Using a 1000 ppm stock Fe solution. I am supposed to prepare the stock solution
ID: 1043396 • Letter: U
Question
Using a 1000 ppm stock Fe solution. I am supposed to prepare the stock solution by weighing accurately 0.72 ferric nitrate nonahydrate, Fe(No3)3 9 H2O and dissolving in 100.0 ml DI water. Calculate the actual concentration of Fe in the stock soluton from the mass of ferric nitrate used. How do I calculate this?
Then I need to take 10 ml of the stock solution from which we need to prepare five standards with concentrations of 50, 75, 100, 150 and 200 ppm Fe. I have to calculate the actual concentrations of the standard solutions based on the concentration of the stock solution. How do I calculate this?
Explanation / Answer
Ans. # Step 1: Moles of Fe(NO3)3.9H2O = Mass / MW
= 0.72 g / (403.99934 g/ mol)
= 0.0017822 mol
# Since 1 mol Fe(NO3)3.9H2O has 1 mol Fe-atom-
Moles of Fe taken = 0.0017822 mol
Mass of Fe taken = Moles x MW = 0.0017822 mol x (55.847 g/ mol)
= 0.0998295 g
= 99.8295 mg
# Now, [Fe],ppm = Mass of Fe in mg / Volume of solution in liters
= 99.8295 mg / 0.100 L
= 998.295 ppm
# Step 2: You did NOT mentioned the final volume of each standard aliquots to be made.
# One example of standard aliquot preparation is provided below. You can do other calculations depending on the required volume of standard aliquots.
# Aliquot 1: 50 ppm : Let the required volume of standard aliquot be 10.0 mL.
Using C1V1 (stock solution) = C2V2 (standard aliquot)
Or, 998.295 ppm x V1 = 50.0 ppm x 10.0 mL
Or, V1 = (50.0 ppm x 10.0 mL) / 998.295 ppm = 0.501 mL
Therefore, required volume of stock solution = 0.501 mL
Aliquot preparation: Transfer 0.501 mL (or 0.50 mL) to a standard 10.0-mL standard volumetric flask. Make the final volume upto the mark with distilled water. The resultant solution is 50.0 ppm Fe.
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