A 16.4 mL aliquot of 0.182 M H3PO4 is to be titrated with 0.267 M NaOH. What volume (in mL) of the base will it take to reach the equivalence point?

• A. 37.4 mL
• B. 41.2 mL
• C. 24.8 mL
• D. 33.5 mL

Answer: (A)

Verified step by step guidance

1. Identify the balanced chemical equation for the reaction: \( \text{H}_3\text{PO}_4 + 3\text{NaOH} \rightarrow \text{Na}_3\text{PO}_4 + 3\text{H}_2\text{O} \). This shows that one mole of \( \text{H}_3\text{PO}_4 \) reacts with three moles of \( \text{NaOH} \).
2. Calculate the moles of \( \text{H}_3\text{PO}_4 \) in the aliquot using the formula: \( \text{moles} = \text{concentration} \times \text{volume} \). Convert the volume from mL to L by dividing by 1000.
3. Determine the moles of \( \text{NaOH} \) required to reach the equivalence point. Since the stoichiometry of the reaction is 1:3, multiply the moles of \( \text{H}_3\text{PO}_4 \) by 3.
4. Calculate the volume of \( \text{NaOH} \) solution needed using the formula: \( \text{volume} = \frac{\text{moles}}{\text{concentration}} \). Ensure the concentration is in mol/L and convert the final volume from L to mL by multiplying by 1000.
5. Verify the calculation by checking the units and ensuring the stoichiometry of the reaction is correctly applied. Compare the calculated volume with the given options to find the correct answer.