Buffer Solutions

Buffer solutions are aqueous systems that resist changes in pH when small amounts of acid or base are added. They neutralise the added acid or base, maintaining a relatively stable pH. This stability is essential in many biological, chemical, and industrial processes.

Types of Buffer Solutions

1. Acidic Buffer:

   • Composition: Weak acid + its salt (with a strong base).
   • Example: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa).
   • Function: Maintains pH below 7.

2. Basic Buffer:

   • Composition: Weak base + its salt (with a strong acid).
   • Example: Ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl).
   • Function: Maintains pH above 7.

Mechanism

1. Acidic buffer (e.g., acetic acid and sodium acetate):
   • Added H⁺ reacts with acetate ions to form acetic acid.
   • Added OH⁻ reacts with acetic acid to form acetate ions and water.
2. Basic buffer (e.g., ammonium hydroxide and ammonium chloride):
   • Added H⁺ reacts with OH⁻ to form water.
   • Added OH⁻ reacts with ammonium ions to form ammonia and water.

Acetate Buffer (ACF Buffer)

• Components: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa).
• Mechanism: Maintains pH by neutralizing added acids or bases.
• Applications: Used in biological studies, fermentation, pharmaceuticals, and food preservation.

Preparation of Buffer Solutions

1. For Acidic Buffers:

   • Mix a weak acid (like acetic acid) with its conjugate base (like sodium acetate).
   • Example: To prepare a buffer with pH 4.76, mix acetic acid (CH₃COOH) and sodium acetate (CH₃COONa) in water.

2. For Basic Buffers:

   • Mix a weak base (like ammonium hydroxide) with its conjugate acid (like ammonium chloride).
   • Example: To prepare a buffer with pH 9.25, mix ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl).

Steps for Preparation:

• Determine the desired pH and select appropriate weak acid/base and their salts.
• Calculate the molar ratio of the acid/base and salt using the Henderson-Hasselbalch equation:

pH = pKa + log ([A−]/[HA] ​)

where [A−] is the concentration of the conjugate base, and

 [HA] is the concentration of the weak acid.

• Dissolve the components in distilled water and adjust the final volume.

Example:

Preparation of 0.1 M Phosphate Buffer (pH 7.4):

1. Select the pKa:
   • Phosphoric acid has multiple pKa values, but for this buffer, the relevant pKa is 7.2, close to the desired pH of 7.4.
2. Henderson-Hasselbalch Equation:

Ph = pKa + log ([Base]/[Acid])

• Since the desired pH (7.4) is slightly higher than the pKa (7.2), the base (Na₂HPO₄) concentration will be slightly higher than the acid (NaH₂PO₄).

3. Calculation of Molar Ratio:
   • The Henderson-Hasselbalch equation gives:

7.4 = 7.2 + log ([Base]/[Acid])

Solving this gives a base-to-acid ratio of approximately 1.6:1.

4. Mixing:
   • Prepare 0.1 M solutions of Na₂HPO₄ and NaH₂PO₄.
   • Mix 61.5 mL of 0.1 M Na₂HPO₄ with 38.5 mL of 0.1 M NaH₂PO₄.
5. Adjust Volume:
   • Add distilled water to bring the total volume to 100 ml.
6. Check pH:
   • Use a pH meter to check the buffer’s pH. Adjust slightly with small amounts of NaOH or HCl if necessary to keep the pH to 7.4.

Applications of Buffer Solutions

1. Biological Systems:
   • Buffers like bicarbonate in blood help maintain physiological pH, which is crucial for enzyme activity and cellular processes.
2. Chemical Reactions:
   • Many reactions, especially in biochemistry and organic chemistry, require a controlled pH to proceed correctly.
3. Pharmaceuticals:
   • Buffers are used to maintain the pH of drug formulations for stability and efficacy.
4. Industrial Processes:
   • Buffers are used in fermentation, dye production, and electroplating to control the pH of the environment.
5. Electrophoresis:
   • Buffers help maintain the pH and charge of molecules during electrophoresis in molecular biology labs.