A student has 100.0 mL of an unknown solution of Vitamin C. She removes 1.00 mL of this solution, dilutes it to 50.0 mL, and titrates this second solution with an iodine solution that is 8.5 × 10-4 M. 6.60 mL of iodine solution is required to reach the endpoint.

1. How many moles of iodine were added to the vitamin C solution to reach the endpoint of this titration?

2. If at the endpoint, the moles of vitamin C and the moles of iodine are equal, how many milligrams of Vitamin C were in the original 100.0 mL?

3. What is the molarity of the original unknown vitamin C solution (prior to dilution)?

http://www.jiskha.com/display.cgi?id=1331598369

To answer these questions, we need to follow a step-by-step approach using the given information and some relevant equations.

1. How many moles of iodine were added to the vitamin C solution to reach the endpoint of this titration?

We know the molarity and volume of the iodine solution used in the titration.
Number of moles = Molarity × Volume (in liters)

First, we need to convert the volume of iodine solution used in the titration to liters:
Volume of iodine solution used = 6.60 mL = 6.60 × 10^(-3) L

Now we can calculate the number of moles of iodine:
Number of moles of iodine = (8.5 × 10^(-4) mol/L) × (6.60 × 10^(-3) L)

2. If at the endpoint, the moles of vitamin C and the moles of iodine are equal, how many milligrams of Vitamin C were in the original 100.0 mL?

Since the moles of vitamin C and iodine are equal at the equivalence point, we can use the mole ratio between vitamin C and iodine to determine the number of moles of vitamin C.

The balanced chemical equation for the reaction between iodine and vitamin C is:
Vitamin C + I2 → Dehydroascorbic Acid + 2I-

The stoichiometry of the reaction tells us that 1 mole of vitamin C reacts with 1 mole of iodine. So, the number of moles of vitamin C is the same as the number of moles of iodine added.

Using the calculated number of moles of iodine, we can get the number of moles of vitamin C.

3. What is the molarity of the original unknown vitamin C solution (prior to dilution)?

We can use the volume and concentration of the diluted solution to find the concentration of the original solution. The diluted solution was prepared by taking 1.00 mL of the original solution and diluting it to 50.0 mL. So, the diluted solution has a concentration that is 1/50th of the concentration of the original solution.

If we let M1 be the molarity of the original solution, we have:
(M1 × 1.00 mL) / 100.0 mL = (8.5 × 10^(-4) M) / 50.0 mL

This allows us to solve for M1, which represents the molarity of the original unknown vitamin C solution before dilution.

By following these steps, we can find the answers to the three questions.