Class 11 Chemistry Ch7 Equilibrium Notes Download

This chapter often feels confusing because reactions seem to “stop” midway — but they actually don’t. That’s where the real concept begins.

In simple words, equilibrium is not about stopping. It’s about balance. Once you understand that reactions continue in both directions at equal rates, the entire chapter starts making sense. Get detailed pattern for CBSE Class 11 Chemistry Syllabus.

In these notes, we’ll break down everything step-by-step so you don’t just memorize formulas - you actually understand how equilibrium works

What is Equilibrium?

Equilibrium is the condition in which the rate of the forward reaction becomes equal to the rate of the backward reaction. At this point, the concentrations of reactants and products remain constant, but both reactions continue at the molecular level. This is known as dynamic equilibrium.

Equilibrium concepts are applied in daily life, industrial chemistry, the atmosphere, biological systems and many natural processes. For example, the solubility of salts, the exchange of gases in the lungs and ammonia synthesis in the Haber process.

Characteristics of Equilibrium

Equilibrium has certain characteristics which are important to form the base of the chapter. 

• It can be attained only in a closed system.
• Both forward and backward reactions continue.
• Concentrations remain constant.
• It can be approached from either side.
• The catalyst does not change equilibrium position.

Types of Equilibrium 

Chemical and physical processes can both attain equilibrium under suitable conditions.

Physical Equilibrium

Physical processes involving two phases of the same substance can reach equilibrium.

• Solid to Liquid Equilibrium

Heat addition melts ice, heat removal freezes water. At equilibrium both phases coexist.

Example: Ice in contact with water at 0°C.

• Liquid to Vapor Equilibrium

Evaporation and condensation occur at the same rate. Vapor pressure depends on temperature.

Example: Water in a closed container.

• Solid to Solution Equilibrium

Dissolution and crystallization occur simultaneously at equilibrium.

Example: Saturated sugar solution.

• Gas to Solution Equilibrium

Governed by Henry’s Law, stating that solubility is directly proportional to the partial pressure of the gas above the liquid.

Example: CO₂ dissolved in soda water.

Chemical Equilibrium

Chemical equilibrium occurs in reversible reactions where reactants and products interconvert continuously.

Example: H₂(g) + I₂(g) ⇌ 2HI(g)

At equilibrium, the rate of formation of HI equals its rate of decomposition.

• Homogenous Equilibrium 

When all reactants and products are in the same phase.

Example: H₂(g) + I₂(g) ⇌ 2HI(g)

• Heterogeneous Equilibrium 

When reactants and products are in different phases.

Example: CaCO₃(s) ⇌ CaO(s) + CO₂(g)

In heterogeneous equilibrium, concentration of pure solids and liquids is taken as constant.

Degree of Dissociation

It represents the fraction of reactant molecules that dissociate at equilibrium.

If α = degree of dissociation,

For weak electrolyte:
K = Cα² / (1 − α)

Equilibrium and Gibbs Free Energy

ΔG° = −RT lnK

If K > 1 → ΔG° is negative → Reaction is spontaneous

If K < 1 → ΔG° is positive → Reaction is non-spontaneous

Law of Mass Action

For a general reversible reaction:

aA + bB ⇌ cC + dD

The equilibrium constant in terms of concentration is:

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

This expression shows how equilibrium depends on the concentration of species involved.

Equilibrium Constant

Kc (Concentration Form)
Used when substances are in solution. Units depend on the reaction stoichiometry.

Kp (Pressure Form)

Used for gaseous reactions: Kp = (pC)^c (pD)^d / (pA)^a (pB)^b

Relation between Kp and Kc
Kp = Kc (RT)^Δn, where Δn = moles of gaseous products minus moles of gaseous reactants.

Reaction Quotient (Q)

Q is calculated using initial concentrations. It predicts the direction of the reaction:

• If Q = K, the system is at equilibrium
• If Q < K, the reaction moves forward
• If Q > K, the reaction shifts backward

Le Chatelier’s Principle

When a change is applied to a system at equilibrium, the system shifts in a direction that counteracts the change.

