Determine the pH of a 1.0 L buffer prepared by adding 0.100 moles of NaOH to 0.250 moles of HF. K_a for HF = 3.5 × 10⁻⁴.

• A. 3.06
• B. 3.28
• C. 3.46
• D. 3.63
• E. 3.85

Answer: (B)

Verified step by step guidance

1. Identify the components of the buffer solution: HF (weak acid) and NaOH (strong base). The reaction between HF and NaOH will produce water and the conjugate base, F⁻.
2. Calculate the moles of HF and NaOH initially present. HF has 0.250 moles, and NaOH has 0.100 moles.
3. Determine the moles of HF and F⁻ after the reaction. Since NaOH is a strong base, it will react completely with HF. Subtract the moles of NaOH from HF to find the remaining moles of HF, and the same amount will be converted to F⁻.
4. Use the Henderson-Hasselbalch equation to find the pH of the buffer: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{F}^-]}{[\text{HF}]} \right) \). Calculate \( \text{pK}_a \) from \( K_a \) using \( \text{pK}_a = -\log(K_a) \).
5. Substitute the concentrations of F⁻ and HF into the Henderson-Hasselbalch equation. The concentrations are the moles of each divided by the total volume of the solution (1.0 L). Calculate the pH using these values.