Introduction
Equilibrium is one of the most important concepts in chemistry. It refers to a state in which the rate of forward reaction becomes equal to the rate of backward reaction.
At equilibrium, the concentrations of reactants and products remain constant with time.
Condition of equilibrium:
Rate (forward) = Rate (backward)
Types of Equilibrium
1. Physical Equilibrium
Occurs in physical processes like change of state.
Example:
H2O (l) ⇌ H2O (g)
2. Chemical Equilibrium
Occurs in reversible chemical reactions.
Example:
N2 + 3H2 ⇌ 2NH3
Characteristics of Equilibrium
- Dynamic in nature
- Occurs in closed system
- Concentrations remain constant
- Both forward and backward reactions continue
Law of Chemical Equilibrium
For a general reaction:
aA + bB ⇌ cC + dD
Equilibrium constant:
Kc = [C]^c [D]^d / [A]^a [B]^b
Equilibrium Constant
In Terms of Concentration
Kc = Product concentration / Reactant concentration
In Terms of Pressure
For gases:
Kp = (PC)^c (PD)^d / (PA)^a (PB)^b
Relation Between Kp and Kc
Kp = Kc (RT)^(Δn)
Where:
Δn = moles of gaseous products − moles of gaseous reactants
Reaction Quotient
Q = same expression as K
If Q < K → forward reaction
If Q > K → backward reaction
If Q = K → equilibrium
Homogeneous and Heterogeneous Equilibrium
Homogeneous
All species in same phase
Example:
H2 + I2 ⇌ 2HI
Heterogeneous
Different phases
Example:
CaCO3 ⇌ CaO + CO2
Le Chatelier’s Principle
When a system at equilibrium is disturbed, it shifts to counteract the change.
Effect of Concentration
Increase reactant → shift forward
Increase product → shift backward
Effect of Temperature
Endothermic reaction:
Increase temperature → forward
Exothermic reaction:
Increase temperature → backward
Effect of Pressure
Increase pressure → shift to side with fewer moles
Ionic Equilibrium
Deals with equilibrium in solutions involving ions.
Electrolytes
Strong Electrolytes
Completely ionized
Example:
NaCl → Na+ + Cl−
Weak Electrolytes
Partially ionized
Example:
CH3COOH ⇌ CH3COO− + H+
Degree of Dissociation
α = moles dissociated / total moles
Acid-Base Theories
Arrhenius Theory
Acid produces H+
Base produces OH−
Bronsted-Lowry Theory
Acid = proton donor
Base = proton acceptor
Ionization Constant
For acid:
Ka = [H+][A−] / [HA]
For base:
Kb = [OH−][B+] / [BOH]
pH Scale
pH measures acidity
pH = −log [H+]
Relation Between pH and pOH
pH + pOH = 14
Ionic Product of Water
Kw = [H+][OH−]
At 25°C:
Kw = 1 × 10^-14
Buffer Solutions
Resist change in pH
Acidic buffer:
Weak acid + its salt
Basic buffer:
Weak base + its salt
Henderson Equation
For acidic buffer:
pH = pKa + log ([salt]/[acid])
Solubility Product
For salt:
AB ⇌ A+ + B−
Ksp = [A+][B−]
Common Ion Effect
Addition of common ion decreases solubility
Important Equations Summary
Equilibrium constant:
Kc = products / reactants
Pressure relation:
Kp = Kc (RT)^(Δn)
Acid constant:
Ka = [H+][A−]/[HA]
pH formula:
pH = −log [H+]
Water constant:
Kw = 1 × 10^-14
Common Mistakes
- Confusing Kc and Kp
- Ignoring stoichiometric powers
- Wrong pH calculation
- Forgetting log rules
Exam Tips
- Practice numericals on pH and Kc
- Learn formulas clearly
- Understand Le Chatelier principle deeply
- Focus on weak electrolyte concepts
Conclusion
Equilibrium is a fundamental concept that explains reversible reactions and ionic behavior. Mastering this chapter helps in solving numerical problems and understanding chemical systems.
SEO Keywords
- Equilibrium Class 11 notes
- Chemical equilibrium formulas
- Ionic equilibrium notes
- pH calculation chemistry
FAQs
Q1. What is equilibrium?
Rate forward = rate backward
Q2. What is pH?
pH = −log [H+]
Q3. What is Kc?
Ratio of product and reactant concentration