Iodometric Redox Titration of Vitamin C Tablets

Introduction
Vitamin C is a vital component of a healthy diet which is why, like many others, my father takes vitamin C supplements. However, I noticed that the vitamin C tablets he takes expired in January 2009. These tablets were bought in the USA and developed a light amber tint. Therefore, I wondered whether this would mean that over time, the concentration of vitamin C has decreased. Hence, I researched a scientific method to determine the concentration of vitamin C in order to see whether my father should continue using the expired tablets or rather buy new ones. Chemically known as ascorbic acid, vitamin C is an organic compound containing of six carbon atoms, of which two can be readily oxidized under aqueous acidic conditions or by air over a longer time period.
The method used to measure the concentration of vitamin C is called a reduction oxidation, known as redox, titration. Ascorbic acid reacts with iodine (I₂) to create dehydroascorbic acid (C₆H₈O₆) under acidic aqueous conditions:
C₆H₈O_{6 (aq)} + I_{2 (aq)}  C₆H₆O_{6 (aq)} + 2 I^–_(aq) + 2 H⁺_(aq)
However, as iodine I₂ is not very soluble in water, a complex created by aqueous iodine I_{2 (aq)} and aqueous iodide anion I^–_(aq) through the following reaction must be used.
I_{2 (aq)} + I^–_(aq)   I₃^–_(aq)
I₃^– is known as triiodide, which is much more soluble in water than iodine. The method used to create the triiodide is the reaction of aqueous iodate IO₃^–_(aq) with aqueous iodide I^–_(aq) under acidic aqueous conditions as shown below.
Reaction 1:IO₃^–_(aq) + 8 I^–_(aq) + 6 H⁺_(aq)  3 I₃^–_(aq) + 3 H₂O _(l)              
The reaction of water soluble starch, being a white solution, with triiodide gives a dark blue complex. This change of colour shows the end of the redox titration.
The redox titration is using the reaction of aqueous triiodide I₃^– _(aq) with aqueous ascorbic acid C₆H₈O_{6 (aq)} to form aqueous dehydroascorbic acid C₆H₆O_{6 (aq)} and aqueous iodide I^–_(aq) under acidic aqueous conditions.
Reaction 2: C₆H₈O_{6 (aq)} + I₃^– _(aq)  C₆H₆O_{6 (aq)} + 3 I^–_(aq) + 2 H⁺_(aq)                           
The method used is an indirect titration, which means it measures the amount of triiodide remaining in the solution after having reacted with the ascorbic acid. Therefore an excess of aqueous triiodide I₃^– _(aq) is needed.
The excess aqueous triiodide I₃^– _(aq) is reduced by aqueous thiosulfate S₂O₃^2- _(aq) to create aqueous iodide I^–_(aq) and aqueous tetrathionate S₄O₆^2- _(aq) as shown below.
Reaction 3:I₃^– _(aq) + 2 S₂O₃^2- _(aq)  3 I^– _(aq) + S₄O₆^2- _(aq)                           
As soon as all the triiodide is reduced to iodide, the colour changes from dark blue (the triiodide starch complex is dark blue) to white. The oxidation of ascorbic acid is a 1:1 reaction, meaning 1 mol of ascorbic acid requires 1 mol of triiodide to form 1 mol of dehydroascorbic acid, whereas the oxidation of thiosulfate is a 1:2 reaction, meaning 2 mol of thiosulfate can be oxidized to 1 mol tetrathionate by 1 mol of triiodide, all under acidic conditions in water. Knowing this one can calculate the amount of vitamin C in various tablets as long as all of them are readily soluble in water, meaning, for example, not coated.
(CH 212 – Quantitative Analysis. 1-2)
Research Question

What is the effect of the age of vitamin C tablets, expressed by the number of months elapsed after the expiry date, on the concentration in percentage weight of vitamin C in said tablets?

Variables

Materials

• 1 g of soluble starch
• 8 M (mol/l) sulfuric acid (H₂SO₄) pure
• Potassium iodide (KI), 95%, pure, DAB
• Potassium iodate (KIO₃), p. A. EMSURE® ACS, ISO, Ph Eurpure, DAB
• Vitamin C (ascorbic acid (C₆H₈O₆))tablets or sachets (it is recommended to use colourless dissolvable products, otherwise the colour change will be difficult to see)
• Sodium thiosulfate pentahydrate (Na₂S₂O₃ ï‚· 5H₂O), 99,5%, pure, DAB
• Sodium carbonate (Na₂CO₃) 99,5%, pure, DAB
• Distilled water (H₂O)

Apparatus

• One 3 ml ± 0,01 ml measuring pipette
• Two 250 ml ± 5% beakers and two 500 ml ± 5% beakers
• One 50 ml ± 0,08 ml and one 100 ml ± 0,1 ml measuring cylinder
• Two 500 ml ± 0,2 ml volumetric flasks with a cork
• Three 250 ml Erlenmeyer flasks to be used for the titrations
• One 50 ml ± 0,1 ml burette with stand and clamp for burette
• Two funnels, each with a diameter of 9,5 cm

It is necessary to clean and rinse all glassware with distilled water beforehand to avoid impurities and contamination of solution used.

