# Determining the Vitamin C Content of Juices

Assignment 1 of Procedure 1

1. Calculate the molarity of the freshly-prepared ascorbic acid standard (known strength) solution:

(a) Mass of ascorbic acid used:0.05g

(b) Moles of ascorbic acid (MW=176.1 g/mol):not sure 0.000284

(c) Total volume of standard solution prepared (mL):100.01mL

(d) Ascorbic acid concentration in standard solution (mol/L):not sure 0.000003

2. For each titration, record and calculate the following. (Be sure to show your data and calculations clearly).

(a) Volume of iodine solution added (mL):#1 21.00mL

#2 20.05mL

(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2 (where V1 and V2 are the volumes of the two concentrated chemical solutions added to the flask):

3. From your two or more best "fine" titrations, calculate the average iodine concentration, thus determining today's standard value for the iodine solution.

(Enter Assignment Report)

Assignment 1 of Procedure 2

1. For each titration of the fresh orange juice, record and calculate the following:

(a) Volume of iodine solution added (mL):#1 11.00mL

#2 10.55mL

(b) Concentration of the ascorbic acid in the juice, using the formula M1*V1 = M2*V2:

2. Calculate the average ascorbic acid concentration for the fresh orange juice.

3. For each titration of the week-old orange juice, record and calculate the following:

(a) Volume of iodine solution added (mL):#1 2.00mL

#2 1.05mL

(b) Concentration of the ascorbic acid in the juice, using the formula M1*V1 = M2*V2:

4. Calculate the average ascorbic acid concentration for the week-old orange juice.

5. Report the average amount of ascorbic acid in each of the 2 containers of commercial orange juice in units of mg per mL of juice. (The molecular weight of ascorbic acid is 176.12).

6. The recommended minimum daily requirement for vitamin C is 60 mg per day.

6a. What percentage of this requirement is in one cup (200 mL) of fresh orange juice?

6b. What percentage of this requirement is in one cup (200 mL) of week-old orange juice?

7. What happens to the the ascorbic acid in orange juice over time? (hint: oxygen makes up 20% of our air.)

THIS MAY HELP

Procedure 1

PLEASE NOTE: Titration requires several repeats in order to obtain exact results. The procedures described in this lab assume that you have already done the TITRATION TUTORIAL and are familiar with the technique.

The prepared iodine solution on the Chemicals shelf with a stated (anticipated) concentration of 0.0015M is standardized (to determine the current actual concentration) by performing the following titration:

1. Prepare a solution of pure ascorbic acid of known strength.

1a. Take a clean volumetric flask from the Glassware shelf and place it on the workbench.

1b. Weight out and add exactly 0.05 g ascorbic acid to the volumetric flask. (This precise weighing can be more difficult in the real world, where one more "chunk" of ascorbic acid can put the total weight over the target.)

1c. Fill the volumetric flask with water. (This is done by adding the water normally and checking the "Fill to the Mark" option instead of entering an amount.) This is the most precise way of making a 100 mL solution. Record the amount of ascorbic acid used and the total volume prepared.

2. Do a rough (approximate) titration of the iodine stock solution.

2a. Take a 150 Erlenmeyer flask from the Glassware shelf and place it on the workbench.

2b. Pour 10.0 mL of the ascorbic acid solution into the flask.

2c. Add 10.0 mL of distilled water to the flask to dilute the ascorbic acid solution added.

2d. Add 1.0 mL of the starch indicator to the flask.

2e. Take a burette from the Glassware shelf and place it on the workbench.

2f. Fill the burette with 50.00 mL of the iodine stock solution (with nominal strength of 0.0015M.)

2g. Drag the flask and drop it on the lower half of the burette (the "base of the stand".)

2h. Open the Properties window and move the window to a convenient location on the computer screen. Click back on the burette, and then click the pushpin in the corner of the Properties window to lock it on the burette.

2i. Titrate the ascorbic acid sample by adding iodine until the solution in the flask turns dark blue. (As you learned in the Titration Tutorial, you should do this first rough titration by adding the iodine solution 1-2 mL at a time in order to quickly find the range in which the endpoint is reached.)

2j. Refill the burette with additional iodine solution to bring it back to the 50.00 mL mark. Be careful not to overfill the burette, since it would make a mess on a real lab bench.

3. Do two or more "fine" (or precise) titrations of other samples of the standard ascorbic acid solution. Strive to cause the starch indicator change color when adding only one more drop (the "last 0.05 mL") of iodine solution.

4. Calculate the current molarity (today's "standardized value") of the iodine solution using the average of your best titration results of the ascorbic acid solution.

Procedure 2

1. Determine the ascorbic acid concentration in commercial orange juice from a freshly opened container.

1A. Prepare a sample of orange juice from the new container by adding 10.0 mL of the juice to a clean Erlenmeyer flask. Add 10.0 mL of water and 1.0 mL of starch indicator.

(Please excuse the very light color of the OJ - the pulp was COMPLETELY filtered out)

1B. Titrate the orange juice with the recently-standardized iodine solution. First do a rough titration and then two accurate titrations. Record the volume of iodine delivered in each titration and then refill the burette.

2. Repeat steps 1A and 1B above using the week-old orange juice.

© BrainMass Inc. brainmass.com June 18, 2018, 9:42 am ad1c9bdddf#### Solution Preview

1. Calculate the molarity of the freshly-prepared ascorbic acid standard (known strength) solution:

(a) Mass of ascorbic acid used: 0.05 g

(b) Moles of ascorbic acid (MW=176.1 g/mol):

(0.05 g)(1 mol/176.1 g) = 0.000284 mol

(c) Total volume of standard solution prepared (mL): 100.01 mL

(d) Ascorbic acid concentration in standard solution (mol/L):

0.000284 mol/0.10001 L = 0.00284 M

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2. For each titration, record and calculate the following. (Be sure to show your data and calculations clearly).

(a) Volume of iodine solution added (mL):#1, 21.00 mL

#2, 20.05 mL

(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2 (where V1 and V2 are the volumes of the two concentrated chemical solutions added to the flask):

Here, we know that we added 10.0 mL of ascorbic acid solution to the flask, right? Therefore, what is MV for ascorbic acid? It is (0.00284 mol/L)(0.010 L) = 0.0000284 mol. Therefore, we know that we must have the same amount of moles of iodine at the endpoint of the titration, since MV=MV. Therefore, we know that MV (for iodine) = 0.0000284 mol.

0.0000284 mol = (x M)(0.02100 L)

If we solve for ...

#### Solution Summary

A highly detailed step by step explanation on how to perform these calculations involving titrating juice to find out the ascorbic acid concentration.