Calculate Ph of 0.10m C4h9nh2

Introduction

Butylamine (C4H9NH2) is an organic amine with a butyl group attached to the amino nitrogen. In aqueous solution, it behaves as a weak base, accepting protons from water to form the butylammonium ion (C4H9NH3+).

The pH of a butylamine solution depends on its concentration and the equilibrium between the undissociated amine and its conjugate acid. This calculator provides a precise method to determine the pH of a 0.10 molar solution of butylamine.

Formula

The pH of a weak base solution can be calculated using the Henderson-Hasselbalch equation, which relates the pH to the concentration of the weak base and its conjugate acid:

Where:

• pKa is the acid dissociation constant of the conjugate acid (C4H9NH3+)
• [C4H9NH2] is the concentration of butylamine
• [C4H9NH3+] is the concentration of butylammonium ion

For a 0.10M solution of butylamine, we assume complete dissociation to form 0.10M C4H9NH3+ and negligible concentration of OH-.

Calculation

To calculate the pH of a 0.10M solution of butylamine:

1. Determine the pKa value for butylammonium ion (C4H9NH3+). Typical values range from 10.6 to 10.8.
2. Assume complete dissociation to form 0.10M C4H9NH3+.
3. Apply the Henderson-Hasselbalch equation with [C4H9NH2] = 0.10M and [C4H9NH3+] = 0.10M.

The result is the pH of the solution, which should be slightly above the pKa value due to the logarithmic relationship.

Interpretation

The calculated pH of a 0.10M butylamine solution will typically be between 11.0 and 11.2. This indicates a strongly basic solution, consistent with the weak base properties of butylamine.

Key points to consider:

• The pH is higher than the pKa of the conjugate acid, as expected for a weak base solution.
• The solution is basic, which is typical for amine solutions.
• For more concentrated solutions, the pH will increase further.