9.3 Effect of Heat (Temperature) on a Chemical Equilibrium System

This topic explains how temperature affects a chemical equilibrium system, using two important examples — hydrated copper(II) sulphate and hydrated cobalt(II) chloride.

Definition

Effect of heat on equilibrium: When the temperature of a system at equilibrium is changed, the equilibrium shifts in the direction that counteracts the temperature change.

Simple explanation: Heating or cooling an equilibrium system shifts the balance in a predictable way.

Main Idea

• Heating favours the endothermic direction (the direction that absorbs heat).
• Cooling favours the exothermic direction (the direction that releases heat).

Simple explanation: The system tries to "use up" the extra heat by shifting in the endothermic direction when heated.

Core Concept

The Two Examples

1. Hydrated Copper(II) Sulphate

CuSO₄ · 5H₂O (blue solid) ⇌ CuSO₄ (white solid) + 5H₂O

• Heating: blue → white (equilibrium shifts right)
• Adding water: white → blue (equilibrium shifts left)

2. Hydrated Cobalt(II) Chloride

CoCl₂ · 6H₂O (pink solid) ⇌ CoCl₂ (blue solid) + 6H₂O

• Heating: pink → blue (equilibrium shifts right)
• Adding water: blue → pink (equilibrium shifts left)

Simple explanation: Heating drives off water, changing the colour. Adding water reverses the change.

Detailed Explanation

Activity from the Book

Materials required: Test tube, dropper, heating system, hydrated copper(II) sulphate.

Procedure:

• Place 5 g of hydrated copper(II) sulphate in a test tube and heat slowly.
• Observe the colour change from blue to white.
• Allow the test tube and its contents to cool to room temperature.
• Add a few drops of water to the test tube using a dropper.
• Observe the colour change from white to blue again.

Explanation:

When copper(II) sulphate is heated, the water in it is removed, forming anhydrous copper(II) sulphate, which is a white solid.

This copper(II) sulphate changes back to the hydrated form on adding water.

CuSO₄ · 5H₂O (blue solid) ⇌ CuSO₄ (white solid) + 5H₂O

Simple explanation: Heating removes water (blue → white); adding water restores the hydrated form (white → blue).

Cobalt(II) Chloride Equilibrium

Likewise, hydrated cobalt(II) chloride is a pink solid. When heated, it loses water and becomes anhydrous cobalt(II) chloride, a blue solid. So the equilibrium shifts towards the right. But when water is added to it, it absorbs water and the equilibrium shifts to the left to form hydrated cobalt(II) chloride again.

CoCl₂ · 6H₂O (pink solid) ⇌ CoCl₂ (blue solid) + 6H₂O

Simple explanation: Heating removes water (pink → blue); adding water reverses it (blue → pink).

How Temperature Affects Equilibrium

Simple explanation: The system "fights" the temperature change by shifting in the opposite direction.

Applying This to the Examples

Copper(II) sulphate equilibrium:

CuSO₄ · 5H₂O ⇌ CuSO₄ + 5H₂O

• Forward reaction (blue → white): endothermic (requires heat)
• Reverse reaction (white → blue): exothermic (releases heat)

Heating: Shifts right (endothermic direction) → blue → white.

Cooling: Shifts left (exothermic direction) → white → blue.

Cobalt(II) chloride equilibrium:

CoCl₂ · 6H₂O ⇌ CoCl₂ + 6H₂O

• Forward reaction (pink → blue): endothermic
• Reverse reaction (blue → pink): exothermic

Heating: Shifts right (endothermic direction) → pink → blue.

Cooling: Shifts left (exothermic direction) → blue → pink.

Key Point from the Book

The book states:

> "When copper(II) sulphate is heated, the water in it is removed, forming anhydrous copper(II) sulphate, which is a white solid. This copper(II) sulphate changes back to the hydrated form on adding water."

This shows that temperature and water concentration both affect the equilibrium.

Intermediate Level

Why Heating Shifts the Equilibrium

• Heating adds thermal energy to the system.
• The system responds by absorbing this energy.
• The endothermic reaction absorbs heat, so it is favoured.
• So the equilibrium shifts in the endothermic direction.

Simple explanation: The system uses the extra heat by shifting in the direction that absorbs heat.

