Lecture Notes: Chemistry of Buffers

Introduction

• Instructor: Mr. Johnson
• Topic: Chemistry of buffers
• Chapter: 6 in the textbook
• Objective: Understand buffers as an extension of acids and bases

Definition of a Buffer System

• Buffer System: A solution resisting pH changes with small additions of strong acids/bases.
• Example: Blood is a well-buffered system, maintaining a pH of ~7.35 for effective oxygen transport.

Components of a Buffer System

• Components: Weak acid and its conjugate base in high, approximately equal concentration.
• Example Buffer System: Acetic acid (weak acid) and acetate (conjugate base).
  • Concentrations set at 0.1 molar.
  • Example of an acidic buffer (pH < 7) because the acid's Ka > base's Kb.

Importance of Concentration in Buffer Systems

• Higher concentrations enhance resistance to pH changes.
• Equal concentrations buffer well against both acids and bases.

Mathematical Derivation: Buffer System Equations

• Ka for Weak Acid:
  • Related to the concentration of acid and conjugate base.
  • Hydronium concentration based on Ka and ratio of acid to conjugate base.
  • For equal concentrations, hydronium equals Ka.
• Henderson-Hasselbalch Equation:
  • Derives from the relationship of hydronium concentration, Ka, and concentrations of weak acid/base.

Application: Buffer System Reactions

• When Adding Strong Acid:
  • Conjugate base reacts with added hydronium, forming weak acid.
  • Results in minimal pH decrease.
• When Adding Strong Base:
  • Strong base reacts with weak acid, forming conjugate base.
  • Results in minimal pH increase.

Practice Problems

• Buffer Systems: Selecting appropriate weak acid/conjugate base pairs for effective buffering.
• Problem-Solving: Calculations involving concentrations and pH adjustments.

Basic Buffer System

• Example: Ammonia (NH3) and Ammonium (NH4+)
  • Basic buffer as NH3's Kb > NH4+'s Ka.
  • Described with similar mathematical proofs as acidic buffers.

Advanced Topic: Henderson-Hasselbalch Equation

• Equation: Relates pH, pKa, and concentration ratios.
• Use Cases:
  • Designing buffer systems with desired pH.
  • Determines buffer capacity based on concentrations.

Preparing Buffer Solutions

• Requirements: Equal or nearly equal concentrations of weak acid/base.
• Practical Application: Calculating appropriate concentrations for desired buffer characteristics.

Conclusion

• Buffer systems are vital in chemistry for maintaining stable pH environments.
• Understanding buffer capacity and the Henderson-Hasselbalch equation is essential for practical applications in designing buffer solutions.
• Reminder: Higher concentration buffers provide greater resistance to pH changes.