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.

http://www.jiskha.com/display.cgi?id=1495566182

To calculate the average number of moles of ascorbic acid present in the titrated samples, you need to use the volume of iodine added, the molarity of the iodine solution, and the balanced equation for the reaction.

In this case, the balanced equation for the reaction is:
ascorbic acid(aq) + I2(aq) → dehydroascorbic acid(aq) + 2 I−(aq)

From the data provided, you have the volume of iodine added for each trial. You can calculate the moles of iodine added using the molarity of the iodine solution.

For example, let's calculate the moles of iodine added for trial 1:
Volume of iodine added = Final Burette Reading - Initial Burette Reading = 11.08 mL - 0.05 mL = 11.03 mL
Moles of iodine added = Volume of iodine added (in L) * Molarity of iodine solution
= 0.01103 L * 0.00200 mol/L
= 2.206 x 10^-5 mol

Repeat this calculation for each trial and calculate the average moles of iodine added.

To determine the moles of ascorbic acid in the titrated samples, you need to use the stoichiometry of the balanced equation.

From the balanced equation, you can see that for every 1 mole of ascorbic acid, 1 mole of iodine is consumed. This means that the moles of ascorbic acid in the sample is equal to the moles of iodine added.

Therefore, the average number of moles of ascorbic acid present in the titrated samples is equal to the average moles of iodine added.

The volume of juice tested was 10.00 mL. To calculate the amount of ascorbic acid that would be ingested if the entire juice box (200 mL) were consumed, you can use stoichiometry.

Using the average number of moles of ascorbic acid (which is equal to the average moles of iodine added) and the volume of juice tested:
Moles of ascorbic acid in juice box = Average number of moles of ascorbic acid * (200 mL / 10 mL)

For example, if the average number of moles of ascorbic acid is 2.5 x 10^-5 mol:
Moles of ascorbic acid in juice box = 2.5 x 10^-5 mol * (200 mL / 10 mL) = 5 x 10^-4 mol

To calculate the mass of ascorbic acid in the juice box, you need to use the molar mass of ascorbic acid (176.14 g/mol).

Mass of ascorbic acid in juice box = Moles of ascorbic acid in juice box * Molar mass of ascorbic acid

For example, if the moles of ascorbic acid in the juice box is 5 x 10^-4 mol:
Mass of ascorbic acid in juice box = 5 x 10^-4 mol * 176.14 g/mol = 0.08807 g

To compare this value to the mass of vitamin C you expected in the juice (based on the information provided on the juice box), you need to convert the mass of ascorbic acid to a percentage of the recommended daily amount.

What is the mass of ascorbic acid you expect to find in the juice box if the recommended daily amount is 90 mg?

To calculate the mass of ascorbic acid based on the recommended daily amount, you need to convert the 90 mg to grams and then calculate the percentage.

Mass of ascorbic acid expected = 90 mg * (1 g / 1000 mg) = 0.09 g

Percentage of ascorbic acid in juice box = (Mass of ascorbic acid in juice box / Mass of ascorbic acid expected) * 100%

For example, if the mass of ascorbic acid in the juice box is 0.08807 g:
Percentage of ascorbic acid in juice box = (0.08807 g / 0.09 g) * 100% = 97.86%

Finally, to calculate the average volume of iodine used, you need to calculate the average of the volumes of iodine added in each trial.

Average volume of iodine used = (Volume of iodine added in trial 1 + Volume of iodine added in trial 2 + Volume of iodine added in trial 3 + Volume of iodine added in trial 4) / 4

For example, if the volumes of iodine added in the trials are 11.03 mL, 10.19 mL, 10.23 mL, and 10.22 mL:
Average volume of iodine used = (11.03 mL + 10.19 mL + 10.23 mL + 10.22 mL) / 4 = 10.9175 mL