Posted by **Anonymous** on Monday, April 9, 2012 at 3:02pm.

The bottle of bleach lists the percentage of sodium hypochlorite as 6.0%. If the density of commercial bleach is 1.084 g/mL, how many mL of .150 M sodium thiosulfate is required to reach the endpoint in a titration if a student analyzed a 2.0 mL sample of bleach.

- Chemistry 2 -
**DrBob222**, Monday, April 9, 2012 at 5:57pm
The first thing I would do is calculate the molarity of the bleach.

That is 1.084 g/mL x 1000 mL x 0.06 = 65.04 g and that divided by the molar mass NaOCl = 65.04/74.44 = 0.874M. Therefore, a 2 mL sample will have mols = M x L = 0.874 x 0.002L = 0.00175 mols.

Are you treating the NaOCl with I^- to oxidize it to I2 then titrating the liberated I2 with thiosulfate?

OCl^- + 2I^- ==>I2 + Cl^-

2S2O3^2- + I2 ==>2I^- + S4O6^2-

1 mol OCl^- = 1 mol I2 = 2 mol S2O3

Therefore 1/2 mol NaOCl = 1 mol S2O3

We have 0.00175 mol NaOCl. That will use 1/2 that of thiosulfate.

M thiosulfate = mols/L

You know M and mols; solve for L and convert to mL.

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