Volumetric Analysis of Vitamin C by Titration
The objective of this experiment is to use a redox reaction titration to accurately determine the amount of vitamin C in a sample of lemon juice, orange juice, or grapefruit juice. Chemistry of Vitamin C The chemical name for vitamin C is L-ascorbic acid. Its molecular formula is C6H8O6; its molar mass is 176. 12 g/mole. Ascorbic acid is found throughout the plant and animal kingdoms, occurring in citrus fruits, hip berries (such as rose hips), fresh tea leaves, tomatoes, broccoli, other fruits and vegetables, paprika, and the adrenal cortex of oxen.
It can be obtained from any of these sources but was originally isolated from and identified in oxen. It was the first vitamin to be prepared in pure form. Ascorbic acid is a white solid that has a sharp, sour taste and dissolves in water. The pure compound is stable to air oxidation when dry, but when impure (as it is in many natural forms) it is readily oxidized when exposed to air and light. Vitamin C is a fairly strong reducing agent and decolorizes many dyes. Its aqueous solutions are rapidly oxidized by air; this reaction is accelerated in basic solution and in the presence of iron and copper ions.
The vitamin-C content of juices can decrease rapidly with time once the juice is exposed to air. (Much of the information reported here about vitamin C was obtained from the Merck Index, Susan Budavari, Ed. , Merck and Co. , Inc. , Rahway, NJ, 1989, which is a good source for physical, chemical, and physiological properties of chemicals, drugs, and biological agents. ) Vitamin C is essential to humans. It is involved in the synthesis of collagen, which constitutes about one-third of the total protein in the human body.
A deficiency of vitamin C results in a disease called scurvy, which is characterized by weakness, swollen joints, bleeding gums and loose teeth, and delayed healing of wounds. Scurvy was common in sailors, who had no fresh fruits or vegetables for long periods. In 1735 James Lind found that scurvy could be avoided if sailors were provided citrus fruit. To this day British sailors are called “limeys” because they were provided limes to prevent scurvy. A quantity of 60 mg vitamin C per day is enough to prevent the disease, and this is the recommended daily dietary allowance (RDA).
Vitamin C is also involved in iron metabolism, and many believe that very large doses are effective in preventing or curing the common cold. In 1970 Linus Pauling, a two-time Nobel Prize winner, published “Vitamin C and the Common Cold”, which stated that doses of 1-2 g per day (18 times the RDA) would prevent colds, and 4010 g/day would cure an existing cold. However, some recent studies do not support this hypothesis. At large doses vitamin C causes problems such as diarrhea and the induction of kidney stones.
Titration Reactions Titrations can be completed on both Acid-base reaction and oxidation-reduction reaction systems. Both types of reactions occur rapidly in aqueous solution, the balanced equations for such reactions can be determined, and there exist techniques (such as changes in color of indicators or the color of the reactants themselves) for determining when the reactants have been mixed in stoichiometric ratios. Since vitamin C is a weak acid and also a good reducing agent, either type of reaction might be used.
This experiment makes use of an oxidation-reduction reaction in which elemental iodine oxidizes ascorbic acid. Iodine is chosen because it is a weak oxidizing agent so it will not oxidize substances other than the ascorbic acid in the sample of fruit juice. As a strong reducing agent, ascorbic acid will reduce I2 to I- very easily. We will use this reaction in conjunction with a starch indicator to determine the number of moles of Vitamin C present. A number of reactions occur during a single titration. The solution to be titrated will consist of KI, acetic acid, starch solution, and ascorbic acid.
In this reaction, the ascorbic acid molecule gains oxygen (in the form of OH groups). Each iodine atom in the I2 molecule accepts an electron and becomes a negatively charge iodide ion. Thus the ascorbic acid molecule is oxidized and the iodine molecule is reduced. Without showing the molecular structure, this equation could be written as: C6H8O6 + I2 + 2H2O ? C6H10O8 + 2I – + 2H+(1) From this equation it is apparent that one mole of iodine is required to react with one mole of ascorbic acid. In other words, the appropriate stoichiometric ratio is: (2)
Titrating to measure the amount of ascorbic acid in a sample of fruit juice can be done by adding a solution of iodine from a buret to the sample containing ascorbic acid in a flask. As the iodine-containing solution is added, the iodine will react with ascorbic acid in the sample according to equation 1 shown above. As long as there is ascorbic acid present iodine will react with it. This reaction will continue until all of the ascorbic acid has been used up, that is, until a mole of iodine has been added for every mole of ascorbic acid that was in the sample to begin with.
