Calculate Ph of 0.11m Solution of Ch3coona Solution

Introduction

Sodium acetate (CH3COONa) is a salt formed by the reaction of acetic acid (CH3COOH) and sodium hydroxide (NaOH). When dissolved in water, it partially dissociates into acetate ions (CH3COO⁻) and sodium ions (Na⁺). The acetate ions can react with water to form acetic acid and hydroxide ions, creating a buffer solution.

The pH of a sodium acetate solution depends on its concentration and the dissociation constants of acetic acid and its conjugate base. The Henderson-Hasselbalch equation allows us to calculate the pH of a buffer solution based on the ratio of the concentrations of the weak acid and its conjugate base.

Calculation Method

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

Where:

• pKa is the acid dissociation constant of acetic acid (pKa = 4.76)
• [CH3COO⁻] is the concentration of acetate ions
• [CH3COOH] is the concentration of acetic acid

For a 0.11M solution of sodium acetate, we assume that the concentration of acetate ions equals the concentration of sodium acetate (since it fully dissociates in water). The concentration of acetic acid is initially zero, but as the solution reaches equilibrium, some acetate ions will convert to acetic acid.

To calculate the pH, we need to determine the equilibrium concentration of acetic acid. This can be done using the dissociation constant of acetic acid (Ka = 1.8 × 10⁻⁵) and the initial concentration of sodium acetate.

Worked Example

Let's calculate the pH of a 0.11M solution of sodium acetate:

1. Assume the initial concentration of acetate ions ([CH3COO⁻]) is equal to the concentration of sodium acetate (0.11M).
2. Let x be the concentration of acetic acid at equilibrium. According to the dissociation reaction, the concentration of acetate ions will be (0.11 - x).
3. The dissociation constant of acetic acid is Ka = 1.8 × 10⁻⁵.
4. Set up the equilibrium expression: Ka = [CH3COOH][OH⁻]/[CH3COO⁻]
5. Assume the solution is dilute, so [OH⁻] ≈ [CH3COOH] = x.
6. Substitute the values: 1.8 × 10⁻⁵ = x² / (0.11 - x)
7. Solve for x using numerical methods or approximation techniques.
8. For a 0.11M solution, x ≈ 0.00018M (very small compared to 0.11M).
9. Now, apply the Henderson-Hasselbalch equation: pH = pKa + log₁₀((0.11 - x)/x)
10. Substitute the values: pH = 4.76 + log₁₀((0.11 - 0.00018)/0.00018)
11. Calculate the logarithm: log₁₀(0.10982/0.00018) ≈ log₁₀(610.1) ≈ 2.785
12. Add to pKa: pH ≈ 4.76 + 2.785 ≈ 7.545

The calculated pH of a 0.11M solution of sodium acetate is approximately 7.55.

Interpreting Results

A pH of 7.55 indicates that the solution is slightly basic. This is expected because sodium acetate is a weak base that dissociates in water to form hydroxide ions. The pH of the solution will depend on the concentration of sodium acetate and the dissociation constants of acetic acid and its conjugate base.

If you need to adjust the pH of the solution, you can add small amounts of a strong acid or base. Adding a strong acid will decrease the pH, while adding a strong base will increase the pH. However, you should be cautious not to exceed the buffering capacity of the solution.