CALCULATION
Standardization of the given Na2S2O3
No. of moles of KIO3 = = 3.224310-3
6H+ + IO3- + 5I- → 3I2 + 3H2O
No. of moles of I2 formed = 3.224310-3 3 = 9.6729 10-3
2S2O32- + I2 → S4O62- + 2I-
No. of moles of I2 used in 25 cm3 solution = 9.6729 10-4
No. of moles of Na2S2O3 used = 9.6729 10-4 2 = 1.9346 10-3
Molarity of the given Na2S2O3 solution = 1.9346 10-3(19.451000)= 0.100M
Finding the amount of Vitamin C by adding excess I2 and back titrate I2 against standard Na2S2O3.
No. of moles of I2 in 25cm3 solution = = 9.6729 10-4
No. of moles of I2 left = (0.1)(8.411000)2 = 4.205 -4
No. of moles of I2 used = 9.6729 x 10-4 – 4.205 x 10-4 = 5.4679 x 10-4
No. of moles of vitamin C used in the 25cm3 solution= 5.4679 x 10-4
Total no. of moles of vitamin C in the 250cm3 solution = 5.4679 x 10-3
Mass of Vitamin C = 5.4679 x 10-3 x ( 2+12x6+6+16x6 )= 0.9624 g
Percentage error = ( 0.9624 – 1.0000 ) /1 X 100% = -3.76%
DISCUSSION
CHOICE OF INDICATOR
In this experiment, starch solution is used to determine the end point of titration of iodine solution against sodium thiosulphate solution. It is because iodine solution could turn starch into blue black, it could help indicating whether iodine is completely reacted. It is because starch could form a blue-black complex with iodine. As the colour change of iodine to iodide is insignificant and difficult to observe, so starch is used as the indicator. However, Starch solution is being added in the later stage of titration. It is because starch would combine with iodine irreversibly when iodine is at a very high concentration. At the endpoint, the blue-black colour would completely disappear.
REASON OF STANDARDIZING SODIUM THIOSULPHATE SOLUTION
We have to standardize the given sodium thiosulphate solution before use. It is because thiosulphate is unstable in acidic medium, it would decompose easily. Also, it would be unstable in the presence of microorganism, Cu2+ and sunlight. So the concentration of the sodium thiosulphate solution might be inaccurate. Therefore, we have to standardize it before use.
REASON OF NOT USING IODINE SOLUTION DIRECTLY
In this experiment, we dissolve known mass of KIO3 into excess KI and H2SO4 solution to form iodine instead of using iodine solution directly.
IO3- + 5I- + 6H+ → 3I2 + 3H2O
This is because iodine is volatile, I2 would escape and cause inaccurate concentration of the solution. Also, iodine would reduce easily, so its amount in the solution would decrease with time. Hence, iodide can be oxidized to iodine in air with the presence of acid, heat and light. As a result, the amount of I2 would increase with time.
4I- + O2 + 4H+ → 2I2 + 2H2O
Moreover, we cannot prepare standard iodine solution by weighing and dissolving certain amount of iodine solid in water because iodine is slight soluble in water. Therefore, we have to prepare standard iodine solution by dissolving known mass of KIO3 into excess KI and dilute sulphuric acid.
SOURCE OF ERROR
ERROR IN READING THE END POINT
When using starch solution as the indicator of titration, it is not easy to determine the end point because the colour difference between yellow and slightly bluish-yellow is insignificant. We have to compare the colour change for several times so as to know whether the sodium thiosulphate has completely reacted with the iodine solution. It would be better if we could find the end point by testing the electric conductivity of the solution. Also, when titrating the solution with vitamin C reacted with iodine solution against sodium thiosulphate, it’s difficult to find the right time to add the starch solution because the vitamin C solution was initially yellow. It’s hard to determine the colour change.
ERROR IN TITRATION
We have encountered problems in finding the titration result because when whirling the solution, some solution left on the flask and they cannot react with the sodium thiosulphate solution. We have to add distilled water to wash the solution into the flask. However, after adding distilled water, the colour of the solution becomes paler. We can hardly make comparison of solution colour with the previous one. We have to add less distilled water to wash the droplets back into the flask next time.
ERROR IN CONCENTRATION OF SOLUTIONS
After standardizing the sodium thiosulphate solution, it might decompose by air and sunlight. As a result, the concentration of the solution would be lower than we expected and the calculated amount of vitamin C would be smaller.
OTHER CHEMICALS IN THE TABLET MIGHT REACT
In this experiment, we have assumed that only vitamin C in the tablet would react with I2. However, we don’t know whether other chemicals and ingredients inside the tablet would react with iodine solution. If some of the ingredients would react with iodine solution, the calculated amount of vitamin C in the tablet would be greater.
CONCLUSION
The mass of vitamin C inside the tablet is 0.9624 g, the tag on the vitamin C tablet claimed that it contain 1.0000 g vitamin C. The percentage error is 3.76%, which is insignificant. So the claim of the tablet maker is justified.
