Posted by **Cassandra** on Monday, March 12, 2012 at 8:28pm.

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 solution that 8.5x1^-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 end point, 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)?

## Answer this Question

## Related Questions

Chemistry - A student has 100.0 mL of an unknown solution of Vitamin C. She ...

Chemistry - A student has 100.0 mL of an unknown solution of Vitamin C. She ...

Chemistry - a potassium dichromate is prepared by 0.3525grams of K2Cr2O7 is ...

Chemistry - A sample of fresh grapefruit juice was filtered and titrated with ...

AP Chemistry - A student titrates 25.0 mL of a 0.100 M solution of acetic acid ...

Chemistry - A student dissolves 22.4 g of sodium phosphate to prepare a 2.98 L ...

Math - A druggist mixes a 10% solution of iodine with a 15% solution of iodine. ...

Chemistry - I am having trouble answering this. please help! Q: Cereol is a ...

CHEMISTRY/ - A student takes a 1.00 mL aliquout of 5.00*10^-4 M HCl solution and...

Chemistry - To illustrate the calculation, assume that 21.95 mL of a thiosulfate...