• One electric scale set in grams and preferably to four decimal places g ± 0,1 mg to allow for maximum accuracy (available scale had two decimal places g ± 10 mg)
• One magnetic stirrer, one stirring rod and a mortar with a pestle
• Small spoons, scalpels and cups, in total 5 of each
• Rubber gloves and safety glasses as the reactants used can irritate skin and eyes

Method

Preparation of starch indicator

1. Fill a 250 ml beaker with 100 ml of distilled water, measured with a 50 ml cylinder.
2. Weigh 1 g of soluble starch using the scale and a spoon. Add the starch to the beaker. Using the stirring rod, stir until dissolved. Every day a new solution should be made.

Preparation of sodium thiosulfate

3. Use a 100 ml measuring cylinder to fill 450 ml of distilled water into a 500 ml beaker.
4. Weigh 0,05 g of Na₂CO₃ using the scale, a spoon, and a cup and add to beaker.
5. Weigh 8,7 g of Na₂S₂O₃ ï‚· 5H₂O using the scale, a spoon, and a separate, equally clean small cup. Add to the same beaker.
6. Dissolve the chemicals compounds through swirling the beaker.
7. Once dissolved, pour the solution into a clean 500 ml volumetric flask and add distilled water up to exactly 500 ml. Seal it tightly with the cork. Label the solution as sodium thiosulfate. Keep the flask closed when not in use.

Preparation of standard iodate solution

8. Use a 100 ml measuring cylinder to fill 450 ml of distilled water into a 500 ml beaker.
9. Weigh 1,01 g of KIO₃ in a small, clean cup using the scale and a clean spoon.
10. Pour the KIO₃ into the 450 ml of distilled water. Swirl the beaker until the potassium iodate has completely dissolved.
11. Once dissolved, pour the solution into a clean 500 ml volumetric flask and add distilled water up to exactly 500 ml. Seal it tightly with the cork. Label the solution as potassium iodate. Keep the flask closed when not in use.

Standardising the sodium thiosulfate solution

12. Set up the stand and clamp for the 50 ml burette.
13. Fill the closed 50 ml burette with the previously prepared sodium thiosulfate solution using a clean funnel. It is vital that the burette contains precisely 50 ml.
14. Using a 50 ml measuring cylinder pour exactly 50 ml of the KIO₃ solution into a clean 250 ml Erlenmeyer flask.
15. Weigh 2 g of KI in a small cup using the scale and a spoon. Place the KI into the flask.
16. Add 5 ml of 8 M H₂SO₄ into the flask using a 50 ml measuring cylinder.
17. Place the 250 ml Erlenmeyer flask onto a magnetic stirrer and begin stirring it. This is to ensure that all the reactants in the solution have reacted to form the triiodide molecule.
18. The solution should have a dark red colour due to the presence of triiodide. Titrate the solution with sodium thiosulfate until the solution has lost most of its red, i.e. a light shade of yellow appears. Using the 3 ml measuring pipette, add 2 ml of the starch indicator to the solution. The starch is only added shortly before the end point of the titration as prior to this, the triiodide starch complex locks onto the triiodide and thus the triiodide might not react with the sodium thiosulfate.
19. Continue titrating the solution until the solution has become colourless. Record the amount of ml of sodium thiosulfate solution used.
20. Repeat the titration three times in order to obtain reliable values, as this titration tells us the exact concentration of sodium thiosulfate, which allows us to determine the amount of triiodide.

