Multiple Choice
A 30.00-mL sample of 0.125 M HCOOH is being titrated with 0.175 M NaOH. What is the pH after 30.0 mL of NaOH has been added? Ka of HCOOH = 1.8 × 10⁻⁴
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18. Aqueous Equilibrium
Titrations: Weak Acid-Strong Base
Multiple Choice
Consider the titration of 100.0 mL of 0.016 M HOCl (Ka=3.5×10⁻⁸) with 0.0400 M NaOH. Calculate the pH after the addition of 10.0 mL of 0.0400 M NaOH.
A
7.00
B
8.00
C
7.40
D
4.00
Verified step by step guidance1
Step 1: Determine the initial moles of HOCl in the solution. Use the formula \( \text{moles} = \text{concentration} \times \text{volume} \). Calculate the moles of HOCl using the initial concentration (0.016 M) and volume (100.0 mL).
Step 2: Calculate the moles of NaOH added. Use the formula \( \text{moles} = \text{concentration} \times \text{volume} \). Calculate the moles of NaOH using its concentration (0.0400 M) and the volume added (10.0 mL).
Step 3: Determine the moles of HOCl remaining after the reaction with NaOH. Since NaOH is a strong base, it will react completely with HOCl, forming water and the conjugate base OCl⁻. Subtract the moles of NaOH from the initial moles of HOCl to find the remaining moles of HOCl.
Step 4: Calculate the concentration of OCl⁻ formed. Since NaOH reacts completely with HOCl, the moles of NaOH added will equal the moles of OCl⁻ formed. Use the total volume of the solution (initial volume of HOCl plus the volume of NaOH added) to find the concentration of OCl⁻.
Step 5: Use the Henderson-Hasselbalch equation to find the pH of the solution. The equation is \( \text{pH} = \text{pKa} + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \), where \( \text{pKa} = -\log(\text{Ka}) \). Substitute the concentrations of OCl⁻ and HOCl into the equation to calculate the pH.
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