Acid/Base Titration of Ascorbic Acid
Ascorbic acid is a water-soluble vitamin (vitamin C) that must be consumed regularly to ensure proper body
function. Lack of vitamin C may result in scurvy. Citrus fruits, tomatoes, cantaloupes, and fresh vegetables are
good sources of vitamin C. Although many people can obtain their minimum daily requirement (60 mg) from
these sources quite readily, many people prefer to consume larger quantities of this vitamin. This practice of
consuming megadoses (500 mg or more) of vitamin C was popularized in the 1960’s by Linus Pauling who
claimed that consumption of larger doses of ascorbic acid would help cure the common cold and prevent certain
kinds of cancer. Recent demographic studies suggested that men who took daily megadoses of vitamin C have a
significantly longer life expectancy than men who did not. The question, however, remains – is it necessary to
consume vitamin C in pill form, or can one obtain the desired dosages via ingestion of foods naturally rich in
vitamin C? How do we determine the amount of vitamin C contained in a given portion of food? To answer
this question you will perform an acid/base titration of grapefruit juice.
Like any other, acid ascorbic acid will react with sodium hydroxide, a base, to produce a salt and water. The
salt produced is sodium ascorbate.
HC6H7O6
+ NaOH 
H2O + NaC6H7O6
This equation suggests that we can determine the amount of vitamin C is a sample of juice via titration with a
standardized solution of sodium hydroxide. Because the mole ratio of ascorbic acid to sodium hydroxide is 1:1
the number of moles of acid present in the sample is equal to the number of moles of base that must be added to
reach the equivalence point of the titration.
M = moles of the solute / volume of the solution in liters
One problem in performing the titration of acids with stock solutions of sodium hydroxide is that the
concentration of the sodium hydroxide tends to change with time. This results from the reaction between
sodium hydroxide and atmospheric carbon dioxide, which produces sodium bicarbonate. For this reason, you
will standardize your sodium hydroxide solution.
Materials
Buret
Buret clamp and stand
Ascorbic acid
Distilled water
Grapefruit juice
0.1 M Sodium hydroxide
Phenolphthalein indicator
100 mL graduated cylinder
10 mL graduated cylinder
2 Erlenmeyer flasks, 250 mL
2 beakers
Standardization of Sodium Hydroxide
1. Mass out approximately 1 g of powdered ascorbic acid (record the exact mass in the data table) and place it in a
clean 250-mL Erlenmeyer flask. Dissolve the ascorbic acid in approximately 50 mL of water.
2. Add 3 drops of phenolphthalein indicator.
3. Note: If the buret has already been prepared skip to step 3. Prepare a buret for titration by rinsing it with two
small portions of distilled water followed by two 10-mL portions of the sodium hydroxide solution. Fill the
buret. Eliminate any air bubbles.
4. Record the initial buret level.
5. Use the approximate concentration of the NaOH to estimate the volume of the NaOH that will be required to
react completely with the ascorbic acid.
6. Run somewhat less than this amount into the flask containing the ascorbic acid while swirling the flask. Use
your wash bottle to clean the flask walls of drops of base that may have splattered out of the titration mixture.
7. Continue adding the NaOH solution. As you approach the volume estimated to be needed for complete reaction,
add the NaOH more slowly, while continuing to swirl and wash down the sides of the flask.
8. Stop the titration when the addition of a single drop of base changes the color of the solution to a light pink,
indicating that you have reached the end point. Record the final buret reading.
9. Repeat the titration. Calculate the molarity of the NaOH from each trial and take the average.
Titration of a Sample of Juice
1. Use a graduated cylinder to measure a 20-mL sample of juice. Transfer to a clean 250-mL Erlenmeyer flask.
Wash down the sides of the graduated cylinder with distilled water and add this water to the flask. Add an
additional ~65 mL of water to the flask.
2. Add 3 drops of phenolphthalein indicator.
3. Titrate, as above, with sodium hydroxide, stopping when the titration mixture has acquired a permanent, pinkishorange color.
4. Repeat the titration. Calculate the amount of acid in each sample of juice. If all the acid is ascorbic acid, how
many mg of vitamin C are there in a 1 cup Serving?
Lab: Acid/Base Titration of Ascorbic Acid
Name _________________________
Date _________ Per ____ Lab Station____
Balanced Equation: Write a balanced equation. Identify the acid, base and salt.
Data & Calculations Show all work in the space below
Standardization of Sodium Hydroxide:
Trial 1
Trial 2
Mass of ascorbic acid
Moles of NaOH to be delivered *
Estimate of the volume of NaOH
required *
Initial buret reading, NaOH
Final buret reading, NaOH
Volume of NaOH delivered
Molarity of NaOH (standardized) *
Average molarity of NaOH
(standardized) *
1.
Moles of NaOH to be delivered:
2.
Estimate of the volume of NaOH required: (How many mL of NaOH (~0.1 M) will be required to react with your mass
of ascorbic acid?)
3.
Molarity of NaOH:
4.
Average molarity of NaOH:
Titration of a sample of juice:
Trial 1
Trial 2
Volume of juice
Initial buret reading, NaOH
Final buret reading, NaOH
Volume of NaOH delivered
Moles of NaOH delivered *
Moles of ascorbic acid titrated *
Molarity of ascorbic acid in the juice *
Average Molarity of juice *
Average mg vitamin C per cup *
Number of cups to obtain RDA *
1.
Moles of NaOH delivered (use the average molarity you calculated on the front):
2.
Moles of ascorbic acid titrated:
3.
Molarity of ascorbic acid in the juice:
4.
Average Molarity of juice:
5.
Average mg vitamin C per cup (1 liter = 4.226 752 82 cups):
6.
The recommended daily allowance (RDA) of vitamin C is 60 mg. How many cups of grapefruit juice would you need to
consume to obtain your RDA of vitamin C ?