Understanding Chemical Equilibrium in Class 11 Chemistry: Concepts, Laws, Types, and Le Chatelier’s Principle Explained with Real-Life Applications

🧪 Class 11 Chemistry: Chemical Equilibrium Explained for Students

🔷 What is Chemical Equilibrium?

Chemical equilibrium is a state in a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction.

➡️ At this point, the concentrations of reactants and products remain constant (not necessarily equal, just constant).

In this Haber process, the reaction can proceed in both directions. When equilibrium is reached, ammonia is formed at the same rate it's decomposed.

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🔷 Characteristics of Equilibrium

1. Dynamic Nature: Reactions are still happening — but at equal rates.

2. No Change in Concentration: Amount of reactants/products stays constant.

3. Can Be Achieved from Either Side: Whether you start with reactants or products, equilibrium will be reached.

4. Occurs in Closed Systems: No exchange of matter with surroundings.

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🔷 Types of Equilibrium

1. Physical Equilibrium: Involves physical processes.

Example: Water ⇌ Vapor

2. Chemical Equilibrium: Involves chemical reactions.

🔷 Law of Chemical Equilibrium (Law of Mass Action)

For a general reaction:

The equilibrium constant (K) is:

Where [X] denotes concentration in moles/litre.

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🔷 Equilibrium Constant (K)

Kc: When concentrations are in mol/L.

Kp: When partial pressures are used (for gases).

🔹 K > 1: Products are favored

🔹 K < 1: Reactants are favored

🔹 K = 1: Moderate amounts of both

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🔷 Factors Affecting Equilibrium (Le Chatelier’s Principle)

📌 When a system at equilibrium is disturbed, it shifts in the direction to minimize the change.

1. Change in Concentration:

Adding reactants → shift right

Removing reactants → shift left

2. Change in Pressure (gases):

Increase pressure → shift to fewer gas molecules

Decrease pressure → shift to more gas molecules

3. Change in Temperature:

For exothermic reactions: ↑ Temp = shift left

For endothermic reactions: ↑ Temp = shift right

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🔷 Ionic Equilibrium (In Aqueous Solutions)

Involves weak acids, weak bases, salts in water:

Acids release H⁺, bases release OH⁻

Equilibrium is established in ionisation reactions

Example:

\ce{CH3COOH <=> CH3COO^- + H^+}

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🔷 Equilibrium in Real Life

Industrial reactions like Haber’s process use pressure/temp tweaks for more product.

Biological systems use equilibrium to maintain pH, respiration balance, etc.

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