Chemical Equilibrium - Complete Summary

Introduction

Chemical equilibrium is a state in which the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant over time.

Chemical Equilibrium is the state at which the concentration of reactants and products do not change with time, i.e., concentrations of reactants and products become constant.

Characteristics of Equilibrium State

Constant Properties

Pressure, density, color, concentration become constant

Measurable properties don't change with time

Closed System

Equilibrium can only be achieved in closed vessels

No exchange of matter with surroundings

Dynamic Nature

Reactions continue in both directions

Forward rate = Backward rate

Thermodynamic Condition

At equilibrium, ΔG = 0

ΔH = TΔS

Reversible and Irreversible Reactions

Reversible Reactions

Reactions in which the entire amount of reactants is not converted into products.

Characteristics

Can be started from either side

Never go to completion

Tend to attain equilibrium (ΔG = 0)

Represented by ⇌ or = symbol

Examples

CH₃COOH + NaOH ⇌ CH₃COONa + H₂O

PCl₅(g) ⇌ PCl₃(g) + Cl₂(g)

CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O

Irreversible Reactions

Reactions in which the entire amount of reactants is converted into products.

Characteristics

Proceed only in one direction

Can proceed to completion

ΔG < 0

Represented by → symbol

Examples

NaOH + HCl → NaCl + H₂O

BaCl₂ + H₂SO₄ → BaSO₄↓ + 2HCl

2KClO₃ → 2KCl + 3O₂↑

Law of Mass Action and Equilibrium Constant

Law of Mass Action

"The rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants at a constant temperature at any given time."

Equilibrium Constant Expressions

For a general reaction: aA + bB ⇌ cC + dD

K_c = [C]^c[D]^d / [A]^a[B]^b

K_p = (P_C)^c(P_D)^d / (P_A)^a(P_B)^b

K_x = (X_C)^c(X_D)^d / (X_A)^a(X_B)^b

Relations Between Equilibrium Constants

K_p = K_c(RT)^Δn

K_p = K_x(P)^Δn

Where Δn = (number of moles of gaseous products) - (number of moles of gaseous reactants)

Value of Δn	Relation between K_p and K_c	Units of K_p	Units of K_c
0	K_p = K_c	No unit	No unit
> 0	K_p > K_c	(atm)^Δn	(mol L⁻¹)^Δn
< 0	K_p < K_c	(atm)^Δn	(mol L⁻¹)^Δn

Characteristics of Equilibrium Constant

Concentration Independence

Value is independent of original concentration of reactants

Temperature Dependence

Has definite value at particular temperature

Varies with change in temperature

Forward and Backward Reactions

K_forward = 1/K_backward

Extent of Reaction

Value tells extent of forward or reverse direction

Catalyst Independence

Unaffected by presence of catalyst

Stoichiometry Dependence

Value depends on stoichiometry of chemical equation

Van't Hoff Equation

log K₂ - log K₁ = -ΔH/(2.303R) [1/T₂ - 1/T₁]

Endothermic Reactions (ΔH = +ve)

K increases with temperature

T₂ > T₁ ⇒ K₂ > K₁

Exothermic Reactions (ΔH = -ve)

K decreases with temperature

T₂ > T₁ ⇒ K₁ > K₂

Applications of Equilibrium Constant

Judging Extent of Reaction

K_c > 10³

Products predominate over reactants

Reaction proceeds almost to completion

K_c < 10⁻³

Reactants predominate over products

Reaction proceeds hardly at all

10⁻³ < K_c < 10³

Appreciable concentrations of both reactants and products

Equilibrium mixture contains significant amounts of both

Predicting Direction of Reaction

Reaction Quotient (Q) = [X][Y]/[A][B] (ratio of product of concentrations of products to that of reactants)

Q < K

Reaction proceeds in forward direction

Toward products

Q > K

Reaction proceeds in reverse direction

Toward reactants

Q = K

Reaction is at equilibrium

No net change

Types of Equilibria

Homogeneous Equilibrium

All reactants and products are in the same phase.

Examples

C₂H₅OH(l) + CH₃COOH(l) ⇌ CH₃COOC₂H₅(l) + H₂O(l)

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

Heterogeneous Equilibrium

Reactants and products are present in different phases.

