Tuesday
March 28, 2017

Post a New Question

Posted by on .

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)?

  • Chemistry - ,

    Refer to your earlier post.

Answer This Question

First Name:
School Subject:
Answer:

Related Questions

More Related Questions

Post a New Question