Analysis of Vitamin C

Analysis of Vitamin C
Taylor Noeller
Noeller |1
Vitamins are a group of small molecular compounds that are essential nutrients
in many multi-cellular organisms, and humans in particular. The name "vitamin" is a
contraction of “vital amine”, and came about because many of the first vitamins to be
discovered were members of this class of organic compounds. And although many of
the subsequently discovered vitamins were not amines, the name was retained.
Fruits, vegetables, and organ meats (e.g., liver and kidney) are generally the
best sources of ascorbic acid; muscle meats and most seeds do not contain significant
amounts of ascorbic acid. The amount of ascorbic acid in plants varies greatly,
depending on such factors as the variety, weather, and maturity. But the most
significant determinant of vitamin C content in foods is how the food is stored and
prepared. Since vitamin C is easily oxidized, storage and the cooking in air leads to the
eventual oxidation of vitamin C by oxygen in the atmosphere. In addition, ascorbic
acid’s water-solubility means that a significant amount of vitamin C present in a food
can be lost by boiling it and then discarding the cooking water (Kramer et. al).
A deficiency of Vitamin C leads to the disease scurvy, at one time commonly
occurring during long sea voyages. British sailors combatted scurvy by carrying limes,
rich in Vitamin C, on their voyages, thus engendering the name “limey.” Although the
Food and Drug Administration recommends a daily intake of 60 mg of Vitamin C, Linus
Pauling suggested that amounts of 1-2 grams daily are instrumental in fighting the
common cold (Blake). Ascorbic acid is an important biological antioxidant (reducing
agent). It helps to keep the iron in the enzyme prolyl hydroxylase in the reduced form
and thereby it helps to maintain the activity of the enzyme. Prolyl oxidase, on the other
hand, is essential for the synthesis of normal collagen. In scurvy, abnormal collagen is
synthesized that causes skin lesions and broken blood vessels (Morasch).
This experiment illustrates how titration can be used to determine the ascorbic
acid content of a Vitamin C tablet containing about 500 mg of Vitamin C. The mass
percentage of ascorbic acid in Vitamin C is determined by titrating the Vitamin C
samples with standardized sodium hydroxide solution.
Noeller |2
In some acid-base neutralization reactions, an acid reacts with a metal
hydroxide base to produce water and a salt:
metal hydroxide
The protons (H+) from the acid react with the hydroxide ions (OH–) from the base
to form the water. The salt forms by combining the cation from the base and the anion
from the acid. Acids often react with bases; the solubility of the salt does not determine
whether the reaction occurs or not. When carrying out an acid-base neutralization
reaction in the laboratory, you observe that most acid solutions and base solutions are
colorless, and the resulting water and soluble salt solutions are also colorless. Thus, it is
impossible to determine when a reaction has occurred, let alone when it is complete.
If the appropriate indicator has been chosen, the endpoint of the titration (i.e., the
color change) will occur when the reaction is complete, or when the acid and base
are stoichiometrically equivalent (Reed):
moles of acid = moles of base
A Vitamin C tablet contains ascorbic acid, HC6H7O6 (aq), as well as binder
material that holds the tablet together. The balanced equation for the reaction
between ascorbic acid and sodium hydroxide is shown below:
HC6H7O6 (aq)
ascorbic acid
sodium hydroxide
NaC6H7O6 (aq)
sodium ascorbate
The objectives in this lab were:
Determine the experimental molar mass of ascorbic acid
Determine the amount (in mg) of ascorbic acid in a vitamin C tablet
Noeller |3
Experimental Procedure
1. Placed a preweighed vitamin C tablet in a preweighed Erlenmeyer flask
2. Added 50 mL of distilled water and let tablet soften for 10 minutes
3. Used stirring rod to break up tablet (rinsed stirring rod into flask)
4. Added 3 drops of phenolphthalein and magnet
5. Titrated to light pink colored solution
