Consider a buffer solution that is 0.50 M in NH3 and 0.20 M in NH4Cl. For ammonia, pKb = 4.75. Calculate the pH of 1.0 L of the solution upon the addition of 30.0 mL of 1.0 M HCl to the original buffer solution.

• A. pH = 8.25
• B. pH = 7.95
• C. pH = 8.05
• D. pH = 8.75

Answer: (C)

Verified step by step guidance

1. Identify the components of the buffer solution: NH3 (ammonia) is the weak base, and NH4Cl provides the conjugate acid NH4+.
2. Use the Henderson-Hasselbalch equation for a buffer solution: \( \text{pH} = \text{pK}_\text{a} + \log \left( \frac{[\text{Base}]}{[\text{Acid}]} \right) \). First, convert \( \text{pK}_\text{b} \) to \( \text{pK}_\text{a} \) using \( \text{pK}_\text{a} = 14 - \text{pK}_\text{b} \).
3. Calculate the moles of NH3 and NH4+ initially present: \( \text{moles of NH3} = 0.50 \text{ M} \times 1.0 \text{ L} \) and \( \text{moles of NH4+} = 0.20 \text{ M} \times 1.0 \text{ L} \).
4. Determine the moles of HCl added: \( \text{moles of HCl} = 1.0 \text{ M} \times 0.030 \text{ L} \). HCl will react with NH3 to form NH4+, changing the concentrations of NH3 and NH4+.
5. Recalculate the concentrations of NH3 and NH4+ after the reaction with HCl, then substitute these values into the Henderson-Hasselbalch equation to find the new pH of the solution.