A 350.0-mL buffer solution is 0.150 M in HF and 0.150 M in NaF. What mass of NaOH can this buffer neutralize before the pH rises above 4.00? If the same volume of the buffer were 0.350 M in HF and 0.350 M in NaF, what mass of NaOH could be handled before the pH rises above 4.00?

Verified step by step guidance

1. Identify the components of the buffer solution: HF (weak acid) and NaF (its conjugate base).
2. Use the Henderson-Hasselbalch equation to find the pH of the buffer: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \), where \( \text{pK}_a \) is the negative logarithm of the acid dissociation constant of HF.
3. Determine the amount of NaOH that can be added before the pH exceeds 4.00 by rearranging the Henderson-Hasselbalch equation to solve for the ratio \( \frac{[\text{A}^-]}{[\text{HA}]} \) when \( \text{pH} = 4.00 \).
4. Calculate the moles of HF and NaF initially present in the buffer using their concentrations and the volume of the solution.
5. For each buffer concentration scenario (0.150 M and 0.350 M), calculate the moles of NaOH that can be added by using the change in moles of HF and NaF required to reach the pH of 4.00, and convert this to mass using the molar mass of NaOH.