Textbook Question
A 25.0-mL sample of 0.125 M pyridine is titrated with 0.100 M HCl. Calculate the pH at each volume of added acid: one-half equivalence.
453
views
18. Aqueous Equilibrium
Titrations: Weak Base-Strong Acid
Multiple Choice
A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution after the addition of 200.0 mL of HNO3. The Kb of NH3 is 1.8 × 10^-5.
A
5.00
B
9.25
C
1.00
D
7.00
Verified step by step guidance1
Start by calculating the initial moles of NH3 in the solution. Use the formula: \( \text{moles} = \text{concentration} \times \text{volume} \). For NH3, \( \text{moles of NH3} = 0.10 \text{ M} \times 100.0 \text{ mL} \). Convert mL to L by dividing by 1000.
Next, calculate the moles of HNO3 added to the solution using the same formula: \( \text{moles of HNO3} = 0.10 \text{ M} \times 200.0 \text{ mL} \). Again, convert mL to L.
Determine the reaction between NH3 and HNO3. NH3 is a weak base and HNO3 is a strong acid. They react in a 1:1 ratio to form NH4+ and NO3-. Calculate the moles of NH3 remaining after the reaction by subtracting the moles of HNO3 from the initial moles of NH3.
Calculate the concentration of NH4+ formed in the solution. Since the total volume of the solution is now 300.0 mL (100.0 mL NH3 + 200.0 mL HNO3), use the formula: \( \text{concentration of NH4+} = \frac{\text{moles of NH4+}}{\text{total volume in L}} \).
Use the Henderson-Hasselbalch equation to find the pH of the solution. The equation is: \( \text{pH} = \text{pKa} + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right) \). First, find the pKa from the Kb of NH3 using the relation \( \text{pKa} = 14 - \text{pKb} \). Then, substitute the concentrations of NH4+ and NH3 into the equation to find the pH.
Related Videos
Related Practice