Examples

2NaHCO₃(s) ⇌ Na₂CO₃(s) + CO₂(g) + H₂O(g)

CaCO₃(s) ⇌ CaO(s) + CO₂(g)

H₂O(l) ⇌ H₂O(g)

Important Note

For heterogeneous equilibria, the concentrations of pure solids and pure liquids are not included when writing the equilibrium constant expression.

Le Chatelier's Principle

Le Chatelier's Principle: "Change in any of the factors that determine the equilibrium conditions of a system will shift the equilibrium in such a manner to reduce or to counteract the effect of the change."

Effect of Various Factors on Equilibrium

Change Imposed	Equilibrium Shifts	Effect on K	Remarks
Concentration of reactants increased	To right (forward)	No change	-
Concentration of products increased	To left (backward)	No change	-
Pressure increased (Δn < 0)	To right	No change	Favors side with fewer moles
Pressure increased (Δn > 0)	To left	No change	Favors side with fewer moles
Temperature increased (exothermic)	To left	Decreases	Favors endothermic direction
Temperature increased (endothermic)	To right	Increases	Favors endothermic direction
Catalyst added	No change	No change	Equilibrium achieved faster

Industrial Applications

Haber's Process

N₂ + 3H₂ ⇌ 2NH₃ + 23 kcal

High pressure (Δn < 0)

Low temperature (exothermic)

Excess N₂ and H₂

Contact Process

2SO₂ + O₂ ⇌ 2SO₃ + 45 kcal

High pressure (Δn < 0)

Low temperature (exothermic)

Excess SO₂ and O₂

PCl₅ Dissociation

PCl₅ ⇌ PCl₃ + Cl₂ - 15 kcal

Low pressure (Δn > 0)

High temperature (endothermic)

Excess PCl₅

Physical Equilibria Applications

Melting of Ice

Ice ⇌ Water - x kcal

Volume decreases (1.09 → 1.01 cc/g)

High pressure favors water formation

High temperature favors water formation

Boiling of Water

Water ⇌ Water Vapour - x kcal

Volume increases

High temperature favors vapor formation

High pressure favors liquid formation

Solubility of Salts

Endothermic dissolution: solubility increases with temperature

Exothermic dissolution: solubility decreases with temperature

Examples: NH₄Cl, KNO₃ (endothermic)

Examples: Ca(OH)₂, NaOH (exothermic)

Relation Between Vapour Density and Degree of Dissociation

x = (D - d) / [d(y - 1)]

x = (M - m) / [(y - 1)m]

Where:
x = degree of dissociation
D = initial vapour density
d = vapour density at equilibrium
M = initial molecular mass
m = molecular mass at equilibrium
y = number of moles of products from one mole of reactant

For PCl₅ dissociation: PCl₅ ⇌ PCl₃ + Cl₂ (y = 2)

x = (D - d)/d

Important Points to Remember

Key Points for JEE Main

At equilibrium, rate of forward reaction = rate of backward reaction
Equilibrium constant K is independent of initial concentrations but depends on temperature
For pure solids and liquids, activity = 1 (not included in K expression)
K_p = K_c(RT)^Δn where Δn = moles of gaseous products - moles of gaseous reactants
Le Chatelier's principle: System counteracts any change imposed on it
Catalyst speeds up attainment of equilibrium but doesn't change K value
For exothermic reactions, K decreases with increase in temperature
For endothermic reactions, K increases with increase in temperature

Do's

Remember the relationship between K_p and K_c

Apply Le Chatelier's principle correctly

Include only gases and solutions in K expression

Check stoichiometry when comparing K values

Don'ts

Don't include solids/liquids in K expression

Don't confuse Q and K

Don't forget units for K_p and K_c

Don't ignore temperature effect on K

JEE Main Weightage

This chapter typically carries 2-3 questions in JEE Main, making it a high-weightage chapter. Questions often focus on equilibrium constant calculations, Le Chatelier's principle applications, and predicting direction of reactions.

Chapter Weightage in JEE Main

Weightage High (2-3 questions)