INTRODUCTION
Ascorbic acid, commonly called Vitamin C, is a 6-carbon compound that is structurally very similar to the molecule glucose. Vitamin C is a water-soluble vitamin (dissolves in water easily) which is found in many different fruits and vegetables such as oranges, lemons, cabbage, and tomatoes.  This molecule functions in a number of biochemical reactions that involve oxidation (the loss of electrons).  Vitamin C is stable in acid solution but can be rapidly oxidized by oxygen in air when in neutral or alkaline solution. Approximately 3% of the Vitamin C stored in our bodies is lost every day due to oxidation and excretion (I will not define this).  Therefore, for a non-smoking adult a minimum daily requirement of around 60 mg (milligrams) of vitamin C is needed to prevent diseases such as scurvy.  Scurvy was prominent in the past on sailing ships that traveled great distances, due to the lack of fresh produce available during such a long voyage.  British sailing ships would always travel with large barrels of limes that the sailors would eat to help prevent scurvy.  This is where the term "limeys" came from in reference to British sailors.

The results of a ten year study concerning Vitamin C were just released and can be found on the web at the URL below.  I would like you to log onto this site and read through the results of this study.  The pre-lab questions can be answered from these websites.  You should be able to double click on the below URL's or copy and paste them into the web address window of your web browser and hit enter.

http://www.ultimatecitrus.com/pressrel.html
http://www.ultimatecitrus.com/vitaminc.html

In this experiment you will be using a technique called titration to find out exactly how much ascorbic acid is in a sample of fruit juice.  When you titrate a sample you are adding to the sample solution a specific number of moles of a reactant from a buret.  This reactant will combine with your sample, in this case ascorbic acid, and react with it to make the products shown below.

     Ascorbic acid  +  I₂         dehydroascorbic acid   +  2 I¯

A small amount of a 1% starch solution is added to the sample solution to serve as an indicator.  When there are more moles of ascorbic acid in the solution than iodine (I2), all of the I2 will react to form I¯ and the solution will remain light colored.  However, once enough I2 added to the sample solution to react with all of the ascorbic acid, the next drop added will  leave some unreacted I2 in solution.  This free I2 will then react with the starch molecules to give a deep blue or black colored solution. By knowing the volume of the I2 solution added from the buret we will be able to calculate the ascorbic acid content of our unknown sample.  Note that it is important to stop the addition when a single drop of excess iodine has been added (at the very first permanent color change).  The addition of any more iodine will lead to an erroneous result.
 

DEFINITIONS OF TERMS

Molarity - The concentration of the solute in a solution in units of moles of solute per liter of solution (mol/L, abbreviated M).
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Titration – A technique in which a measured volume of a solution of known molarity (the titrant solution, in this case iodine) is added from a buret to a specific volume of a solution of unknown molarity (the analyte, in this case vitamin C).  The reactants then can combine in their stoichiometric proportions allowing us to calculate the molar concentration of the unknown.

Buret - The central piece of equipment in a titration experiment.  It is a long glass tube that is labeled in fractions of milliliters.  It is constructed to deliver precise volumes of a solution through a stopcock valve that you can control precisely.

Titrant - One of the reactants in the chemical reaction.  This reactant of known molar concentration is dispensed from the buret.

Equivalence Point -  When one reactant is slowly added from the buret into the unknown solution there will be a point during the chemical reaction in which there is an equal number of moles of both reactants (for a 1:1 reaction, otherwise a mole ratio corresponding exactly to the stoichiometry of the reaction).  This is called the equivalence point and usually is when the solution being titrated undergoes a color change.

Indicator - A small amount of this substance is added to the unknown solution.  It will change color when the reactants in the titration reach the equivalence point.

In this experiment you will titrate a known volume of fruit juice with an iodine solution of known molarity.  A 1% starch solution will serve as the indicator for this titration. Titrating with iodine is called iodimetry, and at the equivalence point the solution turns a blue-black color which persists for at least 30 seconds.  This color change is due to the adsorption of the excess unreacted iodine onto the starch indicator.  At this point, all of the ascorbic acid has been converted to dehydroascorbic acid.  Your instructor will introduce you to the technique of titration and will teach you how to read a buret correctly.
 

PRE -LAB QUESTIONS

1.  Use your book or the web sites that I have indicated to look up the structure of Vitamin C.  Draw this structure.

2.  What is meant by FCOJ?  What affects the percent of vitamin C content in orange fruits and orange juices (Hint: there are six factors.)

3.  Does the amount of vitamin C in orange juice vary?  Explain.

4.  What kind of juice has the highest amount of vitamin C levels?
 

PROCEDURE:

1.  Each person should carefully clean three Erlenmeyer flasks.  We will work in pairs in this experiment, but each student will have a chance to titrate at least three samples.

2.  Add 100 mL of distilled water to the Erlenmeyer flask.

3.  Fill the buret with the 0.005 M standard iodine solution and RECORD the precise molarity and the initial buret level (to the nearest 0.01 mL).  Remember to keep your eye level with the solution when taking a reading.

4.  Add 10-mL samples of a chosen type of fruit juice into three clean Erlenmeyer flasks.  Record the identity of your fruit juice and the information from the label about the amount of vitamin C present (including the serving size).

5.  Add 10 mL of dilute acetic acid to each fruit juice sample.

6.  Next add 2 or 3 mL of the 1% starch solution to the sample, DO NOT forget this step or the sample solution will never change color.

7.  Slowly add the iodine solution from the buret into your juice sample (swirl the flask as you are adding).   Keep adding iodine solution until the sample turns blue-black color and this color persists for more than 30 seconds.  Be careful; use the correct technique demonstrated to you by your instructor at the beginning of class.  Once the solution turns blue-black and stays that color you have reached the equivalence point of the reaction.

8.  Be sure to RECORD the final buret level to the nearest 0.01 mL.

9.  Add more iodine solution to the buret if necessary and repeat the titration again with a new sample.  (Each student should do at least three titrations).

10.  For calculation purposes, you may assume that the iodine and vitamin C react in a 1:1 stoichiometric ratio.  The molecular mass of vitamin C is 176 g/mole.

11.  Calculate the amount of vitamin C in milligrams that is in each of your samples.   Calculate the average of the three results.  Put the number of milligrams of vitamin C PER SERVING that you found on the board with the rest of the class’s data.
 

FINAL QUESTIONS

1.  How many milligrams of vitamin C did you have in your 10 mL sample of fruit juice?

2.  How many milligrams of vitamin C are there per serving in your fruit juice?  As far as you can tell,  is the manufacturer telling the truth on the label about how much vitamin C is in the fruit juice?

3.  What volume of this fruit juice would you have to consume daily to replenish the daily loss of vitamin C?

4.  Compare all of the different juices listed on the chalkboard.  Which brand has the most vitamin C per serving?  Which brand has the least vitamin C?

5.  What are the possible systematic experimental errors that occurred during your experiment?  How did they affect your data?

6.  Why would this experiment likely be difficult to do with normal, purple grape juice?