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Calculate Ph of 0.01 M Acetic Acid

Reviewed by Calculator Editorial Team

Acetic acid (CH3COOH) is a weak acid that dissociates in water to form acetate ions (CH3COO-) and hydrogen ions (H+). The pH of a solution is a measure of its acidity or basicity, calculated from the concentration of hydrogen ions. This guide explains how to calculate the pH of a 0.01 molar acetic acid solution.

Introduction

The pH of a solution is defined as the negative logarithm (base 10) of the hydrogen ion concentration:

pH = -log[H+]

For weak acids like acetic acid, the hydrogen ion concentration cannot be directly calculated from the acid concentration. Instead, we use the acid dissociation constant (Ka) and the equilibrium expression for the dissociation reaction.

Calculation Method

The dissociation of acetic acid can be represented as:

CH3COOH ⇌ CH3COO- + H+

The acid dissociation constant (Ka) is defined as:

Ka = [CH3COO-][H+]/[CH3COOH]

For a 0.01 M acetic acid solution, we can use the following steps to calculate the pH:

  1. Assume a small amount of acid dissociates, creating equal concentrations of CH3COO- and H+.
  2. Use the Ka value for acetic acid (1.8 × 10-5 at 25°C).
  3. Set up the equation and solve for [H+].
  4. Calculate the pH from the [H+].

Note: This calculation assumes the solution is dilute and that the concentration of water (55.5 M) does not significantly affect the equilibrium.

Example Calculation

Let's calculate the pH of a 0.01 M acetic acid solution:

  1. Assume x = [H+] = [CH3COO-].
  2. The initial concentration of CH3COOH is 0.01 M.
  3. After dissociation, [CH3COOH] = 0.01 - x.
  4. The equilibrium expression becomes:
1.8 × 10-5 = x² / (0.01 - x)

For dilute solutions (x << 0.01), we can approximate:

1.8 × 10-5 ≈ x² / 0.01 x ≈ √(1.8 × 10-5 × 0.01) x ≈ √(1.8 × 10-7) x ≈ 4.24 × 10-4 M

Now calculate the pH:

pH = -log(4.24 × 10-4) ≈ 3.37

Therefore, the pH of a 0.01 M acetic acid solution is approximately 3.37.

Interpretation

A pH of 3.37 indicates that the solution is acidic. This is expected for acetic acid, which is a weak acid with a Ka of 1.8 × 10-5. The relatively high pH value compared to strong acids is due to the limited dissociation of acetic acid in water.

This calculation is most accurate for dilute solutions where the assumption of x << 0.01 holds true. For more concentrated solutions, a more precise method such as solving the quadratic equation would be needed.

FAQ

What is the pH of a 0.01 M acetic acid solution?

The pH of a 0.01 M acetic acid solution is approximately 3.37 at 25°C, calculated using the acid dissociation constant and assuming dilute solution conditions.

Why is acetic acid considered a weak acid?

Acetic acid is considered a weak acid because it only partially dissociates in water, with a relatively low acid dissociation constant (Ka = 1.8 × 10-5). This results in a lower concentration of hydrogen ions compared to strong acids.

How does temperature affect the pH of acetic acid solutions?

The pH of acetic acid solutions increases with temperature because the acid dissociation constant (Ka) increases with temperature. At higher temperatures, more acetic acid dissociates, resulting in a higher concentration of hydrogen ions and a lower pH.