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Bleaching Analysis

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In this lab our group carefully weighed measurements of two bleaching agents; one liquid and one powder. We found their differences by carefully titrating the reducer, sodium thiosulfate (Na2S2O3) and monitoring the triiodide.


The most common bleaching agent that we use at home contains mostly the hypochlorite ion, Cl0-, which is an oxidizer that takes away colors and dyes (turns our clothes white) in our clothes. The way the bleach is able to do this is by oxidizing the hypochlorite ion with the substance that is creating the stain to get rid of the visible radiation. Which simply means, the light energy in the color penetrates the surface, excites the electrons and then causes them to jump to higher energy states inside the molecule and then we can see this unabsorbed radiation. The removal of these extra electrons in bleach is caused by the hypochlorite ion reacting and removing the electrons in a strong bleach and in a weak bleach there are still some remaining but not many electrons to be excited by the light energy..

In order to understand this process we needed to know that a substance that is reduced gains electrons and a substance that donates electrons is oxidized. Our experiment had us mixing ClO- with 3I- which caused an oxidation reaction and caused a negative one charge on 3I- to oxidize to negative one third which caused a less negative charge then we started with.

ClO-(aq) + 3I-(aq) + H2O (l)  I3- + Cl-(aq) + 2OH-(aq)

Next we needed to oxidize the ion in order to reverse the process so we used the sodium thiosulfate to do this. The first step was to use starch as our indicator to the show the points of titration. The solution will take on a dark red color when I3- is present. When we added sodium thiosulfate, the color disappeared showing us the reversal of the reaction.

The amount of ClO- in bleach can be determined by the following reactions:

3I3-(aq) + 3Starch(aq)  3[I3 * Starch]-(aq)

3[I3 * Starch]-(aq) + 2S2O32-(aq) 3I3-(aq) + S4O62-

The hypochlorite ion was reacted with the iodide ion and was oxidized due to the triiodide ion. We then took the triiodide ion and titrated it with the reducer, sodium thiosulfate, until the triiodide ad returned completely to the Iodide ion, 3I-. There was a noticeable red-brown color seen when the triiodide ion was present, this was our indicator and we proceeded to titrate the sodium thiosulfate until the color was completely gone. The amount of triiodide was then measured and we preformed several stoichiometry calculations in order to determine the amount of ClO- there was in the original bleach sample. Since we already had the determined amount of the hypochlorite ion present in chlorine, we were able to do the calculations for the amount of chlorine present in each of our bleach samples.


Part A: We mixed a primary standard solution of KIO3 (potassium iodate) by measuring around 2 grams (approximately) of KIO3 and dissolving it into a 100 mL volumetric flask with distilled water up to the mark. We then calculated the molarity and set the solution aside to be used later with procedure B.

Procedure B: This part involved standardizing approximately 0.1 M of sodium thiosulfate (Na2S203) which involved taking approximately 6g of the sodium thiosulfate and dissolving it in 250 mL of water (H20). In order to complete this we had to titrate the solution into an oxidized solution shown below:

IO3- (aq) + 8I-(aq) + 6H+(aq)  3I3- (aq) + 3H20

We set up and filled a buret to the 0.00 mL mark with the sodium thiosulfate solution



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