Chemistry 12
Determining the Amount of Vitamin C (Ascorbic Acid)
in a Sample
Abstract
Redox reactions are involved in a wide variety of techniques for quantitative analysis of
chemical substances. The methods for calculating the amounts of substances involved in
these reactions (expressed either as mass in grams or as the volume of a solution of
known molarity) will have been introduced to in your studies of stoichiometry. The
balanced equations for many of these redox reactions are complex and must be obtained
either by adding balanced half-reaction equations or by using the oxidation number
method.
A substance that is often used in quantitative redox reactions is iodine, I2. It is easily
formed by oxidizing I-, but the reaction can be reversed just as easily, so a wide variety of
reactions can involve I2 as a reactant or product. In addition, it gives a characteristic deep
blue color with starch solution even when it is in very low concentrations, making it easy
to determine when the I2 is all used up or when it begins to form. Starch is therefore used
as the indicator in these reactions.
In Part I of this experiment, you will prepare a solution of potassium iodate (KIO3) of
known concentration. This solution makes a good primary standard in redox titrations,
since it is stable and can be obtained very pure. In Part II, known volumes of this solution
will have excess H+ and I- added and a reaction will occur in which I2 is produced in a
quantity determined by the moles of KIO3 present initially. A solution of sodium
thiosulfate, Na2S2O3, of unknown concentration will then be titrated into the solution,
which will react with the I2. Starch will be added when most of the I2 has reacted with
S2O32- and the titration will be continued until the blue color disappears. The
concentration of the NaS2O3 solution can now be calculated. This solution can be used to
determine the amount of I2 in other solutions. In Part III, the amount of vitamin C in a
sample will be determined by analyzing it with a standard I2 solution.
Purpose
1. To prepare a solution of potassium iodate, KIO3, of known molarity.
2. To standardize a solution of sodium thiosulphate, Na2S2O3, by titrating it against
standard KIO3 (with excess I-, H+).
3. To use a standard I2 solution to determine the mass of ascorbic acid (vitamin C) in
a vitamin tablet.
Chemistry 12
-1-
Materials
Apparatus
2 x 250 mL beaker
Scale
Pipet bulb
250 mL Erlenmeyer flask
Funnel
250 mL volumetric flask
Stirring rod
Teflon coated magnet
4 x 100 mL beaker
25 mL pipet
Wash bottle
Mortar and pestle
10 mL graduated cylinder
1 L Nalgene reagent bottle
Buret
Magnetic stirrer
Reagents
Potassium iodate (KIO3(s))
1.0 M Potassium iodide (KI(aq))
Starch solution
0.12 M Sodium thiosulphate (Na2S2O3(aq))
1.0 M Sulfuric acid (H2SO4(aq))
Procedure
Part 1 – Preparing a Standard Potassium Iodate Solution
1. Before coming to the laboratory, calculate the mass of KIO3 required to make up
250.0 mL of a 0.0200 M solution of KIO3.
2. Put on your lab apron and safety goggles.
3. Accurately measure the mass of a clean, dry 100.0 mL beaker and record it in
your copy of Table 1 in your notebook.
4. Place solid KIO3 in the beaker until you have approximately the mass calculated
in step 1. Record the actual mass accurately in Table 1.
5. Dissolve the KIO3 in water and pour the solution through a funnel into a 250.0
mL volumetric flask. Rinse the beaker three times and add the rinsings to the
flask.
6. Add water until the level of the solution is up to the mark, stopper the flask, and
invert and swirl thoroughly to make the solution homogeneous.
Part 2 – Standardizing a Sodium Thiosulphate Solution
1. Transfer the standard KIO3 solution to a clean, dry 250.0 mL beaker.
2. Obtain in three 100.0 mL beakers approximately 50.0 mL each of 1.0 M KI,
1.0 M H2SO4, and starch solution. Label each beaker.
-2-
3. Using a suction bulb on your pipet, withdraw 25.00 mL of the KIO3 solution and
transfer it to a 250.0 mL Erlenmeyer flask, touching the tip to the side of the flask
to ensure that the correct volume is delivered.
