The Oxidizing Power of Bleach

The Oxidizing Power of BleachPurpose:To determine the weight percent of NaClO in commercial bleach, through anoxidation-reduction titration (iodometric analysis).NOTE – notice that this is not the same purpose that was provided in the lab.Procedure:1. Rinse the buret with distilled water2. Rinse the buret with two portions of sodium thiosulfate3. Fill the buret with 0.1 M sodium thiosulfate4. Obtain 25 mL of dilute bleach solution in a 125 mL Erlenmeyer flask5. Add approximately 1 g of KI to the dilute bleach solution6. Add 5 mL of glacial acetic acid to the flask (this is done in the hood)7. Titrate the solution until it turns clear yellow8. Add starch solution drop-wise to the flask until the solution is deep blue9. Continue titrating until the solution turns clear and colorless10. Read sodium thiosulfate volume in buret11. Pour waste in fume hood sink12. Repeat titration two more timesData: See attachedCalculations: See attachedSources of Error:One source of error in this experiment is that the original bleach solution was old,therefore some of the NaClO in the bleach had decomposed over time and the weightpercent will be less than the 6% listed on the commercial bleach container. Anothersource of error is the use of a graduated cylinder to measure the volume of bleachsolution used. It was difficult to determine exactly 25 mL of bleach solution. Thecalculations in step five would be affected if either more or less solution was actuallydelivered. For example, if more than 25 mL’s of bleach solution was used during thetitration; the calculated weight percent of NaClO would have been larger than the actualvalue.Conclusions:Iodometric analysis was performed in order to determine the weight percent ofNaClO in commercial bleach. Potassium iodide was added to a dilute solution ofbleach and glacial acetic acid. A redox reaction occurred between the hypochlorite ionand the iodide ion to produce iodine. The iodine was then titrated against sodiumthiosulfate. Calculations involving the volume of 0.100 M sodium thiosulfate allowedfor the determination of the weight percent of NaClO in commercial bleach. Theexperimental value was calculated to be 2.98%, which is a 50.3% difference from the

published value. Two sources of error that account for this large difference are listedabove. Another error in the lab was the execution of the titration. During the firsttitration our group overshot the end point, which eventually affected the value of theweight percent of NaClO in commercial bleach. Overall, I learned that titrations arenot just for acid-base reactions, and that it can be used in context of redox reactions.Also, the actual chemical reaction that occurred during the titration did not involve theion for which we wanted to obtain information. We were able to determine the value ofsodium hypochlorite in commercial bleach using two balanced equations andstoichiometric calculations.