This is the equivalence point. As soon as more than an equivalent amount of iodine has been added, the extra iodine cannot be consumed by the ascorbic acid. Therefore, some iodine will remain unreacted in the flask. One way to detect the equivalence point in this titration is to devise a method by which you can detect iodine in the beaker. (See reaction B in the following reaction scheme. ) Excess iodine (I¬2) reacts with iodide ions (I-) to make a triiodide ion (I3-) which forms a very intense blue color when it comes into contact with starch.
This color is due to incorporation of the ions within the molecular structure of the starch; we refer to this as formation of a starch-iodine complex. (Even after the starch-iodine complex has formed, I2 molecules are still there; if something reacts with iodine, the blue color will disappear as the iodine is used up. ) To detect the end point you will add starch to the solution in the flask at the beginning of the titration. As iodine is added from the buret the iodine will react with ascorbic acid and the blue color will not continue to disappear.
When all the ascorbic acid has been used up, the next drop of iodine solution will have nothing to react with but the starch, and the blue color will remain in the solution. The end point of the titration has been achieved when the blue color remains in the solution for at least 30 seconds. (The reason for waiting 30 seconds is that some blue color might form before ascorbic acid has a chance to react with the iodine; if so, it will take a little while for ascorbic acid to react with the iodine that was in the starch-iodine complex. Review of Reaction Scheme: A)
As I2 is formed, it will react with ascorbic acid. C6H8O6 + I2 + 2H2O ? C6H10O8 + 2I – + 2H+ B) As soon as all of the ascorbic acid is consumed the I2 will react with I- to form I3-I2 + I- > I3- C) This will react with the starch indicator to produce the blue-black starch-iodide complex. I3- + starch > starch (I3-) complex (blue-black) The sudden appearance of the blue-black color will indicate that all of the ascorbic acid is consumed.
Since you know the concentration of the standardized I2 solution (written on the bottle), the volume that you used, and the stoichiometry of the reaction you will be able to calculate the amount of Vitamin C present in your original sample. REMINDER:Tips for titrating can be found in the Titration Appendix! EXPERIMENTAL PROCEDURE The Starch-Iodine Complex Color To become familiar with the indicator color, use a graduated cylinder to place 30 mL of distilled water, 1 mL of starch solution, and 1 mL of 6 M acetic acid into a 150 mL beaker.
Add solid potassium iodide (KI) to the beaker, a few crystals are sufficient. Then add drops of iodine solution until you notice a color change; observe the color. The color of the solution can be attributed to the starch-iodine complex. Now add a drop of juice and observe the color. Add four more drops, one at a time, observing the color for about ten seconds after each addition. What can you conclude about the reaction of the vitamin C in juice with iodine in the starch-iodine complex? Save this solution since it will help you to determine the color of the solution at the endpoint of your titration.
The color of the juice used in your analysis may affect the color observed at the end point of a titration. In other words, the color of the end point may be slightly different than the color of the starch-iodine complex observed above. Therefore, it is helpful to carry out a rough titration on your juice to determine the true color change that will occur at the end point and then compare each trial to this rough titration result. Determining the Concentration of Vitamin C in Juice Samples of lemon juice, orange juice, and grapefruit juice will be available in the laboratory.
Your objective is to determine the concentration of vitamin C in one of these juices. Be sure to record the type of juice you work with and the exact designation of amount of Vitamin C on the label of the bottle from which you take your juice. REMINDER! Be sure to record observations in your laboratory notebook. Remember to record your initial and final volumes from the buret in your laboratory notebook. Using the burets properly dispence 2. 0 mL of one of the juices into a clean, dry 250 mL Erlenmeyer flask. Be sure to record the proper number of significant figures for the buret readings.
Use a graduated cylinder to add about 30 mL of distilled water to the flask. Add 1 mL of starch indicator solution, 1 mL of 6 M acetic acid, and 1. 0 g of KI crystals to the flask. [The amounts of starch and acetic acid added do not need to be exactly 1 mL; you can approximate the amounts BUT all values need to be recorded in the laboratory notebook. ] Make up all three of your juice samples at this time. Be sure to label the flasks correctly so that you know the exact concentration of juice in each flask. You can use the untitrated flasks as a reference point during the titration.
Titrate the juice sample with the standardized iodine solution provided in the laboratory. Record the initial and final buret readings and calculate the volume of iodine solution that was used in the titration. Do another titration with a second 2. 0 mL sample of juice. For each titration calculate the number of moles of ascorbic acid in 1 mL of juice. If the calculated numbers have an error of more than 2%, prepare a third sample and titrate it. % Error = moles/mL (trial 1) – moles/mL (trial 2)moles/mL (average) x 100%