4. Add to the flask containing the KIO3 solution approximately 5 mL of 1.0 M KI
and 5 mL of 1.0 M H2SO4. These are excess amounts, so precise measurement is
not required. Use a 10 mL graduated cylinder for adding the KI and H2SO4. You
will notice that a brown precipitate is formed at first and that it dissolves again to
give a clear brown solution. This behavior is characteristic of iodine, I2. The
amount of precipitate formed is determined by the moles of KIO3 originally
present. The iodine will now be made to react with the sodium thiosulphate
solution, Na2S2O3, in order to determine the molarity of the Na2S2O3.
5. Fill a 1 L nalgene reagent bottle about half full with the 0.12 M Na2S2O3 solution
provided and label it with your name. Use solution from this bottle for all
subsequent parts of the experiment requiring Na2S2O3.
6. Add about 10 mL of this solution the buret, rinse, and discard. Fill the buret with
Na2S2O3, then open the valve briefly to allow some to drain through the tip.
7. Determine the initial buret reading and record it in Table 2.
8. Run the Na2S2O3 solution into the flask until the brown colour has faded to a light
yellow. Add approximately 5 mL of starch solution and continue adding Na2S2O3
drop by drop until the blue/black colour disappears. Note that the endpoint is a
very precise one and that it is therefore important not to add the starch too soon.
Take the reading on the buret and record it in Table 2.
9. Repeat steps 3 to 8 until you obtain consistent results (titrations agree within 0.10
mL). You now have enough data to calculate the molarity of the Na2S2O3
solution.
Part 3 – Determining the Amount of Vitamin C (Ascorbic Acid) in a Sample
1. Obtain in a 250.0 mL beaker about 200.0 mL of the of the solution of iodine (in
Potassium iodine, I2/KI) provided. You will need to standardize it by completing
steps 2 to 5 before doing the rest of Part III.
2. Using a suction bulb, withdraw 25.00 mL of the iodine solution into a pipet and
transfer it to a 250 mL Erlenmeyer flask.
3. Refill the buret with the standardized Na2S2O3 solution, take the initial reading,
and record it in Table 4.
4. Run the Na2S2O3 solution into the flask until the brown colour of the I2 fades to
pale yellow, add 5 mL of starch solution, and continue the titration dropwise until
the dark blue colour just disappears. Again, take the final reading on the buret and
record it in Table 3.
5. Repeat steps 2 to 4 until consistent results are obtained (titrations agree within
0.10 mL). You now have enough data to calculate the molarity of the I2 solution
when you later perform your analysis of results.
-3-
6. Discard the Na2S2O3 solution remaining in the buret, wash the buret with water,
then rinse it with the I2 solution. Discard this, then fill the buret again with I2, and
drain some to refill the tip.
7. Obtain a 250 mg (or 300 mg) tablet of vitamin C (ascorbic acid), grind it with a
mortar and pestle, and place the powder in a 250 mL Erlenmeyer flask. Add 5 mL
of starch solution.
8. Take the initial reading of the I2 solution in the buret and record it in Table 4.
Then allow it to run into the flask. Swirl constantly, watching for the first
appearance of the characteristic dark blue colour, which stays even after swirling.
Again, take the final reading on the buret and record it in Table 4.
9. Repeat steps 7 and 8 until consistent results are obtained (titrations agree within
0.10 mL).
10. Clean all glassware and return to their original locations.
11. Before leaving the laboratory, wash your hands thoroughly with soap and water.
Observations
Part 1 – Preparing a Standard Potassium Iodate Solution
Table 1: Preparing a Standard KIO3 Solution
Calculated mass of Potassium iodate, KIO3, required for
250.0 mL of 0.0200 M solution (g).
