Learning Objectives
- Understand chemical equilibrium and equilibrium constants
- Apply Le Chatelier's principle to predict shifts in equilibrium
- Study ionic equilibrium in aqueous solutions
- Learn about pH, buffer solutions, and solubility product
- Understand acid-base theories and indicators
Key Concepts
Chemical Equilibrium
A dynamic state where the rate of forward reaction equals the rate of backward reaction. Concentrations remain constant but reactions continue.
Equilibrium constant (Kc): For aA + bB ⇌ cC + dD:
Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ (concentrations at equilibrium).
Kp = Kc(RT)^Δn, where Δn = (c+d) - (a+b) moles of gaseous products minus reactants.
K is constant at a given temperature. K >> 1: products favoured. K << 1: reactants favoured.
Reaction quotient (Q): Same expression as K but with any concentrations. If Q < K: forward reaction proceeds. Q > K: backward. Q = K: equilibrium.
Le Chatelier's Principle
If a system at equilibrium is disturbed, it shifts to counteract the disturbance.
- Concentration: Adding reactant shifts right; adding product shifts left.
- Pressure: Increasing pressure shifts toward fewer moles of gas.
- Temperature: Increasing T shifts toward endothermic direction. K changes with temperature.
- Catalyst: No effect on equilibrium position (speeds up both directions equally). K unchanged.
Ionic Equilibrium
Arrhenius Theory: Acid gives H⁺ in water; Base gives OH⁻.
Bronsted-Lowry Theory: Acid is proton donor; Base is proton acceptor. Conjugate acid-base pairs.
Lewis Theory: Acid is electron pair acceptor; Base is electron pair donor.
Ionic Product of Water
Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C
pH = -log[H⁺]; pOH = -log[OH⁻]; pH + pOH = 14 (at 25°C).
Neutral: pH = 7. Acidic: pH < 7. Basic: pH > 7.
Weak Acid/Base Equilibrium
Weak acid HA: Ka = [H⁺][A⁻]/[HA]. [H⁺] = √(Ka × C) for weak acids (if α << 1).
Weak base BOH: Kb = [B⁺][OH⁻]/[BOH]. For conjugate pairs: Ka × Kb = Kw.
Degree of dissociation: α = √(Ka/C). Ostwald's dilution law.
Buffer Solutions
Resist change in pH on addition of small amounts of acid or base.
Acidic buffer: Weak acid + its salt (e.g., CH₃COOH + CH₃COONa).
Henderson-Hasselbalch equation: pH = pKa + log([salt]/[acid])
Basic buffer: Weak base + its salt. pOH = pKb + log([salt]/[base]).
Solubility Product (Ksp)
For sparingly soluble salt AxBy ⇌ xAʸ⁺ + yBˣ⁻:
Ksp = [Aʸ⁺]ˣ[Bˣ⁻]ʸ
If ionic product (IP) > Ksp: precipitation occurs. IP < Ksp: unsaturated (more dissolves). IP = Ksp: saturated.
Common ion effect: Addition of a common ion decreases solubility of a salt.
Summary
Chemical equilibrium is dynamic with constant concentrations. The equilibrium constant K quantifies the position. Le Chatelier's principle predicts shifts due to disturbances. Acids and bases are defined by three theories. pH measures hydrogen ion concentration. Buffer solutions maintain nearly constant pH. Solubility product governs precipitation and dissolution of sparingly soluble salts.
Important Terms
- Equilibrium Constant (K): Ratio of product to reactant concentrations at equilibrium
- Le Chatelier's Principle: System shifts to oppose disturbance
- pH: Negative logarithm of H⁺ concentration
- Buffer Solution: Solution resisting pH change
- Solubility Product (Ksp): Product of ion concentrations for saturated solution
- Common Ion Effect: Decrease in solubility due to common ion
Quick Revision
- Kc = [products]/[reactants]; Kp = Kc(RT)^Δn
- Q < K: forward; Q > K: backward; Q = K: equilibrium
- Kw = 10⁻¹⁴; pH + pOH = 14; pH = -log[H⁺]
- Ka × Kb = Kw; Henderson: pH = pKa + log([A⁻]/[HA])
- IP > Ksp: precipitation; IP < Ksp: dissolution
- Temperature changes K; Catalyst does not