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Homework answers / question archive / Calculate the molarity of the ascorbic acid solution: (a) Mass of ascorbic acid used: 0
Calculate the molarity of the ascorbic acid solution:
(a) Mass of ascorbic acid used: 0.1
(b) Moles of ascorbic acid (MW=176.1 g/mol): not sure (0.000568g/moL)
(c) Volume of solution (mL): 100.01mL
(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 iodine concentration, using the formula M1*V1 = M2*V2. Average your results and report one concentration for your standardized iodine solution.
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): .1mL
(b) Concentration of the ascorbic acid in the juice
2. Calculate the average ascorbic acid concentration for the fresh 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):2 drops at a time from burette into a erlenmeyer flask
(b) Concentration of the ascorbic acid in the juice: new orange juice- 8mL of iodine from burette into flask......week old orange juice- 4mL
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 requirement for vitamin C is 60 mg per day. What percentage of this requirement is in one cup (200 mL) of fresh 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.)
Assignment 2 of Procedure 2
Conclusion
Summarize the experiment in one or two paragraphs. Restate the overall purpose of the experiment. What were your results (from each procedure)? What are possible sources of error? What did you enjoy learning about this experiment?
DOES THIS HELP ANY WITH FIGURING OUT THE PROBLEMS?
PLEASE NOTE: Titration requires several steps in order to obtain exact results.
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 from the Glassware shelf and place it on the workbench.
2. Add 0.1g ascorbic acid to the volumetric flask.
3. 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.
4. Take a 150 Erlenmeyer flask from the Glassware shelf and place it on the workbench.
5. Pour 20 mL of the ascorbic acid solution into the flask.
6. Add 1 mL of the starch indicator to the flask.
7. Take a burette from the Glassware shelf and place it on the workbench.
8. Fill the burette with 50 mL iodine solution, with an aproximately known concentration of 0.015M.
9. Drag the Erlenmeyer flask and drop it on the lower half of the burette.
10. Open the Properties window, click back on the burette, and then click the Pushpin icon in the blue bar of the Properties window to lock its functions to the burette.
11. Titrate the ascorbic acid by adding iodine until the solution in the Erlenmeyer flask turns dark blue.
12. You should do a 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. You can then do a fine titration one or two times in order to obtain exact results.
13. When standardizing a solution, you should repeat the titration so that your results are within a few drops of each other. The results are then averaged to determine the precise molarity.
Procedure 2 ( click to view assignments for this procedure )
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 juice from the new container by adding 40 mL of the juice to a clean Erlenmeyer flask. Add 1 mL of starch indicator.
(Please excuse the very light color of the OJ - the pulp was REALLY filtered out)
3. Titrate the orange juice with the 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.
4. Repeat steps 2-3 with the week-old orange juice.
Molarity of ascorbic acid solution:
0.1 g x 1 mole/ 176.1 g = 0.000568 moles (your answer was correct but units are moles not grams/mole)
0.000568 moles/ 0.10001 L = 0.00568 mol/L
For the standardization of the iodine solution, you have not given the information for the number of mL of iodine solution added in the titration of 20 mL of a 0.00568 mol/L solution of ascorbic acid (calculated above). I will thus give you the generalized method for determining the concentration of your standard iodine solution from this titration:
The idea of any titration is that you are reacting two substances together in order to determine the exact point at which a reaction between them is completed. This, along with the stoichiometry of the known reaction, can give information about an unknown amount of one of the reactants given the addition of a known quantity of the other reactant. In this case, you are looking at the oxidation-reduction reaction between iodine and ascorbic acid... the ascorbic acid is oxidized by the iodine until it is completely oxidized to dehydroascorbic acid, at which point the excess iodine added will react with the starch to produce a blue colour, indicating the endpoint of the reaction. Since you know how many moles of ascorbic acid have been added in the first procedure (0.00568 mol/L x 0.020 L = 1.14 x 10-4 moles), you can calculate the concentration of your standard iodine solution if you know the volume of iodine solution added:
The stoichiometry of the reaction (probably given in your lab manual) shows that 2 electrons are released for every molecule of ascorbic acid reduced, and that these 2 electrons are required to oxidize one I2 molecule, so there is a one-to-one ratio between the number of moles of ascorbic acid to moles of iodine in this reaction (making calculations very easy).
To calculate the concentration of your iodine solution:
You have given two concentrations for the iodine standard - 0.0015 M and 0.015 M; I will assume the first one is a mistake and use 0.015 M (moles/L).
Also, the number of mL of iodine added in the standardization titration are not given anywhere that I can see, I will give you the calculation using VI to represent this value. You know that at the endpoint of the titration, the number of moles of iodine is equal to the number of moles of ascorbic acid, and that you have added 0.000114 moles ascorbic acid in 20 mL of solution initially to the flask. There are then 0.000114 moles of iodine at the endpoint, when the solution turns blue. You therefore know that there are 0.000114 moles of iodine in whichever volume VI you have added from the buret, because at this point the reaction is complete and the starch indicator turns the solution blue.
0.000114 moles iodine / VI = molarity of iodine solution in the buret. If the initial iodine concentration was supposed to be around 0.015 mol/L, you would have needed somewhere around 7.6 mL of iodine to reach the endpoint in the first standardization titration.
In the titrations of the fresh and week-old orange juice, you have given values of 8 mL and 4 mL, respectively for the average volumes of iodine solution added. We will assume that the standardization titration confirmed 0.015 mol/L for the iodine solution concentration.
To calculate the ascorbic acid in the orange juice, use the following:
0.015 mol/L iodine x 0.008 L = 0.00012 moles x 1 mole ascorbic acid/1 mole iodine x 176.1 g/mole = 0.0211 g ascorbic acid in fresh orange juice.
0.015 mol/L iodine x 0.004 L = 0.00006 moles x 1 mole ascorbic acid/1 mole iodine x 176.1 g/mole = 0.0106 g ascorbic acid in week old orange juice.
This indicates that there is two times less ascorbic acid in orange juice after exposure to air for one week under refrigeration. The oxygen in the air is a powerful oxidizer - this should tell you what has happened to the ascorbic acid over time.
If the daily intake is recommended to be 60 mg/day, and there is 0.0211 g 21.1 mg in the 40 mL of fresh orange juice added to the flask for the orange juice titrations, there would be:
21.1 mg/40 mL = x mg / 200 mL = 105.5 mg in one 200 mL glass of fresh orange juice. This is 105.5/60 x 100 = 175.8 % of the recommended daily allowance. The amount in the week old orange juice is calculated the same except you put in 10.6 mg in the calculation.