Mass of beaker (g)
Mass of beaker + KIO3 (g)
Mass of KIO3 (g)
Part 2 – Standardizing a Sodium Thiosulphate Solution
Table 2: Volume of Na2S2O3 Needed to React with Iodine from 25 mL of KIO3
Trial 1
Trial 2
Trial 3
Trial 4
(if necessary)
Initial reading of Buret
(mL)
Final reading of Buret
(mL)
Volume of Na2S2O3
required (mL)
Average volume of
Na2S2O4 (mL)
-4-
Part 3 – Determining the Amount of Vitamin C (Ascorbic Acid) in a Sample
Table 3: Volume of Na2S2O3 Needed to React with 25 mL of I2 Solution
Trial 1
Trial 2
Trial 3
Trial 4
(if necessary)
Initial reading of Buret
(mL)
Final reading of Buret
(mL)
Volume of Na2S2O3
required (mL)
Average volume of
Na2S2O4 (mL)
Table 4: Volume of I2 Needed to React with Ascorbic Acid (Vitamin C)
Mass of Vitamin C in Tablet = _________ mg
Trial 1
Trial 2
Trial 3
Trial 4
(if necessary)
Initial reading of Buret
(mL)
Final reading of Buret
(mL)
Volume of I2 required
(mL)
Average volume of I2
(mL)
Post Lab Considerations
In order to calculate the results for this experiment, you will need to work out the
balanced overall redox equation for each reaction that is occurring, then use the mole
relationships in the equation to relate reacting quantities The balanced half-reactions
occurring in this experiment are as follows:
Production of I2 from KIO3, KI, H2SO4:
2 IO3- + 12 H+ + 10 e- → I2 + 6 H2O
2 I- → I2 + 2e-
Standardization of Na2S2O3:
2 S2O32- → S4O62- + 2 e-
You need not show an equation for the reaction of vitamin C (ascorbic acid, C6H8O6)
with I2. They react in a 1:1 mole ratio.
The deep blue/black color formed by starch and iodine is very useful in a variety of
situations. Its formation is not really a chemical reaction. The iodine molecules just
happen to fit closely inside the long spiraling starch molecule and interact strongly with
-5-
it. The combination is unstable above 50°C. It is interesting to note that here you are
using starch to test for the presence of iodine, but in food chemistry the reverse occurs,
iodine solution is used to test for the presence of starch. You may have performed such a
test in an earlier science course.
Note that iodine is only slightly soluble in water, but does dissolve well if I- ions are
present. The reaction that occurs is I2 + I- → I3-. However, the complex I3- breaks down
very readily when any chemical is present that can react with I2 (such as S2O3-) and the
solution therefore acts as though it were a solution of I2.
Questions
Part 1 – Preparing a Standard Potassium Iodate Solution
1. Calculate the concentration of KIO3 solution formed when you dissolved your
calculated mass of KIO3 in water and made the volume up to 250.0 mL.
2. Calculate the number of moles of KIO3, and therefore of IO3-, in 250.0 mL of this
solution.
Part 2 – Standardizing a Sodium Thiosulphate Solution
1. Work out the overall redox equation for the reaction of IO3-, I-, and H+ to give I2.
2. Work out the overall redox equation for the reaction of S2O32- and I2 to give
S4O62- (tetrathionate ion) and I-.
3. State the relationship between moles of IO3- and moles of S2O32- working through
moles of I2.
4. From this relationship and the number of moles of KIO3 in 25.00 mL of solution,
calculate the number of moles of S2O32- with which the KIO3 reacts.
5. Knowing the average volume and the number of moles of S2O32- used, calculate
the [S2O32-].
Part 3 – Determining the Amount of Vitamin C (Ascorbic Acid) in a Sample
1. Calculate the number of moles of S2O32- used from the average volume used and
the molarity from Part II.
2. Calculate the number of moles of I2 present, using the relationship obtained in
Part II.
3. Calculate [I2] from the number of moles used and the volume in litres.
4. From the average volume of I2 solution used in the vitamin C titration, calculate
the number of moles of I2 required to react with the vitamin C.
-6-
5. Recalling that it was stated in the Post Lab Considerations that vitamin C and I2
react in a 1:1 mole ratio and that vitamin C (ascorbic acid) has the formula
C6H8O6, calculate the mass in milligrams of vitamin C present.
6. Compare your result with the rated amount of vitamin C for that tablet and
calculate the percentage deviation between your result and the rated amount.
Conclusions
State the mass of vitamin C present in the tablet and the percentage deviation from the
given value on the bottle.
-7-