Introduction to Chemical Equilibrium
What is Chemical Equilibrium?
Chemical equilibrium is a state in which the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products.
The system appears static, but actually, both forward and reverse reactions continue to occur at equal rates. This is called dynamic equilibrium.
Equilibrium can only be established in a closed system where no reactants or products can escape.
Key Concept: At equilibrium, the concentrations of reactants and products remain constant, but the reaction continues in both directions at equal rates.
Reversible Reactions and Equilibrium
Characteristics of Reversible Reactions
Reversible reactions can proceed in both forward and backward directions under the same conditions. They are represented by double arrows (⇌).
Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (Haber process for ammonia synthesis)
Initially, only forward reaction occurs as reactants are converted to products. As products accumulate, reverse reaction begins. Equilibrium is reached when both rates equalize.
| Type of Reaction | Characteristics | Example |
|---|---|---|
| Irreversible | Goes to completion, only forward reaction | Burning of fuel |
| Reversible | Does not go to completion, both directions | Esterification |
Equilibrium Constant (Kc and Kp)
Law of Mass Action
For a general reaction: aA + bB ⇌ cC + dD
Equilibrium constant Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ (for concentrations in mol/L)
Kp = (P_C)ᶜ(P_D)ᵈ / (P_A)ᵃ(P_B)ᵇ (for partial pressures in atm)
The value of K indicates the extent of reaction. Large K (>1) means products favored. Small K (<1) means reactants favored.
Important Notes
Pure solids and liquids do not appear in equilibrium constant expressions.
K depends only on temperature, not on concentrations or pressures.
The equilibrium constant has no units for reactions where moles of gases are equal on both sides.
Le Chatelier’s Principle
The Principle
“If a system at equilibrium is subjected to a change, the equilibrium will shift to counteract that change.”
This principle helps predict the direction of equilibrium shift when conditions change.
| Change | Effect on Equilibrium | Example Reaction |
|---|---|---|
| Increase concentration of reactants | Shifts to right (more products) | N₂ + 3H₂ ⇌ 2NH₃ |
| Increase pressure (gases) | Shifts to side with fewer moles | N₂ + 3H₂ ⇌ 2NH₃ |
| Increase temperature | Shifts to absorb heat (endothermic direction) | N₂O₄ ⇌ 2NO₂ (ΔH = +ve) |
| Add catalyst | No shift, only speeds up attainment | All reactions |
Industrial Applications
Haber Process (Ammonia Synthesis)
Reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (ΔH = -92 kJ/mol)
Conditions: High pressure (200 atm), moderate temperature (450°C), iron catalyst
High pressure favors forward reaction (4 moles → 2 moles). Low temperature favors exothermic forward reaction but is compromised for reasonable rate.
Contact Process (Sulfuric Acid)
Reaction: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (ΔH = -197 kJ/mol)
Conditions: Moderate pressure (1-2 atm), temperature (450°C), V₂O₅ catalyst
Excess oxygen used to drive equilibrium forward. Temperature compromise between equilibrium and rate.
Study Strategies for Chemical Equilibrium
1
Master Equilibrium Expressions
Practice writing Kc and Kp expressions for various reactions. Remember to exclude solids and liquids.
2
Understand Le Chatelier’s Principle
Create scenarios and predict shifts. Focus on understanding WHY the shift occurs, not just memorizing rules.
3
Solve Numerical Problems
Practice ICE tables (Initial, Change, Equilibrium) for concentration calculations. Work through industrial process optimizations.