Chemistry 1156
Chemistry of Vitamin C
Purpose
In this experiment, you will observe the oxidation-reduction
reaction that takes place between vitamin C (ascorbic acid,
C6H8O6) and iodine, wherein vitamin C is the reducing agent
and iodine is the oxidizing agent. One objective of this
experiment is to use the quantitative redox titration to
determine the amount of ascorbic acid in a commercial
vitamin C tablet.
Introduction
While the importance of vitamin C in preventing scurvy was
recognized as early as 1795 by the British Navy, it was not
until 1928 that ascorbic acid was isolated and identified as the
vital anti-scurvy factor. Since that time, vitamin C
consumption has also been attributed to prevention of the
common cold.
Vitamin C is one of the water-soluble vitamins, capable of
significant hydrogen bonding with water. Because it is watersoluble, vitamin C is continually excreted from the body,
unlike the fat-soluble (nonpolar) vitamins which are stored in
the body’s fat deposits. As a result, the body’s supply of
vitamin C must continually be replenished through ingestion
of foods high in vitamin C and vitamin supplements. Since
excess vitamin C is rapidly excreted in the urine, there is little
risk in consuming more vitamin C than required by the body.
This experiment takes advantage of ascorbic acid’s watersolubility and strong reducing ability to quantitatively
determine the amount of ascorbic acid present in a
commercially available vitamin C tablet. In particular,
ascorbic acid’s ability to reduce iodine to the iodide ion will
be exploited to determine the amount of vitamin C present in
the sample. Following the procedure of R. W. Ramette (R. W.
Ramette, Chemical Equilibrium and Analysis, 1981), iodine will
be produced in solution through the reaction of iodide ion
with iodate ion in acidic solution. To this end, our experiment
combines solid potassium iodide (KI), hydrochloric acid, and
starch solution with the solution of ascorbic acid. This
solution is titrated with potassium iodate, (KIO3), which reacts
with the iodide ion as follows:
IO3–(aq) + 5 I–(aq) + 6 H+(aq)  3 I2 + 3 H2O
Eqn. 1
As quickly as the iodine is produced, it is reduced back to
iodide ion by the ascorbic acid (equation 2, below).
When all of the ascorbic acid has been oxidized, the iodine
concentration will increase to the point where it will react
with iodide ion to form the triiodide ion, I3–
I2 + I–  I3–
Eqn. 3
The triiodide ion in turn combines with the starch to form the
dark blue triiodide-starch complex. The appearance of this
dark blue color signals the endpoint of the titration.
MATERIALS
Supplies: None
Reagents: vitamin C tablets (cut in half), 0.03 M KIO3, solid KI,
1.0% starch solution, 1.0 M HCl (diluted from stock solution)
1|Page
General Chemistry II Lab for Majors
Chemistry of Vitamin C
Procedure
1. Prepare 50-mL buret for redox titration of vitamin C. Rinse
the buret 2–3 times with distilled water, making sure that the
stop-cock and tip are both free of obstructions. Rinse the buret
twice with 10–15 mL portions of 0.03 M KIO3 solution. After
the buret has been rinsed thoroughly, fill the buret with the
0.03 M KIO3 solution. Read the initial volume of KIO3 solution
to the nearest 0.01 mL and record this initial volume in your
laboratory notebook.
2. Prepare vitamin C solution. Weigh ½ of a vitamin C tablet.
Place this half tablet in a 125 mL Erlenmeyer flask and
carefully crush the tablet with a glass stirring rod. Once the
tablet is crushed, dissolve it in 50 mL of distilled water. Any
undissolved material that remains after 2–3 minutes is inert
binder material that does not need to dissolve.
Rinse the stirring rod with a small amount of distilled
water as you remove it from the vitamin C solution.
To the vitamin C solution, add 1 gram of solid potassium
iodide and 5 mL of 1.0 M HCl. Swirl the flask until all of the
potassium iodide dissolves.
Finally, add 1–2 drops of 1% starch solution.
3. Titration. Titrate the vitamin C solution with the potassium
iodate (KIO3) solution until the dark blue color persists. As the
redox titration reaches completion, excess iodine (I2) reacts
with iodide ion (I–) to form the triiodide ion (I3–). Triiodide ion
in turn reacts with starch, forming the dark blue starchtriiodide ion complex, signaling the endpoint of the titration.
Record the final volume of titrant (in the buret) to the nearest
0.01 mL in your laboratory notebook.
Each lab partner should conduct one titration until your
values for the percentage of vitamin C per gram agree to
within 10% of each other. In other words, if your first two
titrations do not agree to within 10%, both students should
repeat the titration a second time, striving for better accuracy.
4. Disposal. When you are finished with the experiment, all of
the solutions may be flushed down the drain.
2|Page
Chemistry 1156
Chemistry of Vitamin C
Vitamin C Content of Commercial Tablets
Data
Trial 1
Trial 2
Mass of tablet
Trial 3
Trial 4
g
g
g
g
mL
mL
mL
mL
Initial buret reading
mL
mL
mL
mL
Net volume of 0.03 M KIO3 solution
mL
mL
mL
mL
Volume of 0.03 M KIO3
Final buret reading
Calculate the mass of vitamin C (ascorbic acid) present in each ½ tablet, taking into account the reaction stoichiometry shown in
equations 1 and 2. Show a sample calculation in the space provided below.
Data
Mass of vitamin C
Trial 1
mg
Trial 2
Trial 3
mg
Average mass
mg
Trial 4
mg
mg
Post-Lab Questions
1. Calculate the percent purity of ascorbic acid for each trial. Are the tablets pure
ascorbic acid, or is there evidence that they contain some inert material? Record your
observations made when dissolving the tablets, and compare the mass of the ½ tablet
with the mass of vitamin C found in the tablet.
mass C6 H 8O6
100  % purity
mass of tablet
2. Sodium nitrite is used as a meat preservative. The amount of NO2– in a sample can be determined by acidifying the solution to
form nitrous acid, HNO2, allowing the nitrous acid to react with an excess of iodide ion, and then titrating the triiodide ion (I 3–)
in the resulting solution with thiosulfate solution in the presence of a starch indicator. The balanced equations are:
2 H+ + 2 HNO2 + 3 I– → 2 NO + I3– + 2 H2O
I3– + 2 S2O32– → 3 I– + S4O62–
When a nitrite containing sample with a mass of 3.000 g was analyzed, 16.65 mL of 0.1500 M Na 2S2O3 solution was required to
completely titrate the sample. What is the mass percent of NaNO2 in the sample?
3|Page