Answer to Question #186762 in General Chemistry for Kyle

Question

The standardized Na2S2O3 titrant is then used to determine the amount of hypochlorite in a 5.0 grams commercial bleach sample via reduction with potassium iodide in acidic medium as shown in the succeeding equations.

 

(A) If 77.5 milliliters of the Na2S2O3 titrant is required to reach the starch endpoint for the bleach sample while 2.25 milliliter of the titrant is needed for the blank analysis, what is the percent NaOCl in the commercial bleach? Write answer in THREE SIGNIFICANT FIGURES. Write the full computation of your solution.

Answer 1

The reaction in that acid-base titration was:

CH₃COOH + NaOH --> CH₃COO^- + Na⁺ + H₂O

When the reaction was complete, the number of moles of base used equals the number of moles of acid originally present; knowing the number of moles of acid allowed you to compute the concentration of acid in the vinegar sample.

 We determine how much iodine is formed (and therefore how much hypochlorite was in the bleach) by carefully adding sodium thiosulfate (Na₂S₂O₃) solution to the iodine until its reaction is complete.

OCl^- + 2 I^- + 2 H₃O⁺ --> I₂ + Cl^- + 3 H₂O

I₂ + 2 S₂O₃^2- --> 2 I^- + S₄O₆^2-

Hence,

%$NaOCl \hspace{1mm}in\hspace{1mm} commercial\hspace{1mm} bleach = \dfrac{mass\hspace{1mm} NaOCl \hspace{1mm}titrated}{mass\hspace{1mm} commercial \hspace{1mm}bleach \hspace{1mm}titrated} \times100$ %