Calculate the pH of 0.250 L of a 0.36 M formic acid–0.30 M sodium formate buffer before and after the addition of (a) 0.0050 mol of NaOH. Assume that the volume remains constant. (b) 0.0050 mol of HCl. Assume that the volume remains constant.

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1. Identify the components of the buffer system: formic acid (HCOOH) and sodium formate (HCOONa). Formic acid is the weak acid, and sodium formate is its conjugate base.
2. Use the Henderson-Hasselbalch equation to calculate the initial 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 formic acid.
3. For part (a), calculate the change in moles of formic acid and formate after adding 0.0050 mol of NaOH. NaOH will react with formic acid to form water and formate ions, changing their concentrations.
4. Recalculate the pH using the Henderson-Hasselbalch equation with the new concentrations of formic acid and formate after the addition of NaOH.
5. For part (b), calculate the change in moles of formic acid and formate after adding 0.0050 mol of HCl. HCl will react with formate ions to form formic acid, changing their concentrations. Recalculate the pH using the Henderson-Hasselbalch equation with the new concentrations.