What mass of NaC₂H₃O₂ (molar mass = 82.03 g/mol) would need to be added to 0.450 moles of HC₂H₃O₂ to prepare a 1.0 L buffer with a pH = 4.95. K_a for HC₂H₃O₂ = 1.8 × 10⁻⁵.

• A. 0.277 g
• B. 0.730 g
• C. 22.7 g
• D. 59.9 g
• E. 295 g

Answer: (D)

Verified step by step guidance

1. Identify the components of the buffer system: HC2H3O2 (acetic acid) and NaC2H3O2 (sodium acetate). The buffer solution is composed of a weak acid and its conjugate base.
2. Use the Henderson-Hasselbalch equation to relate the pH of the buffer to the concentrations of the acid and its conjugate base: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \). Here, \( \text{pK}_a = -\log(1.8 \times 10^{-5}) \).
3. Rearrange the Henderson-Hasselbalch equation to solve for the ratio \( \frac{[\text{A}^-]}{[\text{HA}]} \): \( \frac{[\text{A}^-]}{[\text{HA}]} = 10^{(\text{pH} - \text{pK}_a)} \). Substitute the given pH and calculated \( \text{pK}_a \) to find this ratio.
4. Calculate the concentration of the conjugate base \([\text{A}^-]\) using the ratio from the previous step and the given moles of acetic acid \([\text{HA}] = 0.450 \text{ moles/L}\). Use the equation \([\text{A}^-] = \text{ratio} \times [\text{HA}]\).
5. Convert the concentration of \([\text{A}^-]\) to mass of NaC2H3O2 using its molar mass: \( \text{mass} = [\text{A}^-] \times \text{molar mass of NaC2H3O2} \times \text{volume of solution in L} \).