Introduction
Titrations are often performed with acid and base solutions in order to determine their molarity. Once the enough acid is titrated into the base solution, the mixture will become colorless, indicating the mixture has become neutral. Depending on the chemical reaction at hand, there may be a certain ratio of moles of the acid to the base needed for neutrality. The same process can be formed with a redox reaction. IO3– (aq) + 6S2O3-2 (aq) + 6H+ (aq) —> I– (aq) + 3S4O6-2 (aq) + 3H2O (l) and HOCl (aq) + 2S2O3-2 (aq) + H+ (aq) —> Cl– (aq) + S4O6-2 (aq) + H2O (l) are two general chemical reactions which both can be used in a oxidation-reduction titration. Using the first equation, an unknown molarity of the S2O3-2 component could be calculated and then used in the second equation to find the molarity of an unknown OCl– component. In this experiment, the aforementioned process was performed in order to find the percent composition of household bleach.
Experimental
First, a 0.05 M Na2S2O3 solution was made by dispensing 12.5 mL of a 1.0 M Na2S2O3 into a 250 mL volumetric flask and then filling the volumetric flask with distilled water to the mark on the neck. The solution was then agitated to ensure consistency. Next 50 mL of a 1.507 * 10-2 M KIO3 stock solution was measured into a clean 100 mL beaker. Using a volumetric pipet, 10 mL of the KIO3 solution was dispensed into a 300 mL Erlenmeyer flask. After that, about 25 mL of distilled water and 3 mL of a 3 M KI stock solution were also added and swirled in the Erlenmeyer flask. Finally, 2 mL of a 3 M H2SO4 stock solution was added to the flask, which changed the solution in the flask from being colorless to a deep brown.
Next, the 0.05 M Na2S2O3 solution was put into a buret and was used to titrate the KIO3/KI/H2SO4 mixture. The initial reading on the buret was recorded and titrating began. Once the mixture was a light yellow color, 1.0 mL of starch indicator was dispensed into the Erlenmeyer flask. This changed the mixture to a dark blue-black color. Titrating continued until the mixture became colorless. The final buret reading was recorded. This titrating process was repeated two more times.
For the next part of the experiment, an empty beaker was first weighed and recorded, then about 0.5 mL of bleach was dispensed into it and the new mass of the beaker was recorded. The bleach was then poured into a clean 300 mL Erlenmeyer flask. The beaker was rinsed with distilled water to ensure all the bleach transferred to the Erlenmeyer flask. Then about 25 mL of distilled water was added to the flask along with 3 mL of 3 M KI stock solution and 2 mL of 3 M H2SO4 stock solution. The mixture was swirled to obtain consistency, and then finally 5 drops of 3% ammonium molybdate catalyst was added to the mixture. The mixture was then titrated with the same Na2S2O3 solution from earlier in the same manner for three trials.
