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CHL 212 – Quantitative Analysis
Experiment 7 Redox Titration of Vitamin C Introduction
In this experiment, you will be acting as the quality control laboratory for a pharmaceutical manufacturer. The product line that you support produces 100-mg Vitamin C supplements. You are to determine the average amount of Vitamin C per tablet in a sample of tablets and report this value and its uncertainty to the product line manager.
As in previous experiments, you will determine the amount of Vitamin C in the tablets using titration. In the case of Acid/Base titration, the acidic or basic sample was neutralized with the appropriate titrant and the amount of sample was determined by measuring the volume of titrant required to reach the end point. In the determination of dissolved Ca 2+ and Mg 2+
experiment, you measured the amount of calcium and magnesium dissolved in the laboratory tap water by using a titrant (EDTA) that formed a complex with the ions to be determined.
Although new methods and reagents are being discovered and developed continuously, often a titrant solution is not available that reacts directly with the species of interest. In this experiment, we do not have a convenient solution that reacts or complexes directly with the Vitamin C to be analyzed. In these types of cases, we must perform our analysis indirectly by analyzing species that do react or complex with the species of interest.
Vitamin C (Ascorbic Acid), is a mild reducing agent (it accepts electrons from an electron donor, leaving the Oxidation State of the donor at a value less than original (reduced)). The Ascorbic Acid itself is oxidized to a higher oxidation state. This class of reactions is known as a reduction/oxidation reaction or simply, a redox reaction. One such redox reaction is the reduction of the aqueous iodine molecule (I2(aq)) with Ascorbic Acid, as shown below.
(1) KIO 3 (aq) + 6 H
(aq) + 5 I- (aq) →3 I 2 (aq) + 3 H 2 O (l) + K
(aq) generation of I 2
(2) C 6 H 8 O 6 (aq) + I 2 (aq) →C 6 H 6 O 6 (aq) + 2 I- (aq) + 2 H
(aq) oxidation of vitamin-C
Reaction one generates aqueous iodine, I 2 (aq). This is then used to oxidize vitamin-C (ascorbic
acid, C 6 H 8 O 6 ) in reaction two. Both of these reactions require acidic conditions and so dilute
hydrochloric acid, HCl (aq) , will be added to the reaction mixture. Reaction one also requires a source of dissolved iodide ions, I⎯^ (aq). This will be provided by adding solid potassium iodide, KI (s) , to the reaction mixture.
The two relevant half reactions for reaction (2) above are:
I 2
Reduction half reaction for Iodine at pH 5
Oxidation half reaction for vitamin-C (C 6 H 8 O 6 ) at pH 5
A few drops of starch solution will be added to help determine the titration endpoint. When the vitamin-C (ascorbic acid) is completely oxidized, the iodine, I 2 (aq) , will begin to build up and will
react with the iodide ions, I- (aq) , already present to form a highly colored blue I 3 --starch
complex, indicating the endpoint of our titration.
Each of the iodine atoms is reduced to the I
- ion and the Ascorbic Acid gains two electrons to form Dehydroascorbic Acid as in the chemical reaction below:
Ascorbic Acid + I2(aq) + H 2 O Dehydroascorbic Acid + 2I
This reaction has a high K value and goes to completion. Unfortunately, the solubility of I2(solid) in water is not very high. The saturated I2(aq) concentration is only 1.3 mM.
I2(solid) I2(aq) K = 1.3 x 10
To increase the solubility of the I 2 molecule, we can create a complex between the I2(aq) and the iodide ion (I
I2(aq) + I
I 3
The I 3
- complex is known as triiodide. Rather than starting with solid I 2 and taking the time to dissolve it in solution, triiodide can also be produced by reacting iodate ion (IO 3
) with iodide (I
) as below:
IO 3
8I
+ 6H + 3I 3 - + 3H 2 O
Triiodide is reduced by Ascorbic acid in the same way that the I2(aq) species reacted:
Ascorbic Acid + I 3
H 2 O Dehydroascorbic Acid + 3I
+ 2H +
It is this reaction that we use to measure indirectly the amount of Ascorbic Acid in the sample. We do have a titration solution and indicator that can measure the amount of I 3
- in the sample. The titration solution is sodium thiosulfate (thigh-o-sul-fate), Na 2 S 2 O 3 , and it reacts with the triiodide species in the 1:2 reaction below:
I 3
- 2S 2 O 3 2- 3I - + S 4 O 6 2-
The indicator used is a starch solution. In the presence of the triiodide, starch and triiodide form a complex that is intensely dark blue in color. In the absence of triiodide, the starch indicator is a milky-white.
THE PLAN
So here goes…
The starch indicator is readily biodegraded, so we will need a fresh solution of starch indicator at the start of each lab period.
We will make a standard solution of IO 3
- by weighing out a known amount of KIO 3 and dissolving it in a volumetric flask having a known volume to obtain the known molarity of the IO 3 solution.
Then we will make a solution of sodium thiosulfate that is around 0.07 M. We need to buffer the pH of this solution so we will also add some Na 2 CO3.
We will then take a known amount of IO 3
- and add a large (excess) amount of I - to produce a known amount of triiodide (I 3
) – note the 1:3 stoichiometry. We will titrate this known amount of triiodide with the sodium thiosulfate solution using the starch indicator (1:2 reaction). This will permit us to find the exact concentration of the sodium thiosulfate solution. With this known concentration, we can then determine the amount of triiodide in a sample.