Examples:

Haber Process

Low temperature and high pressure favor ammonia formation.

N₂ + 3H₂ ⇌ 2NH₃

Thermal Decomposition of CaCO₃

CaCO₃(s) ⇌ CaO(s) + CO₂(g), increasing pressure shifts equilibrium toward the reverse direction, reducing CO₂.

Ionic Equilibrium

Ionic equilibrium deals with the ionization of acids, bases and salts in aqueous solutions.

Acid Base Theories

• Arrhenius Theory: Acid produces H⁺ ions, base produces OH⁻ ions
• Bronsted Lowry Theory: Acid donates a proton, base accepts a proton.
• Lewis Theory: Acid accepts an electron pair, base donates an electron pair.

Ionization Constants

Weak Acid

 HA ⇌ H⁺ + A⁻

 Ka = [H⁺][A⁻] / [HA]

Weak Base

 BOH ⇌ B⁺ + OH⁻

 Kb = [B⁺][OH⁻] / [BOH]

Ostwald’s Dilution Law

For weak acids:

Ka = Cα² / (1 − α)

If α is very small, then:

Ka ≈ Cα²

So,

α = √(Ka / C)

pH Scale

pH = −log[H⁺]

pOH = −log[OH⁻]

pH + pOH = 14 (at 25°C)

Ionization of Water

H₂O ⇌ H⁺ + OH⁻

Kw = 1 × 10⁻¹⁴ at 25°C

Ka × Kb = Kw for conjugate acid-base pairs.

Common Ion Effect

Ionization of a weak acid or base is suppressed by the presence of a strong electrolyte containing a common ion.

Example: CH₃COOH is less ionized when CH₃COONa is added.

Buffer Solutions

A buffer solution resists change in pH when small amounts of acid or base are added.

Acidic Buffer

Made from weak acid and its salt.

Example: CH₃COOH + CH₃COONa

Basic Buffer

Made from a weak base and its salt.

Example: NH₄OH + NH₄Cl 

Henderson Hasselbalch Equation

pH = pKa + log([Salt]/[Acid])

Solubility Product Ksp

Ksp depends only on temperature.

If Ionic Product (Q) > Ksp → precipitation occurs

If Q < Ksp → no precipitation

For AB₂ type salts:

Ksp = [A²⁺][B⁻]²

A salt precipitates if Q > Ksp.

Conclusion

That’s a wrap on Equilibrium Class 11 Chemistry. If you understand the logic behind equilibrium constants, Le Chatelier’s Principle, and pH calculations, this chapter becomes one of the highest scoring ones in Physical Chemistry. Don’t just memorize formulas - practice numericals and revise key concepts regularly. If this helped, share it with a friend who’s stressing over equilibrium right now.

FAQs

Q1. What is dynamic equilibrium in chemistry?

Ans. Dynamic equilibrium is a state in a reversible reaction where the forward and backward reactions occur at equal rates. The concentrations remain constant, but reactions continue at the molecular level.

Q2. How is equilibrium constant related to Gibbs free energy?

Ans. The equilibrium constant (K) is related to standard Gibbs free energy change by ΔG° = −RT lnK. If K is greater than 1, the reaction is spontaneous in the forward direction.

Q3. Why does increasing pressure favor ammonia formation in the Haber process?

Ans. In the Haber process, the product side has fewer moles of gas compared to the reactant side. Increasing pressure shifts equilibrium toward the side with fewer gaseous molecules, favoring ammonia formation.

Q4.  What is the difference between strong and weak acids?

Ans. Strong acids ionize completely in aqueous solution, producing a large concentration of H⁺ ions. Weak acids ionize only partially and establish an equilibrium between ions and unionized molecules.

Q5. What is the common ion effect?

Ans. The common ion effect is the decrease in ionization of a weak electrolyte when a strong electrolyte containing a common ion is added. This shifts the equilibrium backward according to Le Chatelier’s Principle.