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Question 1 – iodine solution standardisation

a) Calculate the mean (average) volume of the titre values you have chosen. Justify any

exclusions you have made.

b) What is the number of moles of Vitamin C present in the 25.00mL you pipette into

each conical flask?

c) Using the balanced Equation 2.1, how many moles of iodine, I2, must have been

present in the amount of iodine solution you titrated?

d) Given this number of moles and the average titre value, what is the concentration of

your iodine solution?

Foundations of Chemistry Laboratory Manual VITAMIN C DETERMINATION 3F




1 Thoroughly clean and dry a 250mL beaker and transfer to it roughly 100mL of apple juice. Note the Vitamin C concentration advertised with the nutritional information on the apple juice bottle.

2 Clean and rinse a 250mL conical flask and pipette into it a 25.00mL aliquot of the apple juice.

3 Add 10 drops of starch indicator to this conical flask.

4 Ensure there is a sufficient volume of iodine solution in your burette and that it is set up correctly still from Part One.

5 Perform the titrations until you have a minimum of two concordant titres. Record all your titre values in Table 2.2.

Question 2 – Apple juice investigation

a) Calculate the number of moles of iodine, I2, that was involved during the redox

reaction. (Hint: you are calculating n because you know c and V. What equation

should you use? Refer to Experiment 1F Quantitative Techniques for help.)

b) The equation for the redox reaction between iodine and Vitamin C is provided again

below. Using this balanced equation, how many moles of ascorbic acid (Vitamin C) in

the apple juice reacted with the I2 on average in each titration?

C6H8O6 + I2  C6H6O6 + 2I – + 2H+

c) Given the number of moles of ascorbic acid and original pipette volume, what is the

concentration of Vitamin C in the apple juice you tested?

d) The concentration you calculated in Part c) is in mol.L-1. Convert your concentration