Buffer solutions

Buffer Solutions

Buffer solutions are aqueous systems that resist changes in pH when small amounts of acid or base are added. They are crucial in maintaining the pH of biological and chemical systems. Buffer solutions are typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

Principles of Buffer Action

The ability of a buffer solution to resist pH changes is due to the equilibrium between the weak acid (HA) and its conjugate base (A⁻) in solution:

HA ⇌ H⁺ + A⁻

When an acid (H⁺) is added to the buffer, the equilibrium shifts to the left, minimizing the increase in H⁺ concentration. Conversely, when a base (OH⁻) is added, it reacts with H⁺ to form water, and the equilibrium shifts to the right to replace the H⁺ ions.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution:

pH = pKₐ + \log\left(\frac{[A⁻]}{[HA]}\right)

where:

• pH is the hydrogen ion concentration,
• pKₐ is the negative logarithm of the acid dissociation constant,
• [A⁻] is the concentration of the conjugate base,
• [HA] is the concentration of the weak acid.

This equation is derived from the expression for the acid dissociation constant (Kₐ) and is useful for estimating the pH of buffer solutions.

Types of Buffer Solutions

Acidic Buffers

Acidic buffers are composed of a weak acid and its salt with a strong base. A common example is the acetic acid (CH₃COOH) and sodium acetate (CH₃COONa) buffer system. These buffers maintain a pH less than 7.

Basic Buffers

Basic buffers consist of a weak base and its salt with a strong acid. An example is the ammonia (NH₃) and ammonium chloride (NH₄Cl) buffer system. These buffers maintain a pH greater than 7.

Applications of Buffer Solutions

Buffer solutions are essential in many fields:

• Biological Systems: Buffers maintain the pH of blood and cellular environments, crucial for enzyme activity and metabolic processes. The bicarbonate buffer system is a key component in blood pH regulation.

• Chemical Analysis: Buffers are used in titration to maintain a stable pH, ensuring accurate results.

• Industrial Processes: Buffers are used in fermentation, dyeing, and electroplating processes to maintain optimal pH conditions.

Preparation of Buffer Solutions

To prepare a buffer solution, one must select a weak acid or base with a pKₐ close to the desired pH. The ratio of the concentrations of the acid and its conjugate base (or base and its conjugate acid) is adjusted according to the Henderson-Hasselbalch equation to achieve the desired pH.

Limitations of Buffer Solutions

Buffers have a limited capacity, known as buffer capacity, which is the amount of acid or base the buffer can neutralize before the pH changes significantly. This capacity depends on the concentrations of the buffering agents and the pH range over which the buffer is effective.

See Also

• Acid-base equilibrium
• pH scale
• Titration

References

• Lehninger, A. L., Nelson, D. L., & Cox, M. M. (2008). Principles of Biochemistry. W.H. Freeman.
• Skoog, D. A., West, D. M., Holler, F. J., & Crouch, S. R. (2013). Fundamentals of Analytical Chemistry. Brooks/Cole.