Titration of ascorbic acid

21. Use a 50 ml measuring cylinder to fill 15 ml of 8 M H₂SO₄ into a clean 250 ml Erlenmeyer flask used for titration. Using a 100 ml measuring cylinder add 75 ml of distilled water.
22. Grind vitamin C tablets separately and thoroughly with a mortar and pestle. Put them into the flask and stir until fully dissolved. If needed, for example if part of the tablet is not soluble anymore, filter the solution by using a funnel and filter paper.
23. Carefully pour the solution into a clean 250 ml Erlenmeyer flask used for titration.
24. Using a 50 ml measuring cylinder pour 50 ml of KIO₃ solution into the 250 ml Erlenmeyer flask used for the titration.
25. Weigh 2 g of KI in a small cup using the scale and a spoon. Place the KI into the
    250 ml Erlenmeyer flask used for the titration. At this point, the solution should develop a dark shade of red due to the presence of triiodide.
26. Swirl the flask to make sure the reaction between triiodide and ascorbic acid has been completed.
27. Set up the stand and clamp for the 50 ml burette.
28. Place the 250 ml Erlenmeyer flask used for the titration onto a magnetic stirrer, and begin stirring it. This ensures that the vitamin C has truly completely reacted.
29. Fill the closed 50 ml burette with sodium thiosulfate solution using a clean funnel. Ensure that the burette was cleaned beforehand and remove any excess solution.
30. Begin titrating the triiodide solution with sodium thiosulfate. It should start out being red due to the presence of access triiodide. When the solution changes to a pale yellow, add 2 ml of the starch indicator using a 3 ml measuring pipette. Similar to the previous titration, the starch might hold onto the triiodide and prevent it from reacting with sodium thiosulfate. Continue titrating the solution until it has become colourless. Due to various colourings of the tablets, this might be an off-shade of white.
31. Denote this volume as the end point of the titration.
32. Repeat steps 21-31 for all tablets and sachets available. Each sample of tablets or sachets should be titrated at least five times in order to calculate a representative amount of vitamin C contained in the tablets. It is important to note that the flask containing the vitamin C solution and the magnet of the magnetic stirrer must be washed before each trial to avoid impurities.  (CH 212 – Quantitative Analysis. 5-7)

Safety Considerations

Make sure that safety glasses and gloves are used during the experiment. Appropriate safety clothes must be worn, like laboratory coats with long sleeves. Content of solutions prepared need to be clearly marked with water proof pencils and locked away in laboratory cupboards. Any solution not used anymore needs to be placed it appropriate waste disposal units. Neutralise any acids before disposing of them.
Raw Data
Table 1 shows the overview of all the samples used in the iodometric redox titrations.

Table 1: Samples used in the titration
Table 2 shows the volume of sodium thiosulfate required to titrate 50 ml of potassium iodate.

Table 2: Volume of sodium thiosulfate solution

Note: As the percentage uncertainty of titration 1 0,25% and titration 1 and 2 is 0,26%, these uncertainty are not taken into account, as there is very little impact on the results.

Table 3 shows the results of all titrations of the vitamin C samples.

Table 3: Volume of sodium thiosulfate solution in ml used in the titrations

It was observed that the colour changed from dark blue to a pale yellow as the sample was titrated due to the reaction of triiodide with thiosulfate.

Calculations and Processed Results

In order to find the percentage of vitamin C in each sample, one must calculate the exact concentration of sodium thiosulfate used in all titrations.

8,7 g Na₂S₂O₃ ï‚· 5H₂O with molar mass 248,2 g/mol in 500 ml water = 0,070109 mol/l S₂O₃²

Molar ratio of the oxidation of S₂O₃^2- is S₂O₃^2-:I₃^– = 2:1

38,73 ml of 0,070109 mol/l S₂O₃^2- = 0,002715 mol S₂O₃^2- are oxidized by 0,001358 mol I₃

Molar ratio of the creation of I₃^– is IO₃^–: I₃^– = 1:3

0,001358 mol I₃^–  = 0,000453 mol IO₃^– in 50 ml = 0,009051 mol/l IO₃^–

Exact concentration of KIO₃ = 1,01 g KIO₃/500 ml = 0,009439 mol/l IO₃^–

∴ Exact concentration of S₂O₃^2- is higher by a factor of 0,009439/0,009051 = 1,042911
∴ Exact concentration = 0,070109 ∙ 1,042911 = 0,073117 mol/l S₂O₃^2-

Table 4 shows the volume and concentration of sodium thiosulfate used.

Table 4: Calculation of concentration of sodium thiosulfate

Note: As the uncertainty of 38,73 ml ± 0,3 gives a percentage uncertainty of 0,77%, this uncertainty are not taken into account, as there is very little impact on the results.

In knowing both the precise concentration of sodium thiosulfate and the volume needed to titrate the remaining triiodide, one can determine the concentration of vitamin C in each sample as there is an excess of KIO₃ present in each titration.

Reaction 2: C₆H₈O_{6 (aq)} + I₃^– _(aq)  C₆H₆O_{6 (aq)} + 3 I^–_(aq) + 2 H⁺_(aq)

Molar ratio of I₃^–:C₆H₈O₆ = 1:1
            Molar ratio of IO₃^–:I₃^– = 1:3
            Molar ratio of IO₃^–:C₆H₈O₆ = 1:3

∴ 50 ml 0,009439 mol/l KIO₃ = 0,000472 mol IO₃^– = 3 times more moles of vitamin C
∴ 0,000472 ∙ 3 = 0,001416 mol vitamin C = 249,4 mg vitamin C

Only the first sample had a theoretical maximum content of 500 mg, which is more than the 50 ml of 0,009439 mol/l IO₃^– solution can oxidize, that means no triiodide should have been left over and therefore the blue starch triiodide complex should not have been formed. However, in each titration a blue colour was visible. Therefore, the method applied was valid for all other titrations, otherwise the weight of the sample should have been reduced.

