Determination of Vitamin C


Vitamin C (ascorbic acid) deficiency leads to scurvy, a disease characterized by weakness, small hemorrhages throughout the body that cause gums and skin to bleed, and loosening of the teeth. Sailors that were out at sea for months on end would often develop scurvy unless the captain had the foresight to pack limes and other citrus fruits. Vitamin C is a water soluble antioxidant, and plays a vital role in protecting the body. Oxidizing species attack the body from many directions. Smog and cigarette smoke both contain high levels of oxidizing molecules that cause tissue damage. The body makes oxidizing molecules in response to an infection, and these molecules both kill the infecting organism and cause tissue damage. The body absorbs extra vitamin C in response to an infection. Because it is a water soluble vitamin, any unused vitamin C is excreted. The minimum daily requirement is 30 mg, the recommended daily allowance is 60-70 mg.

The formula for ascorbic acid is C6H8O6 and the structures for the reduced form and for the oxidized form (dehydroascorbic acid) are shown below:

The amount of ascorbic acid can be determined by a redox titration with a standardized solution of iodine. The iodine is reduced by the ascorbic acid to form iodide. As shown in the other half of this redox equation.

The titration end point is reached when a slight excess of iodine is added to the ascorbic acid solution. Thyodene is used to determine the endpoint, excess iodine reacts with the thyodene indicator and forms a highly colored complex. Thyodene does not form this complex with iodide. (Note: Thyodene is available from Fisher Scientific Company, and the CAS number is 9005-84-9.)


Standardize an iodine solution to determine its molarity.

Prepare a standard solution by weighing 0.1200 - 0.1300g of ascorbic acid and dissolving it in exactly (use a volumetric flask) 250 ml of distilled water. Be sure to put the values for the amount you weighed and the calculated molarity in your notebook. For the first titration, pour approximately 50 ml of this standard ascorbic acid solution into a clean 100 ml beaker. Rinse a 10 ml volumetric pipet with this solution, then pipet 10.00 ml into a 250 Erlenmeyer flask. Add approximately 10 ml of distilled water, then add a spatula-tip of the thyodene indicator.

Rinse a clean 50 ml buret with some iodine solution and then fill the buret (somewhere between 0.00 and 1.00). Record this reading to the nearest 0.01ml in your notebook. Add this iodine solution to the Erlenmeyer flask containing the ascorbic acid. Constantly swirl the flask as you add the iodine. Continue adding iodine until the entire flask turns dark blue with the addition of one drop. The first time the entire solution changes is the endpoint. You may have to do a rough titration to determine approximately what volume is needed to reach the end point. On subsequent titrations, you can go quickly to this approximate volume and then slowly add drops and half drops until you reach the exact end point.

Do a total of three accurate titrations with one standard solution. Determine the average molarity of your standardized iodine solution. Determine the standard deviation of that average.

Determine the ascorbic acid concentration in commercial orange juice.

You will determine the ascorbic acid concentration in two preparations of commercial juice.

  1. Juice in a freshly opened container
  2. Juice in a container that was opened several days ago

To prepare each solution for testing, pour about 90 ml of the test solution into a beaker and add about 1.5 gram of oxalic acid. Take 25.00 ml of this prepared solution and titrate with the standardized iodine solution, using a spatula tip of the thyodene indicator.

Do three accurate titrations for each solution and determine an average value (and the standard deviation) for the amount of vitamin C present.


You must report your results using a spreadsheet. Report the average molarity of your standardized iodine solution and the standard deviation for the average. Report the average amount (and standard deviation) of ascorbic acid in the commercial orange juice as mg/100ml. Remember you used 25 ml samples!

All of your values must be placed in a spreadsheet. Do the calculations in the spreadsheet and print out two versions of your completed spreadsheet, one with the values and one with the formulas. (2 points)

You will be graded on the precision (2 points) and accuracy (2 points) of your results. Be sure to report the correct number of significant figures (-1 point).


Please write your conclusion in well organized paragraphs using complete sentences. You must answer the following questions in the Conclusion section. (3 points)

Using the balanced redox equation, how many electrons are lost when ascorbic acid is oxidized to dehydroascorbic acid? Where do those electrons go?
Which solution contained more ascorbic acid?
Explain your results--why would one solution have more vitamin C than the other? What could be happening to the ascorbic acid? (Remember that Oxygen makes up ~20% of our air!)
What conclusions can you draw about the orange juice in your refrigerator?
What other foods contain vitamin C? List three and report their approximate vitamin C content in mg per serving. You may need to do some calculations to determine the amount present in mg, as most labels list the percent of RDA (recommended daily requirement). The RDA for vitamin C is 60 mg. If a food contains 100% of the RDA per serving, it has 60mg of vitamin C.

VitC/kstone/revised 07/10/02