Module 1 Lecture 7 2025 2 per page

Introduction to Buffers

• Definition: A buffer solution is a mixture of a weak acid and its conjugate base (or weak base with its conjugate acid) that maintains a relatively constant pH upon the addition of H3O+ or OH- ions.

• Significance: Essential for maintaining physiological pH levels, e.g., the pH of blood is around 7.4.

Buffer Solution Characteristics

• pH Stability: pH remains constant even with significant dilutions or additions of acids/bases.

• Simple Buffer Example: Acetic acid (CH3COOH) and sodium acetate (CH3COO-) in equal concentrations.

  • Reaction with strong acid:CH3COO- + H3O+ ⟶ CH3COOH + H2O

  • Reaction with strong base:CH3COOH + OH- ⟶ CH3COO- + H2O

Henderson-Hasselbalch Equation

• Purpose: Used to calculate the pH of buffer solutions.

• Equation: [ pH = pK_a + log \left( \frac{[A^-]}{[HA]} \right) ]

• Interpretation: When [A-] = [HA], pH = pK_a, making the buffer equally effective against added acid or base.

pH Calculation Example

• Using Acetic Acid and Sodium Acetate:

  • Given concentrations: [CH3COOH] = 0.100 M, [CH3COO-] = 0.100 M

  • Calculation:[ pH = 4.74 + log \left( \frac{0.100}{0.100} \right) = 4.74 ]

Impact of Strong Acid Addition

• Problem Statement: Adding 100 mL of 0.100 M HCl to a 1.00 L buffer will affect [A-] and [HA] concentrations.

• Buffer Reaction: H3O+ reacts with CH3COO- to form CH3COOH, shifting the concentrations and buffering capacity.

Calculating pH Changes

• Before Reaction:[A-] = 0.100 M (1 L) = 0.100 mol[HA] = 0.100 M (1 L) = 0.100 mol

• After Reaction: Using the Henderson-Hasselbalch equation post-reaction to find new pH reveals only minor pH shifts despite added acid.

  • Calculated pH after addition: 4.65.

Buffer Preparation Techniques

• Options for Creating Buffers:

  • Mixing weak acid with a conjugate base's salt (e.g., acetic acid + sodium acetate).

  • Adding a strong base to a weak acid (e.g., NaOH to CH3COOH).

Buffer Capacity

• Definition: Maximum extent to which a buffer can resist pH changes before exhaustion occurs (when [weak acid] or [conjugate base] is depleted).

• Example Scenario: Adding NaOH to CH3COOH beyond a certain volume depletes CH3COOH, losing buffering ability.

Biological Relevance of Buffers

• Body pH Regulation: Physiological buffer systems such as carbonate, hemoglobin, and phosphate maintain a constant blood pH.

• Hyperventilation Effects: Rapid breathing decreases CO2 levels, increasing blood pH (alkalosis); this can be remedied by re-breathing CO2 into a bag.