Reaction 3:I₃^– _(aq) + 2 S₂O₃^2- _(aq)  3 I^– _(aq) + S₄O₆^2- _(aq)

Molar ratio of I₃^–: S₂O₃^2- = 1:2

The arithmetic average of 10 titrations of the first sample is 3,64 ml of 0,073117 mol/l S₂O₃^2- used to titrate the remaining I₃^–. 3,64 ml of 0,073117 mol/l S₂O₃^2- = 0,000266 mol S₂O₃^2- are oxidized by 0,000133 mol I₃^–. After the reaction of 50 ml of 0,009439 mol/l IO₃^– solution with the vitamin C sample 0,000133 mol I₃^– were left over.

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Molar ratio of IO₃^–:I₃^– = 1:3
50 ml 0,009439 mol/l IO₃^–= 0,000472 mol IO₃^– = 0,001416 mol I₃^–
0,001416 mol I₃^– at the start of the reaction – 0,000133 mol I₃^– found after the reaction = 0,001283 mol reacted with vitamin C

Molar ratio of I₃^–:C₆H₈O₆ = 1:1
0,001283 mol I₃^– = 0,001283 mol C₆H₈O₆ = 0,001283 mol ∙ 176,1 g/mol = 225,9 mg vitamin C
500 mg vitamin C = 100% and 225,9 mg vitamin C = 45,19% active content.

This method is applied to all results. Table 5 shows the results of titrating the four different vitamin C samples.

Table 5: Calculated and measured results of titration of Vitamin C samples
Months expired is the time elapsed between the expiry date and the month of the laboratory work in September 2016. The first titration of sample 1 (see table 3) was an outlier and not considered when calculating the arithmetic average and standard deviation. Only one sachet of sample 2 was available, thus not allowing to calculate average and standard deviation.
Graph 1 shows the percentage of Vitamin C found and the age of the tablets and sachet used.Graph 1:Correlation of concentration of Vitamin C versus number of months expired

Graph 2 shows that the reliabilty of the result depends on the number of titrations carried out. The variation of the results is two times standard deviation (± 2 σ).

Graph 2: Concentration of vitamin C and error bar of ± 2 σ (sample 2 was only titrated once, thus no standard deviation can be calculated)

Assuming that the results are following a normal distribution, 95,4% of the titrations are within the ± 2 σ range as shown above. Thus, the results are accepted for sample 1, 3 and 4.

Conclusion and Evaluation
The results show that the vitamin C content decreases over time. This is demonstrated in graph 1, as the slope is -0,048. The coefficient of determination is 0,989, which is close to 1 and therefore shows a strong linear correlation. This means the results clearly demonstrate that as the number of months expired increases, the concentration of vitamin C decreases. These findings are supported by research of the Applied Sciences Department at the Osun State Polytechnic in Iree, Nigeria, published in 2012. (Oyetade 22) High temperature, exposure to air and sunlight accelerate the oxidation of vitamin C. Thus, the oldest sample shows the highest reduction in percentage concentration of vitamin C with 45,19%, less than half. This means that my father should buy new vitamin supplements, and no longer use his old ones, which were sample 1.
Strengths: The method of investigation delivers fast results, does not require expensive equipment, and works with chemicals that are neither very toxic nor extremely harmful to the environment. A clear relationship between the variables was demonstrated and due to relatively low standard deviation, the results are accepted.
Weaknesses: Potential errors were detected. Only one sachet of sample 2 was available, thus not meeting the minimum criteria of at least 3 titrations of each sample. It is also possible that not all the triiodide reacted with the sodium thiosulfate (see step 18 in methodology). In terms of method, the equipment used was not precise enough, causing possible systematic errors. Instead of a measuring cylinder a volumetric pipette should have been used as well as a high precision scale ± 0,1 mg. Not all samples were easily soluble in water and produced a clear and colourless solution, thus the end of titration was difficult to notice, leading to inaccuracies. The thiosulfate and iodate solutions should not have been stored over a long time exposed to uncontrolled temperature and day light. All titrations should either have been done in one day or the titrants should have been stored in a cool and dark place.

Works Cited

CH 212 – Quantitative Analysis. Philadelphia: La Salle University, n.d. PDF

Oyetade, O. A., G. O. Oyeleke, B. M. Adegoke, and A. O. Akintunde. Stability Studies on Ascorbic Acid (Vitamin C) From Different Sources. N.p.: IOSR Journal of Applied Chemistry (IOSR-JAC), Sept.-Oct. 2012. PDF.