Calculating Moles of IO3- in the Erlenmeyer Flask?

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To calculate the number of moles of IO3- in the Erlenmeyer flask, first determine the moles of KIO3 used, which is 1.6635 × 10^-5 moles based on the given mass and molecular weight. Since the mole ratio is 1:1, the moles of IO3- will be equal to the moles of KIO3. The calculated molarity of the solution is 1.6635 × 10^-4 moles/L. To find the moles of IO3- in the flask, multiply the molarity by the volume of the KIO3 solution transferred, which is 10 mL, not the total volume of 64 mL. This clarification resolves the confusion regarding the volume to use for the calculation.
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Homework Statement



Calculation of the number of moles of IO3- used in the titration(i.e. in the Erlenmeyer flask):
KIO3(s) → K+(aq) + IO3-(aq): balanced eqn of dissolving KIO3 in water

1. Calculate the moles of KIO3 used: Molecular weight of KIO3 = 214.0011 g/mol
Amount of KIO3 used in grams = 0.0356g
# of moles = 0.0356g/(214.0011g/mol) (grams cancels out) = 1.6635 × 10^-5 moles
The mole ratio is 1:1:1. So # of moles of KIO3 will be equal to # of moles of IO3-

2. Calculate the Molarity
Amount of dH2O used for dissolving = 100mL = 0.1L
Molarity = (moles of solute) / (volume of solution in Liters) = 1.6635 × 10^-5 moles/0.1L
= 1.6635 × 10^-4 moles/L
I did until this. But in the question, it says that <Calculate the # of moles of IO3- in the Erlenmeyer flask> And in the Erlenmeyer flask, there was 10mL of KIO3 solution, and other solution of 54ml.

And my question is, do I have to multiply the molarity I gained by calculation by 0.064mL (because (10ml+54ml)/1000 = 0.064L=amount in Erlenmeyer flask) to gain the # of moles of IO3- in the Erlenmeyer flask, Or can I just multiply the molarity by 10mL Or am I doing wrong??
Please help!
 
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10 mL - that was the amount of solution of iodate transferred, right? You have not transferred 64 mL, or any other volume.

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I get it. I was confused when my lab TA emphasized "in the Erlenmeyer flask", so i thought she might mean something else. :)
 
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