Reversible Reactions and Chemical Equilibria

🔄 Reversible Reactions

A reversible reaction is a chemical reaction in which the products can react to form the original reactants again.

• It is represented by a double arrow (⇌), showing that both the forward and backward reactions can occur:
• A+B⇌C+D
• A + B ⇌ C + D

▶️ Forward Reaction

• The forward reaction is the reaction where reactants are converted into products.
• Example: In the reaction N₂(g)+3H₂(g) ⇌ 2NH₃(g)

◀️ Backward Reaction

• The backward reaction is the opposite of the forward reaction, where the products break down to form the original reactants.
• In the same example, the backward reaction is: 2NH₃ → N₂+3H₂

📘 Example of a Reversible Reaction

1. Haber Process: N₂(g)+3H₂(g) ⇌ 2NH₃(g) – Used to produce ammonia from nitrogen and hydrogen.
2. Thermal decomposition of ammonium chloride: NH₄Cl(s) ⇌ NH₃(g)+HCl(g)

⚖️ Chemical Equilibrium

• Chemical equilibrium occurs when the forward and backward reactions happen at the same rate, and the concentrations of reactants and products remain constant (but not necessarily equal).
• It only happens in a closed system (nothing can enter or leave).

🔁 Dynamic Chemical Equilibrium

• The word dynamic means the reactions are still happening, but at equal rates.
• Reactants are being converted to products and vice versa, continuously, but the amounts stay unchanged.
• Example: In the Haber process at equilibrium:
  • Ammonia is being formed and decomposed at the same rate.

⚙️ Factors Affecting Chemical Equilibrium

These factors affect the position of equilibrium, i.e., whether the equilibrium shifts toward the products (right) or reactants (left).

🌡️ 1. Temperature

• Increasing temperature favors the endothermic reaction (absorbs heat).
• Decreasing temperature favors the exothermic reaction (releases heat).

Example (Haber Process: exothermic): N₂+3H₂⇌2NH₃+heat

• ↑ Temperature → equilibrium shifts left → less NH₃.
• ↓ Temperature → equilibrium shifts right → more NH₃.

💨 2. Pressure (only for gases)

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

Example (Haber Process):

• Left side: 1 N₂ + 3 H₂ = 4 molecules
• Right side: 2 NH₃ = 2 molecules
• ↑ Pressure → equilibrium shifts right → more NH₃.
• ↓ Pressure → equilibrium shifts left → more N₂ and H₂.

⚗️ 3. Concentration

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

Example:
If more H₂ is added in the Haber Process, the system makes more NH₃.

⏩ 4. Catalyst Effect

• A catalyst speeds up both the forward and backward reactions equally.
• It does not change the position of equilibrium.
• But it helps the system reach equilibrium faster.

Example: Iron catalyst is used in the Haber Process.

📐 Le Chatelier’s Principle

Le Chatelier’s Principle states:

“If a change is made to a system at equilibrium, the system will adjust itself to oppose that change and restore a new equilibrium.”

In short:

• Add reactant → equilibrium shifts right.
• Remove product → equilibrium shifts right.
• Increase temperature → favors endothermic direction.
• Increase pressure → favors side with fewer gas molecules.

This principle helps in predicting how conditions affect yield in industrial processes.

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