Chemistry Lesson Note On Reversible Reactions and Chemical Equilibrium – Le Chatelier’s Principle

Topic: Reversible Reactions and Chemical Equilibrium – Le Chatelier’s Principle

Class: SS1

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Specific Objectives:

By the end of this lesson, students should be able to:

1. Differentiate between reversible and irreversible reactions.
2. Explain the concept of chemical equilibrium.
3. State and apply Le Chatelier’s Principle.
4. Identify factors that affect equilibrium position.
5. Appreciate the importance of equilibrium in industrial and biological processes.

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Instructional Materials:

• Charts showing forward and backward reactions
• Beakers/test tubes
• Iron(III) chloride and potassium thiocyanate solution
• Cobalt(II) chloride and water
• Heat source and ice
• Whiteboard and marker for illustrating reaction shifts

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Lesson Content:

Step 1: Reversible and Irreversible Reactions

• Irreversible Reactions: Proceed in one direction only; products cannot change back into reactants.
  Example: Burning of fuel.

• Reversible Reactions: Proceed in both forward and backward directions.
  Example:
  

  This double arrow indicates a reversible reaction.

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Step 2: Concept of Chemical Equilibrium

• In a closed system, when the rate of the forward reaction equals the rate of the backward reaction, the system is in dynamic equilibrium.
• Concentrations of reactants and products remain constant but not necessarily equal.

Characteristics of Equilibrium:

• Occurs in closed systems.
• Dynamic in nature.
• No visible changes, but reactions continue.
• Concentrations of all substances remain constant.

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Step 3: Le Chatelier’s Principle

• States that if a dynamic equilibrium is disturbed by changing conditions (temperature, pressure, or concentration), the position of equilibrium shifts to counteract the change.

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Step 4: Factors Affecting Equilibrium Position

1. Change in Concentration:

• Increasing the concentration of reactants shifts equilibrium to the right (more products formed).
• Increasing the concentration of products shifts equilibrium to the left (more reactants formed).

2. Change in Temperature:

• For exothermic reactions, increasing temperature shifts equilibrium to the left.
• For endothermic reactions, increasing temperature shifts equilibrium to the right.

3. Change in Pressure (for gaseous reactions):

• Increasing pressure shifts equilibrium toward the side with fewer gas molecules.
• Decreasing pressure favors the side with more gas molecules.

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Step 5: Practical Demonstrations

1. Cobalt(II) Chloride Solution Reaction:

• Add heat → blue color (forward).
• Add water → pink color (reverse).

2. Iron(III) and Thiocyanate Reaction:

• Increase Fe³⁺ → deeper red (more product).
• Add water → pale color (dilution shifts backward).

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Step 6: Industrial Application – Haber Process

\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) + \text{heat}

• High pressure (favors fewer molecules) → more NH₃
• Moderate temperature (to balance yield and rate)
• Catalyst: Iron

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Evaluation:

1. Differentiate between reversible and irreversible reactions.
2. What is chemical equilibrium?
3. State Le Chatelier’s Principle.
4. How does pressure affect the position of equilibrium in gaseous reactions?
5. Why is a catalyst used in industrial equilibrium reactions?

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Homework:

1. Explain how an increase in temperature affects an exothermic reaction at equilibrium.
2. Draw and explain a reversible reaction involving a color change (e.g., cobalt(II) chloride).
3. Describe the role of Le Chatelier’s Principle in the production of ammonia.

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Conclusion:

Understanding reversible reactions and how equilibrium can be manipulated is crucial in both chemical industry and biological systems. Le Chatelier’s Principle helps predict and optimize the outcomes of chemical processes.

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