Just a coupe of uni level chem questions! If you have any issues especially in the second one, read the uploaded document to get an idea.
Question 1 – iodine solution standardisation:
a) Calculate the mean (average) volume of the titre values you have chosen. Justify any exclusions you have made. [Average is found to be: 0.74 Litres]
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 C6H8O6 + I2 = C6H6O6 + 2I- + 2H+ , 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?
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? [n = CxV]
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 from Part c) to g.L-1 (hint: What equation relates number of moles and mass?)
e) Convert the concentration from Part d) to milligrams per litre (mg.L-1 )
f) Now convert the concentration from Part e) to milligrams per 100 ml.
g) How does your experimentally determined Vitamin C concentration compare with the value given on the juice bottle?
h) List the experimental errors that could lead to a discrepancy between the determined and the advertised value.
i) Considering your experimentally determined value and the possible sources of error, make a comment about the accuracy of the advertised amount of Vitamin C present.
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