Why Cooling Shifts the Equilibrium

• Cooling removes thermal energy from the system.
• The system responds by releasing heat to warm up.
• The exothermic reaction releases heat, so it is favoured.
• So the equilibrium shifts in the exothermic direction.

Simple explanation: The system releases heat to counteract the cooling.

Determining Which Direction Is Endothermic

For the copper(II) sulphate equilibrium:

CuSO₄ · 5H₂O ⇌ CuSO₄ + 5H₂O

• Heating drives off water → forward reaction is endothermic.
• So ΔH is positive for the forward reaction.
• Reverse reaction is exothermic.

For the cobalt(II) chloride equilibrium:

CoCl₂ · 6H₂O ⇌ CoCl₂ + 6H₂O

• Heating drives off water → forward reaction is endothermic.
• Reverse reaction is exothermic.

Simple explanation: Removing water requires heat, so the forward reaction is endothermic.

Le Chatelier's Principle (Concept)

This behaviour is an example of Le Chatelier's principle:

> When a system at equilibrium is subjected to a change, the system shifts in the direction that counteracts the change.

Simple explanation: The system "fights back" against any change imposed on it.

Important Points

• - Temperature affects the position of equilibrium.
• - Heating shifts equilibrium in the endothermic direction.
• - Cooling shifts equilibrium in the exothermic direction.
• - Hydrated copper(II) sulphate: blue; anhydrous copper(II) sulphate: white.
• - Hydrated cobalt(II) chloride: pink; anhydrous cobalt(II) chloride: blue.
• - Heating removes water of crystallization → colour change.
• - Adding water restores the hydrated form.
• - Equations:
• - CuSO₄·5H₂O (blue) ⇌ CuSO₄ (white) + 5H₂O
• - CoCl₂·6H₂O (pink) ⇌ CoCl₂ (blue) + 6H₂O

Remember

• - Heating → shifts equilibrium in endothermic direction.
• - Cooling → shifts equilibrium in exothermic direction.
• - CuSO₄·5H₂O (blue) ⇌ CuSO₄ (white) + 5H₂O.
• - CoCl₂·6H₂O (pink) ⇌ CoCl₂ (blue) + 6H₂O.
• - Heating: blue → white; pink → blue.
• - Adding water: white → blue; blue → pink.

Common Mistakes

• 1. Thinking heating always shifts equilibrium to the right. It depends on whether the forward reaction is endothermic or exothermic.
• 2. Confusing the colours. CuSO₄·5H₂O is blue; anhydrous CuSO₄ is white. CoCl₂·6H₂O is pink; anhydrous CoCl₂ is blue.
• 3. Thinking adding water is the same as heating. Heating drives off water (shifts right); adding water shifts left.
• 4. Forgetting that temperature affects equilibrium. Temperature is a key factor.
• 5. Thinking cooling has no effect. Cooling shifts equilibrium in the exothermic direction.
• 6. Confusing hydrated and anhydrous forms. Hydrated = with water; anhydrous = without water.

Quick Revision

• Effect of heat on equilibrium:
• Heating shifts equilibrium in the endothermic direction.
• Cooling shifts equilibrium in the exothermic direction.
• Copper(II) sulphate equilibrium:
• CuSO₄·5H₂O (blue solid) ⇌ CuSO₄ (white solid) + 5H₂O
• Heating: blue → white; adding water: white → blue.
• Cobalt(II) chloride equilibrium:
• CoCl₂·6H₂O (pink solid) ⇌ CoCl₂ (blue solid) + 6H₂O
• Heating: pink → blue; adding water: blue → pink.
• Why? Heating favours the reaction that absorbs heat (endothermic); cooling favours the reaction that releases heat (exothermic).

60-Second Revision

• Heating → shifts equilibrium in endothermic direction.
• Cooling → shifts equilibrium in exothermic direction.
• CuSO₄·5H₂O (blue) ⇌ CuSO₄ (white) + 5H₂O.
• CoCl₂·6H₂O (pink) ⇌ CoCl₂ (blue) + 6H₂O.
• Warning: don't confuse the colours — CuSO₄·5H₂O is blue; anhydrous CuSO₄ is white. CoCl₂·6H₂O is pink; anhydrous CoCl₂ is blue.