Buffer Solutions

Buffer Solutions

Definition

• Aqueous solution with a weak acid/base and its conjugate.

• Resists pH change upon addition of small amounts of acid or base.

• Used to maintain constant pH in chemical applications.

Chemistry

• Equilibrium between weak acid (HA) and conjugate base (A-): $HA \leftrightarrow H^+ + A^-$

• Adding $H^+$ shifts equilibrium left; adding $OH^-$ shifts it right (Le Chatelier's principle).

• Acid dissociation constant: $Ka = \frac{[H^+][A^-]}{[HA]}$

• Rearranged: $[H^+] = Ka \times (\frac{[HA]}{[A^-]})$

• Henderson-Hasselbalch equation: $pH = pKa + log(\frac{[A^-]}{[HA]})$ or $pH = pKa + log(\frac{[conjugate \ base \ or \ salt]}{[acid]})$

Making a Buffer

• Use Henderson-Hasselbalch equation to determine the ratio of conjugate base to acid needed for a specific pH.

• $pH$ depends on the ratio of acid/salt concentrations, not actual concentrations.

Calculating pH

• Use the Henderson-Hasselbalch equation.

• Example: 0.20M $CH<em>3COOH$ and 0.30M $CH</em>3COONa$ solution, $Ka = 1.8 \times 10^{-5}$, pH = 4.9.

Buffer Capacity

• Measure of buffer's efficiency in resisting pH changes.

• Indicates amount of acid/base a buffer can handle before losing its resistance.

• Depends on:

  • Proximity of buffer pH to its pKa (within 1-2 pH units).

  • Total buffer concentration.

• Higher concentrations provide greater buffer capacity.

Applications

• Chemical manufacturing and biochemical processes.

• Enzymes require precise pH; buffers prevent denaturation and maintain activity.

• Carbonic acid and bicarbonate buffer in blood plasma (pH 7.35-7.45).

• Industrial uses: fermentation, dyes for fabrics, chemical analysis, pH meter calibration.

• Common buffer: PBS (phosphate buffer saline) at pH 7.4 for biological samples.

Demonstration

• Calculate pH of buffer solution (e.g., 0.2 mol $K<em>2HPO</em>4$ and 0.1 mol $KH<em>2PO</em>4$, $Ka= 6.2\times 10^{-8}$).

• Calculate pH change after adding NaOH or HCl.