Determination of ascorbic acid (vitamin C).
Titrations involving iodine are widely used in analytical chemistry. I2 is only slightly soluble in water. Its solubility can be increased by complexation with iodide:
I2 + I- ® I3-
Therefore, when we speak of using iodine as a titrant, we mean that we are using solution of  I3- plus excess of I-. The yellow-orange color of iodine solutions is not intensive enough to detect the excess of iodine. To enhance the color, we use starch. It forms an intense blue complex with iodine. Very small amounts of iodine can be detected with starch, making quantitative titration possible.
Reducing agents such as ascorbic acid (vitamin C) are titrated directly with standard iodine solution in the presence of starch, until reaching the intense blue end point:

One mole of ascorbic acid (F.W.= 176.13) reacts with one mole of iodine.
Procedure:
You will need a standard solution of iodine. Fill the buret with this solution and write its concentration from the label in your notebook.
Your unknown solution is placed in 100-mL volumetric flask. Dilute the unknown to the mark with deionized water. Using a 25 mL pipet, transfer an aliquot part of the solution into an Erlenmeyer flask. Add some colorless starch solution, and titrate until the intense blue color.
Repeat the same titration two times more.

Calculations:
To obtain the mass of ascorbic acid in unknown, you need to find molarity of the ascorbic acid:
CHA = CI2 x VI2/ V HA
mHA = FWHA x C HA x (total volume)
In your case, (total volume)=100 mL= 0.100 L and FW=176.13 g/mol.
Therefore,
mHA = 176.13 x0.100 x CI2 x VI2/V HA  = 17.613 x CI2 x VI2/V HA
Report

Concentration of iodine (copy from the label!): __________
Volume of your pipet _____25.00 mL
Volume of iodine used for titration

# 1 ___________    # 2______________  # 3_____________( mL)

Mass of ascorbic acid
# 1 ___________    # 2______________  # 3______________(g)

Average  mass_____________ g 
Standard deviation _____________g