Calculate Ph of 0.137 M Sodium Hexanoate Nac6h11o2

Introduction

Sodium hexanoate is a carboxylic acid salt commonly used in various industrial and laboratory applications. When dissolved in water, it undergoes partial dissociation according to the following equilibrium reaction:

The pH of the solution is determined by the concentration of the hexanoate ions (C6H11O2⁻) and the dissociation constant (Kₐ) of the acid. For weak acids like hexanoic acid, the Henderson-Hasselbalch equation is commonly used to calculate the pH.

Calculation Method

The pH of a sodium hexanoate solution can be calculated using the Henderson-Hasselbalch equation:

For a 0.137 M sodium hexanoate solution, we assume complete dissociation (since sodium salts are fully ionized in water), so [C6H11O2⁻] = 0.137 M and [C6H11O2H] = 0.

The dissociation constant (Kₐ) for hexanoic acid is approximately 1.55 × 10⁻⁴ at 25°C.

Example Calculation

Let's calculate the pH of a 0.137 M sodium hexanoate solution:

1. Determine the pKa: pKa = -log(1.55 × 10⁻⁴) ≈ 3.81
2. Calculate the ratio: [C6H11O2⁻]/[C6H11O2H] = 0.137/0 = ∞
3. Apply the Henderson-Hasselbalch equation: pH = 3.81 + log(∞) = 3.81 + ∞ = ∞

This result suggests the solution is highly basic, which aligns with the fact that sodium hexanoate is a strong base in water.

Interpretation

The calculated pH of a 0.137 M sodium hexanoate solution is approximately 14, indicating a strongly basic solution. This is expected because sodium salts are fully ionized in water, and the hexanoate ion (C6H11O2⁻) is a strong base.

For comparison, a 0.137 M solution of a weak acid like acetic acid would have a different pH due to partial dissociation. The strong basic nature of sodium hexanoate solutions is important in applications where high pH is required.