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.

Clown Bot is here to answer your questions with a touch of humor!

1. The procedure for this investigation involves titrating the fruit juice with an iodine solution and using starch as an indicator. In comparison to the previous titration you performed, well, at least this time you get to add some color to your experiment with the iodine!

2. To calculate the average number of moles of ascorbic acid present in the titrated samples, you need to use the formula: moles = volume x concentration. So, multiply the volume of iodine added (in liters!) by the concentration of the iodine solution.

3. If the volume of juice tested was 10.00 mL and the entire juice box is 200 mL, then you can multiply the number of moles of ascorbic acid by a factor of 20 to determine the amount of ascorbic acid ingested. Just imagine the small juice box expanding like a balloon!

4. To calculate the mass of ascorbic acid in the juice box, you need to know the molar mass of ascorbic acid. Multiply the number of moles of ascorbic acid you calculated earlier by the molar mass to get the mass. Then compare that mass to the mass of vitamin C you expected in the juice box. Will it be a match made in fruity heaven?

5. If the recommended daily amount of ascorbic acid is 90 mg, you can convert that to moles using the molar mass of ascorbic acid. Then multiply the number of moles by the molar mass to get the mass of ascorbic acid you expect to find in the juice box. Will it be a pleasant surprise or a citrus disappointment?

6. The average volume of iodine used can be calculated by subtracting the initial burette reading from the final burette reading for each trial. Then take the average of these numbers. Remember to convert from mL to L before using it in any calculations. How much iodine will you get to play with on average?

I hope my answers brought a smile to your face!

To compare and contrast this procedure with the titration performed earlier, we need to identify the similarities and differences:

Similarities:
1. Both procedures involve titration.
2. Both procedures use iodine solution as the titrant.
3. Both procedures use starch indicator to signal the endpoint of the titration.

Differences:
1. The earlier titration was performed on an unknown solution, while this titration is performed on fruit juice.
2. The earlier titration aimed to determine the concentration of a substance, while this titration aims to determine the mass of ascorbic acid in the juice.
3. The earlier titration required the calculation of molar concentration, while this titration requires the calculation of moles and mass of ascorbic acid.

To calculate the average number of moles of ascorbic acid present in the titrated samples, we use the formula:

Moles of ascorbic acid = Moles of iodine used

The moles of iodine used can be calculated using the formula:

Moles of iodine = Volume of iodine (L) x Molarity of iodine (mol/L)

Calculating the moles of iodine used in each trial:
Trial 1: 0.01103 L x 0.002 mol/L = 2.206 x 10^-5 mol
Trial 2: 0.01019 L x 0.002 mol/L = 2.038 x 10^-5 mol
Trial 3: 0.01023 L x 0.002 mol/L = 2.046 x 10^-5 mol
Trial 4: 0.01022 L x 0.002 mol/L = 2.044 x 10^-5 mol

Calculating the average moles of ascorbic acid:
Average moles = (2.206 x 10^-5 mol + 2.038 x 10^-5 mol + 2.046 x 10^-5 mol + 2.044 x 10^-5 mol) / 4

The volume of juice tested was 10.00 mL. To calculate the amount of ascorbic acid ingested if the entire juice box (200 mL) were consumed, we use the proportion:

Moles of ascorbic acid in 200 mL = Average moles of ascorbic acid x (200 mL / 10 mL)

To calculate the mass of ascorbic acid in the juice box, we use the formula:

Mass of ascorbic acid = Moles of ascorbic acid x Molar mass of ascorbic acid

To compare the calculated mass of ascorbic acid with the expected mass on the juice box, you can subtract the calculated mass from the expected mass and compare the difference.

If the recommended daily amount is 90 mg, we can calculate the moles of ascorbic acid corresponding to 90 mg using the formula:

Moles of ascorbic acid = Mass of ascorbic acid (g) / Molar mass of ascorbic acid

Then, we can convert the moles of ascorbic acid to volume using the molar concentration of iodine:

Volume of iodine (L) = Moles of ascorbic acid / Molar concentration of iodine

Finally, to calculate the average volume of iodine used, you can find the average of the volumes from all trials.

To answer the question about the average volume of iodine used, we need to calculate the average of the volumes of iodine used in the four trials. Here are the steps to get the answer:

1. Write down the volumes of iodine used in each trial:
- Trial 1: 11.03 mL
- Trial 2: 10.19 mL
- Trial 3: 10.23 mL
- Trial 4: 10.22 mL

2. Calculate the average volume by summing up the volumes from each trial and dividing by the number of trials:
- Total volume = 11.03 mL + 10.19 mL + 10.23 mL + 10.22 mL = 41.67 mL
- Average volume = Total volume / Number of trials = 41.67 mL / 4 = 10.42 mL

Therefore, the average volume of iodine used in the titration is 10.42 mL.

You have typed your entire lab experiment here. I'm not inclined to answer each line by line question but I'll be glad to help if you tell me what I can help you with. Note that you have initial volumes and final reading the buret. Subtract the two to find the volume used in the titration. Basically you will need tis information.

grams vitamin C = mols vitamic C x molar mass.
mols vitamin I2 = M I2 x L I2 = ?
Use the coefficients in the balanced titration equation to convert mols I2 to mols ascorbic acid.
Then grams ascorbic acid = mols x molar mass.
% vitamin C = (grams vitamin C/grams sample)*100 = ?

Show any work you have done.