Ch. 18: Aqueous Solutions and Ionic Equilibrium

Buffers and Their Action

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. This property is crucial in biological and chemical systems where maintaining a stable pH is essential for proper function.

• Definition: A buffer is a solution containing a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Components: Typically, a buffer consists of a weak acid (HA) and its salt (A-), or a weak base (B) and its salt (BH+).

• Buffer Action: Buffers neutralize added H+ or OH- ions, minimizing pH changes.

• Example: The acetic acid/acetate buffer system: CH3COOH/CH3COO-.

Key Equation:

Additional info: This is the Henderson-Hasselbalch equation, used to calculate the pH of buffer solutions.

Buffer Capacity and Preparation

The buffer capacity is the amount of acid or base a buffer can neutralize before the pH changes significantly. Buffers are prepared by mixing a weak acid/base with its conjugate salt in appropriate ratios to achieve the desired pH.

• Calculating pH Change: After adding a small amount of strong acid or base, use an ICE table or the Henderson-Hasselbalch equation to determine the new pH.

• Preparation: Choose a weak acid/base with a pKa close to the desired pH and mix with its salt.

Acid-Base Titrations

Titrations are used to determine the concentration of an unknown acid or base by reacting it with a standard solution. The titration curve shows how pH changes as titrant is added.

• Equivalence Point: The point at which the amount of acid equals the amount of base during titration.

• Types of Titrations: Strong acid-strong base, weak acid-strong base, weak base-strong acid, and polyprotic acid titrations.

• Buffer Region: The region where the solution acts as a buffer (pH ≈ pKa).

• Calculation: Use stoichiometry and equilibrium concepts to calculate pH at various points.

Key Equation for Strong Acid-Strong Base:

Additional info: This equation is used to find the equivalence point in titrations.

Solubility and Solubility Product (Ksp)

Solubility describes how much of a substance can dissolve in a solvent at equilibrium. The solubility product constant (Ksp) quantifies the equilibrium between a solid and its ions in solution.

• Relationship: The higher the Ksp, the more soluble the compound.

• Calculating Molar Solubility: Use Ksp to determine the maximum amount of solute that can dissolve.

• Common Ion Effect: The presence of a common ion decreases solubility.

Key Equation:

Additional info: For a salt AmBn dissolving to mA+ + nB-.

Selective Precipitation and Complex Ion Formation

Selective precipitation is used to separate ions in a mixture by adding a reagent that precipitates one ion while leaving others in solution. Complex ion formation can increase the solubility of certain salts by forming soluble complexes.

• Predicting Precipitation: Compare the ion product (Q) to Ksp. If Q > Ksp, precipitation occurs.

• Minimum Reagent Concentration: Calculate the lowest concentration needed to start precipitation.

• Complex Ions: Formed when metal ions bind to ligands, increasing solubility.

Key Equation for Precipitation:

If Q > Ksp, a precipitate forms; if Q < Ksp, no precipitate forms.

Summary Table: Key Concepts in Aqueous Ionic Equilibrium