Applications of Aqueous Equilibria

Neutralization Reactions

Neutralization reactions occur when an acid reacts with a base to produce a salt and water. The nature of the acid and base (strong or weak) determines the pH of the resulting solution after neutralization.

• Strong Acid + Strong Base: Complete dissociation; pH = 7 after neutralization.

• Weak Acid + Strong Base: Complete neutralization; pH > 7 after reaction.

• Strong Acid + Weak Base: Complete neutralization; pH < 7 after reaction.

• Weak Acid + Weak Base: Does not proceed 100% to completion; pH depends on relative values of Ka and Kb.

The Common-Ion Effect

The common-ion effect describes the shift in equilibrium when a substance providing an ion already present in the equilibrium is added. This effect is governed by Le Châtelier’s Principle and is important in buffer and solubility calculations.

• Example: Adding acetate ions to acetic acid solution suppresses acid dissociation, decreasing [H3O+] and increasing pH.

Buffer Solutions

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

• Mechanism: Addition of OH- is neutralized by the weak acid; addition of H3O+ is neutralized by the conjugate base.

• Examples: CH3COOH/CH3COO-, NH4+/NH3, HF/F-

The Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentrations of the acid and its conjugate base:

• Equation:

• Relationship: ;

Buffer Capacity

Buffer capacity is a measure of how much acid or base a buffer can absorb without significant change in pH. It depends on the concentrations of the acid and base components; higher concentrations yield greater capacity.

• Key Points: Buffer pH does not depend on solution volume, only on the ratio and amounts of acid/base.

Titration Curves and Equivalence Point

Titration is used to determine the concentration of an unknown solution by reacting it with a standard solution. The equivalence point is where stoichiometrically equivalent amounts of acid and base have reacted.

• Strong Acid–Strong Base Titration: pH rises sharply near equivalence point; pH at equivalence is 7.

• Weak Acid–Strong Base Titration: pH at equivalence point > 7 due to basic salt formation.

• Strong Acid–Weak Base Titration: pH at equivalence point < 7 due to acidic salt formation.

Sample Results Table

The following table summarizes pH calculations at various points in the titration of 40.0 mL of 0.100 M HCl with 0.100 M NaOH:

Indicators in Titration

Indicators are used to detect the equivalence point in titrations. The choice of indicator depends on the expected pH at equivalence.

• Phenolphthalein: Suitable for weak acid–strong base titrations (equivalence at pH ~8.7).

• Methyl Red: Suitable for strong acid–strong base titrations (equivalence at pH ~7).

Solubility Equilibria

Solubility Product Constant (Ksp)

The solubility product constant, Ksp, describes the equilibrium between a solid and its ions in a saturated solution. It is used to predict solubility and precipitation.

• General Expression:

• Example: For CaF2:

Calculating Solubility from Ksp

Solubility can be calculated by setting up the equilibrium expression and solving for the concentration of ions.

• Example: For MgF2: If , then M (molar solubility).

Factors Affecting Solubility

• Common Ion Effect: Addition of a common ion decreases solubility due to Le Châtelier’s Principle.

• pH: Solubility of salts containing basic anions increases as solution acidity increases.

• Formation of Complex Ions: Complex formation can increase solubility.

• Amphoterism: Some compounds (e.g., Al(OH)3) are soluble in both strong acid and strong base.

Precipitation and Ion Product (Qc)

Precipitation occurs when the ion product (Qc) exceeds Ksp. If Qc = Ksp, the solution is saturated; if Qc < Ksp, no precipitation occurs.

• Key Points: Compare Qc to Ksp to predict precipitation.

Summary Table: Titration Types and pH at Equivalence

Key Equations

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