Multiple Choice
20.0 mL of 0.500 M HC2H3O2 is titrated with 0.350 M KOH. Determine the pH after 15.0 mL of the base has been added. Ka for HC2H3O2 = 1.8 × 10⁻⁵.
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18. Aqueous Equilibrium
Titrations: Weak Acid-Strong Base
Multiple Choice
A 25.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. What is the pH of the solution after the addition of 15.0 mL of KOH? The Ka of HF is 3.5×10^-4.
A
pH = 9.26
B
pH = 7.00
C
pH = 3.14
D
pH = 4.74
Verified step by step guidance1
Calculate the initial moles of HF using the formula: \( \text{moles of HF} = \text{volume (L)} \times \text{molarity (M)} \). For a 25.0 mL sample of 0.20 M HF, convert the volume to liters and multiply by the molarity.
Calculate the moles of KOH added using the formula: \( \text{moles of KOH} = \text{volume (L)} \times \text{molarity (M)} \). For a 15.0 mL sample of 0.10 M KOH, convert the volume to liters and multiply by the molarity.
Determine the moles of HF remaining and the moles of F\(^-\) formed after the reaction. Since KOH is a strong base, it will react completely with HF: \( \text{HF} + \text{OH}^- \rightarrow \text{F}^- + \text{H}_2\text{O} \). Subtract the moles of KOH from the initial moles of HF to find the moles of HF remaining, and the moles of F\(^-\) formed will be equal to the moles of KOH added.
Use the Henderson-Hasselbalch equation to find the pH of the solution: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{F}^-]}{[\text{HF}]} \right) \). Calculate \( \text{pK}_a \) from \( \text{K}_a \) using \( \text{pK}_a = -\log(\text{K}_a) \). Substitute the concentrations of F\(^-\) and HF into the equation.
Calculate the concentrations of F\(^-\) and HF by dividing the moles of each by the total volume of the solution in liters (initial volume of HF plus volume of KOH added). Substitute these values into the Henderson-Hasselbalch equation to find the pH.
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