Then (FINALLY!) we will dissolve some Vitamin C tablets. We will then add a known amount of the standard IO 3
- solution and an excess of I - to produce a known total amount of triiodide. The Ascorbic Acid will react with some of the triiodide, resulting in a small amount of triiodide remaining in the sample. We will use the sodium thiosulfate solution and the starch indicator to determine the amount of triiodide remaining (1: reaction).
With the knowledge of how much triiodide we started with and the amount of triiodide remaining after reaction with all of the ascorbic acid (1:1), we can determine the amount of ascorbic acid contained in the tablets.
The following indicator will be used for the titration:
Starch solution
dark blue (complexed with iodine) milky-white (uncomplexed) end point is disappearance of dark blue
Procedure
The procedure for this experiment has been adapted from experiment 13 on the following website: www.whfreeman.com/qca, “Iodometric Titration of Vitamin C”.
SAFETY NOTE
The sulfuric acid used in this experiment WILL CAUSE INJURY if in contact with the skin. The proper use of laboratory SAFETY GLASSES and/or SPLASH GOGGLES and GLOVES will be expected and strictly monitored.
Reagents
Starch Solution Sodium Carbonate (Na 2 CO 3 ) 0.5 M Sulfuric Acid (H 2 SO 4 ) 0.3 M Sulfuric Acid (H 2 SO 4 ) Potassium Iodide (KI) Potassium Iodate (KIO 3 ) 100 mg Vitamin C Tablets Sodium Thiosulfate pentahydrate (Na 2 S 2 O 3 5H 2 O)
Standardization of Sodium Thiosulfate Solution
Preparation of Starch Indicator
- Clean and rinse a 100-mL beaker with DI water. Fill the beaker to the 100-mL mark and place on a hot plate until boiling.
- Weigh out around 1 gram of soluble starch and slowly add to the beaker of boiling water with stirring. Continue boiling until the solution is clear.
Preparation of Sodium Thiosulfate Solution
- Clean and rinse a 500-mL beaker with DI water. Fill the beaker to the 500-mL mark and place on a hot plate. Boil the water for 5 minutes to expel dissolved CO 2 gas. Allow solution to cool.
- Weigh out around 0.05 g of Na 2 CO 3 and place into the 500-mL beaker of boiled water.
- Weigh out around 8.7 grams of Na 2 S 2 O 3 5H 2 O and dissolve into the 500-mL beaker of boiled water buffered with Na 2 CO 3.
- Clean and rinse a 1-L brown plastic bottle with DI water. Rinse the bottle with a small amount of the sodium thiosulfate solution. Transfer the remainder of the sodium thiosulfate solution to the brown plastic bottle. Label this bottle “Sodium Thiosulfate Solution” and keep this bottle tightly capped when not in use.
Analysis of Vitamin C
React Triiodide with Ascorbic Acid
- Clean and rinse a 250-mL flask with DI water.
- Place 60 mL of 0.3 M H 2 SO 4 into the clean flask.
- Add 2 Vitamin C tablets to the flask and dissolve in the sulfuric acid. You can use a clean glass stirring rod to help break up the solid tablets. Some solid binding material may not dissolve.
- Accurately pipet 50.00 mL of standard KIO 3 solution into the flask.
- Weigh out around 2 grams of KI and place into the flask.
- Gently swirl the flask for 1 minute to insure complete formation of the triiodide complex and complete reaction between the triiodide and the Vitamin C.
Titrate Solution with Sodium Thiosulfate
- The solution should start out as a “red-ish” solution due to the presence of the triiodide. Titrate the solution with the Sodium Thiosulfate until the solution has lost most of the red-ish color (should be a pale yellow). At this point add 2 mL of the starch indicator (may turn smoky-blue or remain yellowish). Carefully add Sodium Thiosulfate until the solution turns colorless (may be a milky-white). Record this volume as the end point.
The indicator is not added until just before the end point as the triiodide/starch complex can “hold on to” the triiodide in concentrated solutions and prevent it from reacting with the Sodium Thiosulfate.
- Repeat this titration (steps 20 – 26) two more times to obtain 3 good values for the volume of Sodium Thiosulfate required to titrate the remaining triiodide.
Q (90%) Number of Observations 0.94 3 0.76 4 0.64 5
range
gap Qcalc
Data Analysis
- Use your buret calibration graph created in Experiment I to correct your final buret readings.
Using the formula weight of KIO 3 , calculate the molarity of your standard IO 3
solution.Using the known stoichiometry of the Sodium Thiosulfate/IO 3
reaction, calculate the
molarity of your Sodium Thiosulfate solution. Calculate the average and standard deviation of the concentration.Using the known stoichiometries of the Sodium Thiosulfate/IO 3
and Ascorbic
Acid/IO 3 - reactions, calculate the average number of moles of Ascorbic Acid contained in each Vitamin C tablet.- Using the formula weight of Vitamin C (C 6 H 8 O 6 , FW = 176.13), calculate the average number of milligrams of Vitamin C contained in each tablet along with the standard deviation.
Report
Report the concentration of your standard IO 3
Report the average concentration with both absolute error and % relative error of your standardized Sodium Thiosulfate solution.
Report the calculated average concentration with both absolute error and % percent relative error of the determined mass in milligrams of Vitamin C per tablet.
Include the manufacture’s brand and label information from the bottle of Vitamin C tablets.
Question: Why is sulfuric acid used in the titrations?
