Chapter 7: Equilibrium

Introduction to Chemical Equilibrium

Chemical equilibrium refers to the state where the rates of forward and reverse reactions become equal, and the concentrations of reactants and products remain constant over time. This doesn’t mean the reactions have stopped; rather, they continue at the same rate in both directions, creating a dynamic equilibrium.

Key Characteristics of Equilibrium

1. Dynamic Nature: Reactions continue to occur in both directions at equal rates.
2. Constant Concentrations: The concentrations of reactants and products remain unchanged.
3. Closed System: Equilibrium can only be achieved in a closed system where no matter is lost or gained.
4. Reversibility: The reaction must be reversible for equilibrium to establish.

Law of Mass Action

The Law of Mass Action, proposed by Guldberg and Waage (1864), forms the foundation of chemical equilibrium. It states that the rate of a chemical reaction is directly proportional to the product of the active masses (molar concentrations) of the reactants.

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

The equilibrium constant (K) is expressed as: K = [C]^c × [D]^d / [A]^a × [B]^b

Where [A], [B], [C], and [D] represent the molar concentrations of the respective species at equilibrium, and a, b, c, and d are the stoichiometric coefficients.

Equilibrium Constant

The equilibrium constant (K) is a quantitative measure of the extent to which a reaction proceeds toward completion at a given temperature. The value of K provides important information about the reaction:

• K > 1: The equilibrium lies toward the products (forward reaction is favored).
• K < 1: The equilibrium lies toward the reactants (reverse reaction is favored).
• K = 1: Equal concentrations of reactants and products at equilibrium.

Different Forms of Equilibrium Constants

1. Kc: Equilibrium constant in terms of molar concentrations.
2. Kp: Equilibrium constant in terms of partial pressures (for gaseous reactions).

The relationship between Kp and Kc is given by: Kp = Kc(RT)^Δn

Where:

• R is the gas constant
• T is the absolute temperature
• Δn is the change in the number of moles of gas (products – reactants)

Homogeneous and Heterogeneous Equilibria

Homogeneous Equilibria

All reactants and products are in the same phase (e.g., all gases or all in solution). Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Heterogeneous Equilibria

Reactants and products exist in different phases. Example: CaCO₃(s) ⇌ CaO(s) + CO₂(g)

For heterogeneous equilibria, pure solids and liquids are not included in the equilibrium constant expression as their concentrations remain constant.

Le Chatelier’s Principle

Le Chatelier’s principle states that if a system at equilibrium is subjected to a change in concentration, temperature, pressure, or volume, the system will adjust itself to counteract the effect of the change and establish a new equilibrium.

Factors Affecting Equilibrium

1. Concentration Changes:
   • Adding more reactant shifts the equilibrium to the right (toward products).
   • Adding more product shifts the equilibrium to the left (toward reactants).
2. Temperature Changes:
   • For exothermic reactions (ΔH < 0), increasing temperature shifts equilibrium to the left.
   • For endothermic reactions (ΔH > 0), increasing temperature shifts equilibrium to the right.
3. Pressure Changes (for gaseous reactions):
   • Increasing pressure favors the reaction that produces fewer gas molecules.
   • Decreasing pressure favors the reaction that produces more gas molecules.
4. Catalyst Effects:
   • A catalyst does not change the position of equilibrium; it only helps the system reach equilibrium faster.

Ionic Equilibrium

Ionic equilibrium deals with equilibria involving ions in aqueous solutions, particularly for weak electrolytes.

Acids and Bases

1. Arrhenius Concept:
   • Acids produce H⁺ ions in water.
   • Bases produce OH⁻ ions in water.
2. Brønsted-Lowry Concept:
   • Acids are proton donors.
   • Bases are proton acceptors.
3. Lewis Concept:
   • Acids are electron pair acceptors.
   • Bases are electron pair donors.

Ionization Constants

For a weak acid (HA): HA ⇌ H⁺ + A⁻ Acid ionization constant (Ka) = [H⁺][A⁻]/[HA]

For a weak base (B): B + H₂O ⇌ BH⁺ + OH⁻ Base ionization constant (Kb) = [BH⁺][OH⁻]/[B]

The relationship between Ka and Kb for a conjugate acid-base pair is: Ka × Kb = Kw (where Kw is the ionic product of water = 1.0 × 10⁻¹⁴ at 25°C)

pH Scale

pH is a measure of the hydrogen ion concentration in a solution.

pH = -log[H⁺]

Similarly, pOH = -log[OH⁻]

At 25°C: pH + pOH = 14

The pH scale ranges from 0 to 14:

• pH < 7: Acidic solution
• pH = 7: Neutral solution
• pH > 7: Basic solution

Buffer Solutions

Buffer solutions resist changes in pH when small amounts of acid or base are added. They consist of a weak acid and its conjugate base or a weak base and its conjugate acid.

Examples:

• Acetic acid and sodium acetate
• Ammonia and ammonium chloride

Buffer Capacity

Buffer capacity refers to the amount of acid or base that a buffer can neutralize before experiencing a significant pH change.

Salt Hydrolysis

Salt hydrolysis occurs when ions of a salt react with water to produce either an acidic or basic solution.

Types of salt hydrolysis:

1. Salts of strong acids and strong bases: Neutral solution (e.g., NaCl)
2. Salts of weak acids and strong bases: Basic solution (e.g., CH₃COONa)
3. Salts of strong acids and weak bases: Acidic solution (e.g., NH₄Cl)
4. Salts of weak acids and weak bases: pH depends on relative Ka and Kb values

Solubility Product

For a sparingly soluble salt, the solubility product (Ksp) is the product of the molar concentrations of its ions raised to the power of their stoichiometric coefficients.

For example, for a salt AxBy that dissociates as: AxBy ⇌ xA^y+ + yB^x-

The solubility product is: Ksp = [A^y+]^x × [B^x-]^y

The relationship between solubility (s) and Ksp depends on the stoichiometry of the salt.

Common Ion Effect

The common ion effect is the decrease in solubility of an ionic compound when a common ion is added to the solution. This is an application of Le Chatelier’s principle.

For example, the solubility of AgCl decreases when NaCl is added to the solution due to the common Cl⁻ ion.

Complete Chapter-wise Hsslive Plus One Chemistry Notes

Our HSSLive Plus One Chemistry Notes cover all chapters with key focus areas to help you organize your study effectively:

1. Chapter 1 Some Basic Concepts of Chemistry
2. Chapter 2 Structure of Atom
3. Chapter 3 Classification of Elements and Periodicity in Properties
4. Chapter 4 Chemical Bonding and Molecular Structure
5. Chapter 5 States of Matter
6. Chapter 6 Thermodynamics
7. Chapter 7 Equilibrium
8. Chapter 8 Redox Reactions
9. Chapter 9 Hydrogen
10. Chapter 10 The s Block Elements
11. Chapter 11 The p Block Elements
12. Chapter 12 Organic Chemistry: Some Basic Principles and Techniques
13. Chapter 13 Hydrocarbons
14. Chapter 14 Environmental Chemistry