs in a closed system (e.g., sealed container), a dynamic equilibrium is reached.

  • Neither the forward nor reverse reactions are complete; both reactants and products are present.

  • When equilibrium is reached, the reaction appears to have stopped, but the forward and reverse reactions are occurring at the same rate – hence it is known as a dynamic system.

  • A balance has been reached.

  • At equilibrium, the concentrations of reactants and products do not change.

  • The position of equilibrium refers to the proportion of reactants/products (not usually 50:50).

Necessary conditions for dynamic equilibrium:

  • Can only be established in a closed system (no materials added or removed)

  • Forward and backward reactions occur at the same rate

  • No net overall change; concentration of reactants and products remains constant

  • Macroscopic properties (e.g., temperature, pressure) remain constant

Changes That Affect a System in Equilibrium

  • Le Chatelier’s Principle: if a dynamic equilibrium is subjected to a change, the equilibrium moves to counteract the change.

  • Factors to consider:

    • Changing the temperature

    • Changing the pressure

    • Changing the concentration of reactants or products

1. Changing Temperature

  • Example reaction: A + 2B ⇌ C + D, ΔH = –120 kJ/mol (forward reaction is exothermic)

  • Increasing the temperature (adding heat)

    • The equilibrium position shifts in the direction of the endothermic reaction (to remove heat).

    • In the example, equilibrium shifts to the left or towards the reactants.

  • Decreasing the temperature (removing heat)

    • The equilibrium position shifts in the direction of the exothermic reaction (to increase heat).

    • In the example, equilibrium shifts to the right or towards the products.

Example:

  • Dynamic equilibrium: N₂(g) + O₂(g) ⇌ 2NO(g), ΔH = positive (forward reaction is endothermic)

  • Effect of increasing temperature: equilibrium shifts towards the products (endothermic direction), so the yield of NO increases.

  1. Changing the Pressure of the Reactants or Products

  • The pressure of a gas or mixture of gases depends only on the number of molecules of gas present.

  • Example reaction: A(g) + 2B(g) ⇌ C(g) + D(g)

  • Increasing the pressure

    • The equilibrium position shifts to the side with the fewer number of gas molecules.

    • In the example, equilibrium shifts to the right or towards the products.

  • Decreasing the pressure

    • The equilibrium position shifts to the side with the greater number of gas molecules.

    • In the example, equilibrium shifts to the left or towards the reactants.

  • Example: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

  • Effect of increasing pressure: equilibrium shifts towards the products (fewer gas molecules), so the yield of SO₃ increases.

  1. Changing the Concentration of the Reactants or Products

  • Example reaction: A + 2B ⇌ C + D

  • If you add more A, the system removes it by producing more C and D, i.e., the equilibrium position shifts to the right or towards the products.

  • If you remove C as it forms, it will be replaced by more A and B reacting, i.e., the equilibrium position shifts to the left or towards the reactants.

  • Example: [Cu(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) ⇌ [CuCl₄]²⁻(aq) + 6H₂O(l)

    • [Cu(H₂O)₆]²⁺ is blue, [CuCl₄]²⁻ is yellow

  • Effect of adding concentrated hydrochloric acid: equilibrium shifts to the right (towards products), producing more [CuCl₄]²⁻ and turning the solution yellow.

  • Effect of adding a large volume of water to the yellow solution: equilibrium shifts to the left (towards reactants), producing more [Cu(H₂O)₆]²⁺ and turning the solution blue.

Catalysts and Equilibrium

  • A catalyst increases the rate of a reaction without itself being chemically changed by providing an alternative route for the reaction which has a lower activation energy (Ea).

  • When a reaction has reached equilibrium, the rate of the forward and backward reactions are constant, so a catalyst has no effect. This means it does not shift the position of equilibrium in any direction.

  • Before a reaction has reached equilibrium, a catalyst can speed up the time it takes for equilibrium to be achieved.

  • Catalysts are used in industrial processes not to increase the yield of products, but to increase the speed at which equilibrium is reached, cutting down energy costs and reaction time.

  • An increase in temperature before equilibrium is reached also speeds up the reaction because particles have more kinetic energy at higher temperatures.

  • More particles have energy greater than or equal to the required activation energy (Ea), and they collide faster, increasing the frequency of successful collisions and therefore the rate.

  • Once equilibrium is reached, the effect of temperature on the position of equilibrium is determined by whether the reaction is exothermic or endothermic.

Cobalt Equilibrium Demonstration / Class Practical

Safety:

  • Hair tied up, goggles and lab coats should be worn.

  • Do not open the tubes.

  • Follow any additional instructions given by the teacher.

Equilibrium reaction:
[Co(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) ⇌ [CoCl₄]²⁻(aq) + 6H₂O(l), ΔH = +120 kJ/mol

Method:

  1. Set up three small beakers:
    a. Hot (~70°C)
    b. Room temperature
    c. Iced water

  2. Note the colour of each solution.

  3. Swap the tubes between the beakers to observe colour changes.

  4. Complete observations in the table.

Temperature of Water Bath

Colour of Equilibrium Mixture

Substance Present in Greater Amount

Explanation in Terms of Equilibrium

Hot (~70°C)

Blue

[CoCl₄]²⁻

Forward reaction is endothermic; higher temperature favours formation of products

Room Temperature

Purple

Balanced mixture

Neither direction is favoured; forward and reverse reactions occur at the same rate

Cold (iced water)

Pink

[Co(H₂O)₆]²⁺

Reverse reaction is exothermic; lower temperature favours formation of reactants

Equilibrium Position vs Rate

  • Equilibrium reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = –92 kJ/mol (forward reaction exothermic)

Factor Changed

Effect on Equilibrium Position

Explanation

Effect on Yield of NH₃

Effect on Rate

Explanation

Increasing temperature

Shifts to the left

Forward reaction is exothermic; adding heat favours endothermic reverse reaction

Decreases

Increases

Higher temperature increases kinetic energy, so more collisions occur

Decreasing temperature

Shifts to the right

Lower temperature favours exothermic forward reaction

Increases

Decreases

Particles have less kinetic energy, so collisions are less frequent

Increasing pressure

Shifts to the right

Fewer gas molecules on product side (4 → 2), system reduces pressure

Increases

Increases

Higher pressure increases collision frequency

Decreasing pressure

Shifts to the left

More gas molecules on reactant side, system counteracts by favouring reactants

Decreases

Decreases

Lower pressure reduces collision frequency

Actual industrial conditions for ammonia production:

  • 200 atmospheres (high but not too high)

  • 450°C (relatively high)

  • Iron catalyst used to speed up rate without affecting equilibrium position

Reasons for Conditions Used in the Haber Process to Make Ammonia

Pressure of 200 atmospheres

  • N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

  • Higher pressure shifts equilibrium to the right (fewer gas molecules) and increases ammonia yield.

  • Higher pressure increases rate (more collisions).

  • Pressures above 200 atm are expensive and dangerous; 200 atm is a compromise.

Temperature of 450°C

  • N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = –92 kJ/mol

  • Low temperatures favour forward reaction (exothermic), increasing yield.

  • Too low temperatures make reaction slow.

  • 450°C is a compromise between yield and rate.

Iron catalyst

  • Speeds up both forward and reverse reactions without being used up.

  • Does not change equilibrium position or yield.

  • Allows lower temperature to maintain rate, saving energy.