a. Placed a white sheet of paper under the flask
b. Turned on the magnetic stirrer to stir the acid at a steady rate
c. Slowly opened the stopcock and allowed NaOH to drip into flask
d. Stopped the addition of NaOH when the solution turned light pink
6. Recorded values and completed calculations
Data Table
mass of flask and tablet
mass of flask
mass of tablet
final buret reading
initial buret reading
mL of NaOH added
L of NaOH added
90.80 g
90.20 g
0.60 g
31.20 mL
0.10 mL
31.10 mL
.0311 L
Results- Objective 1
moles of NaOH added
L of NaOH x M of NaOH
moles of ascorbic acid neutralized
stoichiometry; 1:1 mole ratio
molar mass of ascorbic acid
g of ascorbic acid / mol of ascorbic acid
percent error
(experimental – true) / true x 100%
3.11 x 10-3 mol
3.11 x 10-3 mol
192.9 g/mol
Results- Objective 2
grams of ascorbic acid neutralized
mol of ascorbic acid x g/mol of ascorbic acid
milligrams of ascorbic acid neutralized
g ascorbic acid x 1000
percent error
(experimental – true) / true x 100%
0.547 g
547 mg
Noeller |4
mass of tablet
90.80 g – 90.20 g = 0.60 g
mL of NaOH added
31.20 mL – 0.10 mL = 31.10 mL
L of NaOH added
31.10 mL / 1000 L/mL = 0.0311 L
moles of NaOH added
0.0311 L x 0.10 M = 3.11 x 10-3 mol
molar mass of ascorbic
0.60 g / 3.11 x 10-3 mol = 192.9 g/mol
percent error (obj. 1)
(192.9 g – 176 g) / 176 g x 100% = 9.6%
grams of ascorbic acid
3.11 x 10-3 mol x 176 g/mol = 0.547 g
milligrams of ascorbic
acid neutralized
0.547 g x 1000 mg/g = 547 mg
percent error (obj. 2)
(547 g – 500 g) / 500 g x 100% = 9.4%
The first objective of this lab was to determine the experimental molar mass of
ascorbic acid. My experimentally determined molar mass was 192.9 g/mol. The true
molar mass of ascorbic acid is 176 g/mol. Our experimental data yielded an 9.6% error.
Human error that may have caused this positive percent error includes adding too little
NaOH during titration. If too little NaOH was added, then my calculations would
conclude a fewer number of moles neutralized than should have been concluded.
With the number of moles calculated being too small, then the ratio of grams per mole
would be wrong and would produce a number that is too large. This could very well be
one of the reasons why my experimental molar mass is larger than the true molar mass.
Another possible error is that the tablet was not pure ascorbic acid. The tablet included
binders which contributed to the mass of the tablet. Given that the mass of the
ascorbic acid was actually less than the mass of the whole tablet, then the ratio of
grams of ascorbic acid per moles would be smaller. With a smaller mass number, the
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calculations for experimental molar mass could have been smaller and therefore closer
to the true molar mass.
The second objective was to determine the amount (in milligrams) of ascorbic
acid in a vitamin C tablet. The true value for the mass of ascorbic acid is 500 mg. My
experimental value was 547 mg. This gives me a 9.4% error. This positive error could be
attributed to too many moles of ascorbic acid calculated. If too much NaOH was
added during the titration, then my calculations would conclude a larger number of
moles of ascorbic acid being neutralized. With a smaller number of moles being
multiplied by the molar mass, the result would be a smaller number of milligrams that
could be closer to the true value. It is also possible that the true value was wrong. The
true value could have been an approximation. If the true value was an approximation,
then it is possible that my experimental value was closer to correct producing a smaller
percent error.
1. Blake, Robert. "Ascorbic Acid in Vitamin C Tablets." CHM 151L. Glendale
Community College. Glendale, Arizona.
2. Kramer, B.K., V.M. Pultz, and J.M. McCormick. "Vitamin C Analysis." CHEM130.
Truman State University. Kirksville, Missouri. 25 Apr 2011.
3. Morasch, Ralph. "Analysis of Vitamin C." Chem-131 Lab-08. Pierce College.
Lakewood, Washington.
4. Reed, Robin. "ANALYSIS OF VITAMIN C." CHEM 1021. Austin Peay State University.
Clarksville, Tennessee.