Results
Standardization of 0.05 M Na2S2O3 Solution:
| Trial 1 | Trial 2 | Trial 3 | |
| Molarity of standard KIO3, M | 1.507 * 10-2 | 1.507 * 10-2 | 1.507 * 10-2 |
| Volume of KIO3 titrated, mL | 10.00 | 10.00 | 10.00 |
| Moles of KIO3 used | 1.507 * 10-4 | 1.507 * 10-4 | 1.507 * 10-4 |
| Initial buret reading, mL | 1.30 | 19.91 | 3.05 |
| Final buret reading, mL | 19.91 | 38.29 | 21.40 |
| Volume of Na2S2O3 added, mL | 18.61 | 18.38 | 18.35 |
| Moles of Na2S2O3 added | 9.042 * 10-4 | 9.042 * 10-4 | 9.042 * 10-4 |
| Molarity of Na2S2O3 | 4.859 * 10-2 | 4.919 * 10-2 | 4.928 * 10-2 |
Average Molarity, M: 4.902 * 10-2
Precision, ppt: 14.08
Determination of the Oxidizing Capacity of an Unknown Liquid Bleach:
| Trial 1 | Trial 2 | Trial 3 | |
| Mass of beaker and bleach, g | 27.37 | 27.17 | 27.10 |
| Mass of beaker, g | 26.53 | 26.62 | 26.54 |
| Mass of bleach, g | 0.84 | 0.55 | 0.56 |
| Initial buret reading, mL | 21.40 | 0.52 | 19.40 |
| Final buret reading, mL | 49.59 | 19.40 | 38.25 |
| Volume of Na2S2O3, mL | 28.19 | 18.88 | 18.85 |
| Moles of Na2S2O3 | 1.382 * 10-3 | 9.255 * 10-4 | 9.240 * 10-4 |
| Moles of NaOCl | 6.910 * 10-4 | 4.628 * 10-4 | 4.620 * 10-4 |
| Mass of NaOCl, g | 0.05144 | 0.03445 | 0.03439 |
| Percent of NaOCl | 6.124 | 6.264 | 6.141 |
Average Percentage NaOCl: 6.176 %
Precision, ppt: 22.67
Calculations
To find the number of moles of KIO3 used, I used the equation Molarity = Moles/Liter. This equation can be rearranged to Moles = Molarity * Liter. Using my data, I subbed in 1.507 * 10-2 M * 0.01000 L = 1.507 * 10-4 moles of KIO3 used. To find the volume of Na2S2O3 used for titration, I subtracted the initial buret reading from the final buret reading. For example, I did 19.91 mL – 1.30 mL = 18.61 mL for my first trial. To find the number of moles of Na2S2O3 added, I looked at the chemical equation and saw that for every mole of KIO3 used, 6 moles of Na2S2O3 were used, so I multiplied the amount of KIO3 used by 6 to find the number of moles of Na2S2O3 used (6 * 1.507 * 10-4 = 9.042 * 10-4). To find the molarity of the Na2S2O3, I used the equation Molarity = Moles/Liter. Using my data, I subbed in 9.042 * 10-4 moles / 0.01861 L = 4.859 * 10-2 M. To find the average molarity, I added the values for the three trials and then divided by three. Then to find precision, I took the highest molarity (4.928 * 10-2 M) minus the lowest molarity (4.859 * 10-2 M), divided by the average molarity (4.902 * 10-2 M), and finally multiplied by 1000.
For the second part of the experiment, I did the same calculations for finding the volume of Na2S2O3, the average, and precision as I did for the first part. To find the mass of the bleach, I subtracted the weight of the empty beaker from the weight of the beaker with the bleach in it. For the first trial, 27.37 g – 26.53 g = 0.84 g. To find the moles of Na2S2O3, I again used the equation Moles = Molarity * Liters. In my first trial, 4.902 * 10-2 M * 0.02819 L = 1.382 * 10-3 moles. To find the moles of NaOCl, I looked at the chemical equation and saw for every mole of NaOCl, there were two moles of Na2S2O3, so I simply divided the number of moles of Na2S2O3 by two (for the first trial, 1.382 * 10-3 / 2 = 6.910 * 10-4). In order to find the mass of NaOCl, I multiplied the number of moles of NaOCl by the molar mass of NaOCl (74.4419 g). Finally, to find the percent of NaOCl, I divided the mass of NaOCl by the mass of the bleach.
Discussion/Conclusions
My results seem very plausible. I know that standard bleach is diluted quite a bit, so 6.176 % of actual bleach could be very close to the correct answer. The lab procedures ran smoothly for the most part. I had to make a new mixture of KI, H2SO4, and bleach for the second part of the experiment when I put the wrong amount of H2SO4 in the Erlenmeyer flask, but that was not a big deal. When I obtained my bleach the first time, I think the dispenser was set at the wrong amount, as my weight of the bleach was significantly higher for my first trial. This did not seem to effect my final results however, as the percentage of NaOCl for that trial was almost exactly the same as trials two and three. Lastly, I think I probably had some error during my titrations. I think I did not stop the buret of Na2S2O3 at the right moment. I am fairly sure a few extra drops came out, which would have effected my results, making the molarity of the Na2S2O3 lower.
