# 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.

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#### 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

----------------------------------------------

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.

Vitamin C Content of Juices Lab

Lab 8

DIRECTIONS:

The minimum daily requirement for vitamin C is 30mg and the recommended daily allowance is 60-70 mg. You can find the vitamin C content on the labels of most commercially available juices. How close is their reported number to the actual amount of the vitamin in the juice? You will test that in this lab by titrating the juice with an iodine solution using starch as an indicator. The formula for ascorbic acid is C6H8O6 and the structures for the reduced form and for the oxidized form (dehydroascorbic acid) are shown, which illustrates the first half of redox titration with a standardized solution of iodine. In this procedure, the iodine is reduced by the ascorbic acid to form iodide as show in the other half of this redox equation:

I2 + 2e ---> 2I-

The titration end point is reached when a slight excess of iodine is added to the ascorbic acid solution. To highlight when the excess iodine, thyodene indicator is added to the titrant. Thyodene is a starch that reacts with the excess iodine to form a bright blue complex. Thyodene does not form this complex with iodide. Before doing a titration on a solution of unknown concentration, you must first prepare a standardized iodine solution. A standardized iodine solution is one whose concentration is known exactly from its preparation and is then verified by a controlled titration.

The iodine solution on the chemicals shelf has been prepared according to a calculation that gives a concentration of 0.015M. Your task in Procedure 1 is to confirm this exact concentration of the iodine solution. In Procedure 2, you will use the standardized iodine solution to determine the concentration of vitamin C in samples of orange juice. As in any titration we use the formula: C1*V1 = C2*V to calculate the unknown concentration where C is the concentration in mol/L and V is the Volume of the solution. When standardizing the iodine, C1 and V1 for the absorbic acid solution are known, as the volume of iodine delivered from the burette and so you can solve for the concentration of the iodine solution.

ASSIGNMENT PROCEDURE 1:

The prepared iodine solution on the chemicals shelf with a stated concentration of 0.0015M is standardized (confirming the concentration) by performing the following titration:

1. Take a clean volumetric flask and put on workbench

2. Add 0.1g ascorbic acid to the volumetric flask

3. Fill the flask with water (making an appx 100ml solution) Record the amount of ascorbic acid used and the total volume prepared

4. Take a 150 Erlenmeyer flask and put on workbench

5. Pour 20ml of the absorbic acid solution( from volumetric flask)

6. Add 1ml Starch Indicator to Erlenmeyer flask

7. Burette fill with 50ml iodine solution with an approximately known concentration of 0.015M

8. Titrate the absorbic acid in Erlenmeyer flask at 1ml increments

QUESTIONS:

1. Calculate the molarity of the ascorbic acid solution:

(a) Mass of ascorbic acid used

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

(c) Volume of solution (mL)

(d) Ascorbic Acid Concentration (mol/L)

2. For each titration record and calculate the following:

(a) Volume of iodine solution added (mL)

(b) Concentration of the iodine solution

3. Calculate the average iodine concentration using the formula M1*V1 = M2*V2

PROCEDURE 2: Scenario:

1. Determine the ascorbic acid concentration in commercial orange juice, from a freshly opened container and from a container that was opened one week ago.

2. Prepare a sample of orange from the NEW container by adding 40ml of the juice to a clean Erlenmeyer flask. Add 1ml Starch Indicator

3. Titrate the orange juice with the standardized iodine solution.

4. Repeat steps 2 and 3 to the WEEK OLD orange juice.

QUESTIONS:

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

(a) Volume of iodine solution added (mL)

(b) Concentration of the ascorbic acid in the juice

2. Calculate the average ascorbic acid concentration for the NEW orange juice, using the formula M1*V1 = M2*V2

3. For each titration of the WEEK OLD orange juice record and calculate the following:

(a) Volume of iodine solution added (mL)

(b) Concentration of the ascorbic acid in the juice

4. Calculate the average ascorbic acid concentration for the WEEK OLD orange juice using the formula M1*V1=M2*V2

5. Report the average amount of ascorbic acid in the 2 samples of commercial orange juice in units of "mg per mL" of juice. The molecular weight of ascorbic acid is 176.12

6. The minimum daily requirements for vitamin C is 60mg per day. What percentage of this requirement is in one cup (200ml) of NEW and WEEK OLD orange juice?

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

LAB FINDINGS PROCEDURE 1:

Clean volumetric flask add 0.1g ascorbic acid

Fill remainder with water (makes 100ml solution)

Clean 150ml Erlenmeyer flask

Pour 20ml ascorbic acid into flask from volumetric flask (above)

Add 1ml Starch Indicator

Burette, fill 50ml Iodine

Titrate with 1ml increments from burette until it turns dark blue

49.00ml - 1ml

48.00ml - 1ml

47.00ml- 1ml

46.00ml - 1ml

45.00ml -1ml

44.00ml - 1ml

43.00ml-1ml

42.00ml - 1ml TURNED DARK BLUE

LAB FINDINGS PROCEDURE 2:

Add 40 ml or juice to Erlenmeyer flask NEW orange juice

Add 1ml Starch Indicator to flask

Burette 50ml Iodine, Titrate in 1ml increments

Repeat steps 2 and 3 for ONE WEEK OLD orange juice

NEW JUICE:

40ml NEW juice

1ml Starch Indicator

Burette 50ml Iodine 1ml increments

49.00ml - 1ml

48.00ml - 1ml

47.00ml - 1ml

46.00ml - 1ml

45.00ml - 1ml

46.00ml - 1ml

45.00ml - 1ml

44.00ml - 1ml TURNED ORANGE

WEEK OLD JUICE:

40ml WEEK OLD juice

1ml Starch Indicator

Burette 50ml Iodine 1ml increments

49.00ml -1ml

48.00ml -1ml TURNED ORANGE