Buffers in Chemistry

Buffers play a critical role in both biological and industrial systems by maintaining pH levels and preventing abrupt changes. An understanding of buffers is essential for various applications.

Buffer Fundamentals

• Buffer Solution: Buffers are solutions that resist changes in pH when small amounts of acids or bases are introduced. They act as stabilisers within systems.

• Importance:

  • Biological Applications: Buffers help maintain enzyme activity, which is crucial for homeostasis. For example, blood buffer systems maintain a pH of approximately 7.4.
  • Industrial Applications: In processes like food preservation and pharmaceuticals, buffers ensure both safety and efficiency.

Differentiation Between Acidic and Basic Buffers

• Acidic Buffers:

  • Composition: Comprised of a weak acid and its conjugate base.
  • Applications: Utilised in food preservation and pharmaceuticals.

• Basic Buffers:

  • Composition: Consist of a weak base and its conjugate acid.
  • Applications: Commonly found in cleaning solutions.

Illustration of Buffer Action with Equations

• General Role: Buffers mitigate the impact of added H⁺ or OH⁻ ions.
  • Reaction with added H⁺:
    • Equation: $\mathrm{CH_3COO^-} + \mathrm{H^+} \rightleftharpoons \mathrm{CH_3COOH}$
    • Explanation: Acetate ions convert to acetic acid, minimising pH changes.
  • Reaction with added OH⁻:
    • Equation: $\mathrm{CH_3COOH} + \mathrm{OH^-} \rightleftharpoons \mathrm{CH_3COO^-} + \mathrm{H_2O}$
    • Explanation: Acetic acid reacts to form water, stabilising the pH.

Diagram Demonstration

• Diagram shows the pH range where buffers are most effective:

  

Henderson-Hasselbalch Equation

Equation Overview:

• A fundamental formula for calculating buffer pH:

  $\mathrm{pH} = \mathrm{p}K_a + \log \left(\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}\right)$

• Derivation: Derivation includes acid dissociation and the equilibrium constant $K_a$ to relate pH.

  

Worked Examples

Acetic Acid/Sodium Acetate Buffer Calculation:

• Problem: Calculate the pH of a buffer containing 0.1 M acetic acid ($\mathrm{CH_3COOH}$) and 0.1 M sodium acetate ($\mathrm{CH_3COONa}$) where $\mathrm{p}K_a$ of acetic acid is 4.76.
• Solution:
  • Using the Henderson-Hasselbalch equation:
  • $\mathrm{pH} = 4.76 + \log \left( \frac{0.1}{0.1} \right)$
  • $\log(1) = 0$
  • Therefore, $\mathrm{pH} = 4.76$

Ammonia/Ammonium Chloride Buffer:

• Problem: Calculate the pH of a buffer containing 0.2 M ammonia ($\mathrm{NH_3}$) and 0.2 M ammonium chloride ($\mathrm{NH_4Cl}$) where $\mathrm{p}K_b$ of ammonia is 4.75.
• Solution:
  • First, calculate pOH: $\mathrm{pOH} = \mathrm{p}K_b + \log \left( \frac{[\mathrm{NH_4^+}]}{[\mathrm{NH_3}]} \right)$
  • $\mathrm{pOH} = 4.75 + \log \left( \frac{0.2}{0.2} \right) = 4.75$
  • Then calculate pH: $\mathrm{pH} = 14 - \mathrm{pOH} = 14 - 4.75 = 9.25$

Practical Experiment Guide

• Goal: Evaluate the efficiency and determine the limits of a buffer.
• Procedure:
  1. Prepare the buffer for testing.
  2. Calibrate the pH meter.
  3. Measure the initial pH.
  4. Incrementally add an acid or base and record the pH changes.

Importance of Buffers in Natural Systems

Biological Systems

• Blood pH Regulation: Essential for preventing conditions like acidosis and alkalosis.

Environmental Systems

• Ocean Buffer Capacity: Maintains the pH balance, which is critical for marine ecosystems.

  

• Impact of CO₂ Rise: Causes ocean acidification, affecting ecosystems like the Great Barrier Reef.

Adaptation Strategies and Future Research

• Biological Adaptations: Some coral species and plants adapt through genetic changes.