Lab: Determining the Quantity of Vitamin C in Fruit Juice
Vitamin C, also called ascorbic acid, is commonly found in commercial fruit juices and drinks. In this activity you will analyze data collected from a titration analysis of a fruit juice.
The nutrition facts for a 200-mL juice box are shown. The value most important to this investigation is for vitamin C. The value given for vitamin C is 80% of the daily value.A structural diagram for vitamin C is shown, along with the molar mass of vitamin C, 176.14 g/mol.
Background Information
Ascorbic acid reacts with iodine in solution, as described by the following reaction:
ascorbic acid(aq) + I2(aq) → dehydroascorbic acid(aq) + 2 I−(aq)
In this procedure a standard aqueous iodine solution is added to a sample of juice. The initial reaction involves iodine reacting until the ascorbic acid in the juice sample depletes. The endpoint of this titration is a blue colour, signified by the reaction of excess iodine with starch (which is added to the juice prior to titration).
Earlier in this lesson you calculated a mass of vitamin C (ascorbic acid) that you would expect to find in the juice if it met the 80% of the daily recommended amount.
Purpose
The purpose of this investigation is to test the manufacturer’s claim that the juice product contains 80% of the daily recommended amount of ascorbic acid.
Problem
What mass of vitamin C (ascorbic acid) is present in a box of juice?
Materials
lab apron
eye protection
fruit juice (200-mL box)
0.002 00-mol/L iodine solution
starch indicator solution
distilled water
50-mL burette and stand
stirring rod
small funnel
10-mL volumetric pipette and bulb
clean, dry beaker
125-mL Erlenmeyer flask
Procedure
Step 1: Assemble the ring stand and burette clamp. Clean the burette using distilled water, and wash using a small quantity of the aqueous iodine solution. Place the burette in the clamp.
Step 2: Fill the burette with the aqueous iodine solution.
Step 3: Read and record the volume in the burette.
Step 4: Clean the pipette using distilled water, and wash using a small quantity of the fruit juice. Use the pipette to transfer 10.00 mL of juice to the Erlenmeyer flask.
Step 5: Add a few drops of the starch indicator solution to the Erlenmeyer flask.
Step 6: Add 40 mL of distilled water to the Erlenmeyer flask.
Step 7: Titrate the juice sample to the desired endpoint—a permanent dark blue colour. Measure and record the burette reading at the endpoint. Record the endpoint colour.
Step 8: Repeat steps 3 to 7 three more times (four trials altogether).
Observations
Titration of 10.00 mL Fruit Juice with 2.00 x 10−3 mol/L I2(aq)
Trial
1
2
3
4
Final Burette Reading (mL)
11.08
21.27
31.50
41.72
Initial Burette Reading (mL)
0.05
11.03
21.27
31.50
Volume of I2(aq) Added (mL)
11.03
10.19
10.23
10.22
Final Colour of Solution
purple
blue
blue
blue
Analysis
Read the background information and procedure for this investigation. Compare and contrast this procedure with the titration you performed earlier.
Use the data to calculate the average number of moles of ascorbic acid present in the titrated samples.
The volume of juice tested was 10.00 mL. What amount of ascorbic acid would be ingested if the entire juice box (200 mL) were consumed?
Calculate the mass of ascorbic acid in the juice box. Compare this value to the mass of vitamin C you expected in the juice (based on the information provided on the juice box).
What is the mass of ascorbic acid you expect to find in the juice box if the recommended daily amount is 90 mg?
What is the average volume of iodine that is used? show your calculations.
I am willing to help but I am not willing to to this complete lab for you.
To calculate the average volume of iodine that is used, we need to find the difference between the final and initial burette readings for each trial.
Here are the calculations for each trial:
Trial 1:
Final Burette Reading = 11.08 mL
Initial Burette Reading = 0.05 mL
Volume of I2(aq) Added = Final Burette Reading - Initial Burette Reading
= 11.08 mL - 0.05 mL
= 11.03 mL
Trial 2:
Final Burette Reading = 21.27 mL
Initial Burette Reading = 11.03 mL
Volume of I2(aq) Added = Final Burette Reading - Initial Burette Reading
= 21.27 mL - 11.03 mL
= 10.24 mL
Trial 3:
Final Burette Reading = 31.50 mL
Initial Burette Reading = 21.27 mL
Volume of I2(aq) Added = Final Burette Reading - Initial Burette Reading
= 31.50 mL - 21.27 mL
= 10.23 mL
Trial 4:
Final Burette Reading = 41.72 mL
Initial Burette Reading = 31.50 mL
Volume of I2(aq) Added = Final Burette Reading - Initial Burette Reading
= 41.72 mL - 31.50 mL
= 10.22 mL
To find the average volume of iodine, we can add up all the volumes and divide by the number of trials:
Average Volume of I2(aq) =
(11.03 mL + 10.24 mL + 10.23 mL + 10.22 mL) / 4
= 41.72 / 4
= 10.43 mL