Experiment G: Volumetric Analysis for Vitamin C

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The purpose of this experiment is to measure the exact mass of a vitamin C tablet by mixing the filtered and diluted vitamin C solution with a standard potassium iodate solution and standardizing the mixture solution with sodium thiosulphate. Vitamin C, also known as ascorbic acid, is quite essential for human wellbeing. Due to lack of fresh vegetables and fruits, many sailors on merchant ships died from scurvy—the symptoms that human body is heavily lacked on vitamin C. Now, the general public could purchase vitamin c pill from many stores. Vitamin C, like many other vitamins, will expire overtime. It’s important to know when do these vitamin C pills expire. This sound like an impossible job to do. However, vitamin C will degrade overtime, causing the tablet to lose weight. Base on this fact, manufacturers can accurately assign an expiration date for vitamin C pills by measure the weight lost.

Since vitamin C came in terms of a tablet and vitamin C pills are not purely made from vitamin C. It also contains a small dose of sodium. Therefore, dissolve the tablet into distilled water and filter it can ensure that the solution is mostly vitamin C and produce a better result in experiment rather than a chunk of tablet that only react on the surface. Deal to the problem of vitamin C degrade overtime, so the exact the mass of vitamin C cannot be determined. To determine the concentration of vitamin C requires a primary solution to titrate with the vitamin C solution. A primary solution is the solution with accurately known concentration that is calculated from the mass of the solute and the volume of the solution. However, vitamin C solution that was added in to the primary solution only has a limited amount which mean there is excess. To solve this problem, only thing required is adding another solution to react with excess chemical. By subtracting the leftover from the overall. The concertation of vitamin C solution can be determined. Thus, mass of the vitamin C can also be determined. This technique is called back titration. To set up the primary solution requires the following formula:


By adding potassium iodide (KI) and Sulfuric acid (H2SO4), this made the solution within the Erlenmeyer flask turn to ember colour. The process will produce a known quantity of iodine. Afterward, add vitamin C solution into the flask react with iodine in the flask based on the following formula:

This process would not use up all iodine. Therefore, another solution is needed to complete react with the excess iodine. Sodium thiosulphate (Na2S2O3) was added to react with the excess iodine base on the following formula:

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Add Sodium thiosulphate to the flask until the colour turn faint yellow. The endpoint of the titration is indicated with the assistance of starch. The solution will turn dark black when starch is added. Continue the titration until the black colour completely fade away. At this point, the titration is completed. The concentration now can be calculated.


The procedure is written on “First Year Chemistry Lab” Lab G. The following procedure is a modified version. Dissolve vitamin C tablet provided by TA in a clean 250ml with 100ml deionized water and stir the solution with a magnetic stirrer for approximately 10 minutes which ensure the tablet is mostly dissolved. Insert filter paper into a funnel which is hanged above a 500ml volumetric flask and pour the vitamin solution into the filter paper. Rinse the beaker and the funnel to ensure all solution are transfer into the volumetric flask. Fill the volumetric flask until the mark and mix the solution completely. Dilute the 0.0187mol/L concentrated Potassium iodate solution to diluted 0.00187mol/L Potassium iodate solution. Fill a 50 ml burette with the sodium thiosulphate solution that is placed in fume hood. Pipet 25ml of diluted Potassium iodate solution into 125ml Erlenmeyer flask and add a scoop of Potassium iodide (≈ 0.2g) and 20 drops of 1M of Sulfuric acid. The current colour is ember. Titrate the solution in the flask until the colour of the solution is faint yellow. At this point add 20 to 25 drops of starch which turn the colour of the solution within the flask black. Continue the titration until the colour is clear. Repeat the above procedure until two trials are within 0.1ml.


When diluted Potassium iodate solution was mixed with Sulfuric acid and Potassium iodide powder, the solution appeared to be a ember colour. After vitamin C solution was added into the mixture, the color became yellow. The yellow colour gradually fade into a faint yellow when the first part of the titration is at the endpoint. The starch was colorless. When strch was added in to the faint yellow solution, the solution inside the flask turn into black colour. The endpoint for the second part of the titration had a colourless solution. The last drop for the titration only required half a drop.

Entire titration of Lab G is based on the color change of the solution within the flask. Within the primary solution, the known concentration of iodine was reacted to vitamin C first. This process made the solution change colour from ember to yellow. This indicates that some of the iodine had reacted with vitamin C, but there was some excess iodine left. In order to let all excess iodine to react, sodium thiosulphate was added. When the colour turned faint yellow, this indicates that most of the iodine was reacted. However, just a last bit of iodine is left over, and the solution requires better indicator. Therefore, the starch was added. The starch turned the solution black. With the dark fade away and left a colourless solution, this marks that all iodine had been react. This experiment result in two trial. They were 0.11ml apart from each other, but still quite accurate. The average volume of sodium thiosulphate that was used is 17.05 ± 0.04ml.

All equipment that were used in lab has more or less some uncertainty. Two volumetric flasks—250ml and 500ml, had 0.02ml and 0.3 ml uncertainty respectively. The burette consists with 0.02ml of uncertainty. These instrument uncertainties did add up and effects the result. However, within the process of experiment, there were also uncertainty. Especially the last bit of titration, being a bit black and completely colourless could be half a drop. Still sometime dual to human error, one drop might be added in to the solution. Furthermore, within calculation uncertainty can be even larger than before. The final result of 480mg has 2.09% of uncertainty, when the beginning of the calculation only consists of 0.36% uncertainty.

The expected mass of the vitamin C tablet was 500mg. The calculated result was 480mg ± 10mg. Despite that the calculated result is 20mg away form the expected value. The calculation was still quite accurate. The expire date for the vitamin C pill was very close to the experiment day. Only four mouth form expire. Dual to degrade of vitamin C over time, it was expected that the pill will be lighter than before. Also consider that Life brand vitamin C pill was not 100% pure vitamin C, it also had small percentage of sodium inside the pill. Therefore, based one these two reason, the final calculated result was reasonable.


The mass of the vitamin pill was calculated to be 480mg ± 10mg. The uncertainty in terms of relative uncertainty was 2.09%. Compare to the expected result of 500mg, the final experimental result of the vitamin pill is decently accurate.

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