Ok, Vitamin C is also known as absorbic acid. Therefore a simple acid-base titration can determine the amount of vitamin C present in the juice of the orange. The procedure you will follow is really close to the procedure used in an AP Biology lab that deals with enzymes and catalysts. If your school has AP Bio, get the procedure for the this lab. You can pretty much use the same procedures, but check with a chemistry teacher to make sure that you use the right base for the titration.
A sample procedure is below:
Experiment: Analysis of Vitamin C
In this lab you will learn to use the back titration procedure to analyze an unknown and a juice sample for the concentration of Vitamin C.
Vitamin C (or ascorbic acid), is a vitamin commonly found in fruits and vegetables which serves a wide variety of biochemical functions. In addition to preventing scurvy (no longer a problem in the developed world), Vitamin C is known to assist in the absorption of iron and to improve resistance to infection.
Vitamin C is a compound that is readily oxidized, and is therefore a good reducing agent. Because it is preferentially oxidized in the body it serves as an antioxidant, protecting other substances in the body from oxidation. Recent studies suggest that the antioxidant properties of Vitamin C are important part of its cancer fighting ability. Its formula is C6H8O6.
Since Vitamin C is an acid it would seem a simple matter to use an acidâ€“base titration to measure its concentration. But juices, fruits and vegetables contain other acids besides Vitamin C, so an acidâ€“base titration would detect not only the Vitamin C, but also the citric and other acids that are present. To circumvent this obstacle we will use a method that relies on Vitamin Câ€™s properties as a reducing agent.
To analyze for Vitamin C content a sample is titrated with a solution of triodide ion, I3-. The triodide ion oxidizes the Vitamin C to dehydroascorbic acid, C6H8O7, as shown in equation (1):
(1) I3- + C6H8O6 + H2O â€”â€”> C6H8O7 + 3 I- + 2 H+
In this experiment, you will add an excess amount of I3- to the solution containing the ascorbic acid. You will then titrate the excess I3- with a solution of thiosulfate, S2O3-2, as shown in equation (2).
(2) I3- + 2 S2O3-2 â€”â€”> 3 I- + S4O6-2
Since the original amount of I3- is known, and the excess is determined by titration with thiosulfate, the amount needed to react with the ascorbic acid can be found by subtraction. This procedure is called a back titration. You will prepare your own solution of sodium thiosulfate, and standardize it with a known quantity of I3-.
There is just one other problem. Solutions of I3- cannot be directly prepared from any solid reagent. In other words, you cannot go to the shelf and find, for example, NaI3. In order to generate I3-, a solution of KIO3 is reduced with solid KI in acidic conditions, as shown in equation (3).
(3) 6 H+ (aq) + KIO3 (aq) + 8 KI (s) â€”â€”> 9 K+ (aq) + 3 I3- (aq) + 3 H2O (l)
The amount of I3- produced can be calculated from the amount of KI and KIO3 used. This amount, minus the excess titrated with the thiosulfate, represents the amount of I3-that reacted with the ascorbic acid.
Your task will be to analyze two samples. One will be a sample of juice and the other will be an unknown sample that you will attempt to analyze to within 5% of the true value. After completing calculations for the juice sample, you will report your results to the instructor who will post them for the class. You will then be able to analyze the class results for precision, and can compare the class average for each juice with the value on the product label.
In this laboratory you will be working individually.
Part I â€” Preparation and Standardization of Sodium Thiosulfate
Prepare 250 mL of a 0.0800 Msolution of sodium thiosulfate, Na2S2O3â€¢5H2O. Stopper and label the flask.
Clean a buret. Rinse and fill with the sodium thiosulfate solution.
For your first titrations, obtain about 80 mL of 0.0200 MKIO3 from the lab cart.
Pipet 25.00 mL of the KIO3 solution into a 250 mL Erlenmeyer flask.
Add 20 mL of 0.5 MH2SO4 and 2.00 g of KI to the flask. Stir to ensure that all of the KI dissolves. The solution should become a redâ€“brown color due to the presence of I3- ion.
Begin titrating the solution with the thiosulfate solution. The redâ€“brown color should fade as the I3- is consumed. When the solution takes on a yellow color, add a few drops of the starch indicator. The solution should turn a deep blue, or possibly a greenâ€“brown. Continue titrating until the blue color has disappeared. Record the amount of thiosulfate used.
Repeat the titration one more time, or as needed to ensure precision.
Calculate the concentration of the thiosulfate solution. You already know the approximate concentration; this calculation will give you the exact value.
Part II â€” Analysis of an Unknown Sample
Select an unknown sample from the lab cart, and record its number in your lab book. Pipet 25.00 mL of the solution into a 250 mL Erlenmeyer flask. Add 20 mL of 0.5 MH2SO4.
Add 1.00 g of KI and 25.00 mL of 0.0200 MKIO3 to the flask. Begin titrating with the thiosulfate solution. As before, add the starch solution when the redâ€“brown color fades and continue titrating until the blue color disappears.
Repeat the experiment as needed for precision. Note that Vitamin C oxidizes readily! If you start your analysis on one day, and return to finish it on another, it is likely that your sample will have changed. Try to start and finish the unknown on the same day.
Part III â€” Analysis of Vitamin C in Juice
Use a graduated cylinder to measure out a 25.0 mL sample of either pineapple or grapefruit juice. Add the juice to a 250 mL Erlenmeyer flask, and then add 20 mL of 0.5 MH2SO4.
Add 1.00 g of KI and 25.00 mL of 0.0200 MKIO3 to the flask. Begin titrating with the thiosulfate solution. As before, add the starch solution when the redâ€“brown color fades and continue titrating until the blue color disappears. However, note that the solution will not turn completely clear, due to the color of the juice. You will have to make your best estimate as to when to stop the titration.
Repeat the experiment a second time.
Analysis and Report
For each titration, determine how many moles of I3- are produced from the reaction between KI and KIO3.
Likewise, determine how many moles of I3- are consumed from the reaction with thiosulfate in each titration.
The difference between the above two quantities will be the moles of I3- that reacted with the ascorbic acid. From the balanced equation, determine how many grams of ascorbic acid were in the original sample for each titration.
Report your average concentration of Vitamin C in the juice sample using units of mg of Vitamin C/L sample. Also report the concentration of Vitamin C in the juice indicated on the bottle label. You will be expected to complete the calculations and report the experimental and label concentrations for the juice sample within two days so that results can be posted for all students.
In your report of the lab analyze the class results for precision and compare the class average for each juice with the value on the bottle label.
Report your average concentration of Vitamin C in your unknown to your instructor, and obtain the correct value. Calculate your percent error, and evaluate your results in your analysis of the lab.
If you use this procedure, instead of taking a "sample off of the cart," use your oranges. Try to figure out these procedures, compare them to the AP Bio procedures, and talk with the chemistry teacher at your school to make sure that the procedure you decide to use will work.