Calculate the pH of a solution that is 0.205 M in CH3NH2 and 1.55 M in CH3NH3Br. Assume the Kb for CH3NH2 is 4.4 x 10^-4.

• A. pH = 9.82
• B. pH = 8.45
• C. pH = 10.18
• D. pH = 11.22

Answer: (C)

Verified step by step guidance

1. Identify the components of the solution: CH3NH2 is a weak base and CH3NH3Br is its conjugate acid. This is a buffer solution.
2. Use the Henderson-Hasselbalch equation for a buffer solution: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{Base}]}{[\text{Acid}]} \right) \).
3. Calculate \( \text{pK}_a \) using the relation \( \text{pK}_a = 14 - \text{pK}_b \). First, find \( \text{pK}_b \) from \( \text{K}_b \) using \( \text{pK}_b = -\log(\text{K}_b) \).
4. Substitute the concentrations into the Henderson-Hasselbalch equation: \( [\text{Base}] = 0.205 \text{ M} \) and \( [\text{Acid}] = 1.55 \text{ M} \).
5. Calculate the pH using the values obtained for \( \text{pK}_a \) and the concentrations of the base and acid.