Calculate the pH of the solution that results from each mixture. a. 150.0 mL of 0.25 M HF with 225.0 mL of 0.30 M NaF b. 175.0 mL of 0.10 M C₂H₅NH₂ with 275.0 mL of 0.20 M C₂H₅NH₃Cl

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1. Identify the type of solution formed in each case. For part (a), mixing HF and NaF forms a buffer solution of a weak acid (HF) and its conjugate base (F⁻). For part (b), mixing C₂H₅NH₂ and C₂H₅NH₃Cl forms a buffer solution of a weak base (C₂H₅NH₂) and its conjugate acid (C₂H₅NH₃⁺).
2. Use the Henderson-Hasselbalch equation to calculate the pH of the buffer solutions. For part (a), use the equation: pH = pKa + log([A⁻]/[HA]), where [A⁻] is the concentration of F⁻ and [HA] is the concentration of HF. For part (b), use the equation: pH = pKa + log([B]/[BH⁺]), where [B] is the concentration of C₂H₅NH₂ and [BH⁺] is the concentration of C₂H₅NH₃⁺.
3. Calculate the moles of each component in the mixtures. For part (a), moles of HF = 0.25 M * 0.150 L and moles of NaF = 0.30 M * 0.225 L. For part (b), moles of C₂H₅NH₂ = 0.10 M * 0.175 L and moles of C₂H₅NH₃Cl = 0.20 M * 0.275 L.
4. Determine the concentrations of each component in the final solution by dividing the moles by the total volume of the mixture. For part (a), total volume = 150.0 mL + 225.0 mL. For part (b), total volume = 175.0 mL + 275.0 mL.
5. Substitute the concentrations into the Henderson-Hasselbalch equation to find the pH of each buffer solution.