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1 Some Basic Concepts of Chemistry 2 Structure of Atom 3 Classification of Elements and Periodicity in Properties 4 Chemical Bonding and Molecular Structure 5 Thermodynamics 6 Equilibrium 7 Redox Reactions 8 Organic Chemistry – Some Basic Principles and Techniques 9 Hydrocarbons

6 Equilibrium

6.1 Equilibrium in Physical Processes 6.2 Equilibrium in Chemical Processes - Dynamic Equilibrium 6.3 Law of Chemical Equilibrium and Equilibrium Constant 6.4 Homogeneous Equilibria 6.5 Heterogeneous Equilibria 6.6 Applications of Equilibrium Constants 6.7 Relationship between Equilibrium Constant K, Reaction Quotient Q and Gibbs Energy G 6.8 Factors Affecting Equilibria 6.9 Ionic Equilibrium in Solutions 6.10 Acids, Bases and Salts 6.11 Ionization of Acids and Bases 6.12 Buffer Solutions 6.13 Solubility Equilibria of Sparingly Soluble Salts

Solubility Equilibria of Sparingly Soluble Salts

NCERT Reference: Chapter 6 – Equilibrium – Page 203

Quick Notes

  • Sparingly soluble salts dissociate only slightly in water.
  • The Solubility Product Constant (Ksp) is the equilibrium expression for the dissolving of a sparingly soluble salt.
  • The common ion effect suppresses solubility when one of the dissociated ions is already present in the solution.
  • Ksp is specific for each salt and varies with temperature.

Full Notes

Introduction

Some salts don’t dissolve well in water – they dissolve just a little, and we call them sparingly soluble salts. Even though they seem almost “insoluble,” a small amount still dissolves and sets up an equilibrium between the solid salt and the ions in solution.

This section helps us understand how to describe that equilibrium using a special constant called the solubility product (Ksp).

Solubility Product Constant (Ksp)

When a sparingly soluble salt dissolves in water, an equilibrium is set up between the solid phase and the ions in solution:

IB Chemistry NCERT Class 11 diagram showing dissolution equilibrium of a sparingly soluble salt between solid and aqueous ions.

The equilibrium constant for this process is called the Solubility Product Constant, or Ksp. It is not the same as solubility – instead, it represents the product of the ion concentrations at equilibrium.

Ksp is temperature-dependent. Although Ksp is derived from activities and technically has no units, we often use molar concentrations in calculations.

Ksp Expression:

IB Chemistry NCERT Class 11 expression for solubility product Ksp as product of equilibrium ion concentrations. IB Chemistry NCERT Class 11 example Ksp constant form for a generic salt MxNy with exponents based on stoichiometric coefficients.

Calculating Solubility from Ksp

Worked Example

What is the molar solubility of PbCl2 if Ksp = 1.6 × 10−5?

  1. Write the dissociation and Ksp expression
    PbCl2(s) ⇌ Pb2+(aq) + 2Cl−(aq)
    Let solubility = x. Then [Pb2+] = x and [Cl−] = 2x.
    Ksp = [Pb2+][Cl−]2 = x(2x)2 = 4x3.
  2. Solve for x
    4x3 = 1.6 × 10−5 ⇒ x3 = 4.0 × 10−6 ⇒ x ≈ 1.6 × 10−2 mol L−1.

Answer: The molar solubility of PbCl2 ≈ 1.6 × 10−2 M.

Key Insight: A low Ksp value indicates a salt is very slightly soluble – but even these salts do dissolve to some extent and create a measurable ionic equilibrium.

Common Ion Effect on Solubility of Ionic Salts

When a solution already contains one of the ions from a dissolving salt, we call it a common ion.

Adding a common ion causes the equilibrium to shift leftward (toward the undissolved salt), according to Le Chatelier’s Principle – thereby reducing the solubility of the salt.

Example AgCl(s) in NaCl(aq)

AgCl(s) ⇌ Ag+(aq) + Cl−(aq)

Adding NaCl increases [Cl−] and equilibrium shifts left this causes more AgCl to precipitate = solubility of AgCl decreases.

IB Chemistry NCERT Class 11 schematic showing common ion effect: added chloride drives AgCl equilibrium toward the solid.

(Note: Ksp remains constant, but the concentrations adjust.)

Calculations

If [Cl−] is known (e.g., from added NaCl), we can calculate the reduced solubility:

Ksp = [Ag+][Cl−] ⇒ [Ag+] = Ksp / [Cl−]

This gives the new [Ag+], i.e., the solubility in the presence of common ion.

Application in Qualitative Analysis

The common ion effect is a powerful tool in selective precipitation – allowing chemists to control which ions precipitate from solution.

By carefully adjusting the concentration of common ions, specific salts can be selectively removed from mixtures – a key step in qualitative ionic analysis.

Summary