Question

B. Preparation of the unknown sample Accurately weigh by difference to 3 decimal places about 0.2 g of potassium iron (III) oxalate into a 100 mL beaker. Add about 50 mL of water, and stir to dissolve the solid. Transfer this solution to a clean 250.0 mL volumetric flask (rinsing the beaker and stirring rod as necessary with small portions of water), and make up to exactly 250.0 mL. This solution is your original unknown solution. After mixing, transfer the original unknown solution to a clean DRY beaker. Rinse out the volumetric flask. Using a pipet, deliver exactly 2.00 mL of the original unknown solution in the beaker into the rinsed 250.0 mL volumetric flask. Add the reagents as described in Part A, using the same quantities of ammonium acetate, hydroquinone and o-phenanthroline. Make to 250.0 mL and mix well. Allow 10 minutes for colour development. This is the diluted unknown solution. C. Preparation of the calibration graph Prepare a series of sample holders by cleaning out regular 18 mm x 150 mm test tubes. Rinse these twice with deionized water and twice with the solution that it will contain. Fill these test tubes about one-third full with the blank, the standards and the diluted unknown samples respectively. Once the colour has developed for 10 minutes, ask your laboratory instructor to calibrate the spectrophotometer using the blank solution, record this reading in your data table and then determine and record the absorbance values of the rest of the solutions and samples on the spectrophotometer at 490 nm. Note that the absorbance scale is the bottom scale, it is an exponential (logarithmic), NOT a linear, scale and that Absorbance has no units. Each of you will need to record the absorbance of the blank and standard solutions but only the absorbance of your own unknown sample. CALCULATIONS 1. Use the given value for the concentration of iron in the stock solution and the volume you used to calculate the concentration of Iron(II) ions (Fe2+) in mg/L in the blank and each of the four standard solutions. 2. Convert the % Transmittance of each solution to absorbance. [A=2 – log(%T)] 3. Present the concentration of Iron(II) ions (Fe2+) in mg/L and absorbance of the blank and standards in a table for graphing. 4. Construct a calibration graph from the above values and draw the best-fitting straight line through the data points. Determine the concentration of iron with its unit in the diluted unknown solution on the calibration graph. 5. From the diluted unknown concentration, calculate the concentration of iron in the original unknown solution. 6. Calculate the mass of iron in the solid unknown sample. 7. Calculate the percentage of iron in the unknown sample where: % iron in unknown sample = (mass of iron in unknown sample / total mass of unknown sample) X 100% 8. Calculate the theoretical % Fe in one mole of K3Fe(C2O4)3 • 3H2O. Is your result from Question 7 higher or lower than the theoretical value?

B. Preparation of the unknown sample
Accurately weigh by difference to 3 decimal places about 0.2 g of potassium iron (III) oxalate into a 100 mL beaker. Add about 50 mL of water, and stir to dissolve the solid. Transfer this solution to a clean 250.0 mL volumetric flask (rinsing the beaker and stirring rod as necessary with small portions of water), and make up to exactly 250.0 mL. This solution is your original unknown solution.
After mixing, transfer the original unknown solution to a clean DRY beaker. Rinse out the volumetric flask. Using a pipet, deliver exactly 2.00 mL of the original unknown solution in the beaker into the rinsed 250.0 mL volumetric flask. Add the reagents as described in Part A, using the same quantities of ammonium acetate, hydroquinone and o-phenanthroline. Make to 250.0 mL and mix well. Allow 10 minutes for colour development. This is the diluted unknown solution.
C. Preparation of the calibration graph
Prepare a series of sample holders by cleaning out regular 18 mm x 150 mm test tubes. Rinse these twice with deionized water and twice with the solution that it will contain. Fill these test tubes about one-third full with the blank, the standards and the diluted unknown samples respectively.
Once the colour has developed for 10 minutes, ask your laboratory instructor to calibrate the spectrophotometer using the blank solution, record this reading in your data table and then determine and record the absorbance values of the rest of the solutions and samples on the spectrophotometer at 490 nm. Note that the absorbance scale is the bottom scale, it is an exponential (logarithmic), NOT a linear, scale and that Absorbance has no units.
Each of you will need to record the absorbance of the blank and standard solutions but only the absorbance of your own unknown sample.
CALCULATIONS
1. Use the given value for the concentration of iron in the stock solution and the volume you used to calculate the concentration of Iron(II) ions (Fe2+) in mg/L in the blank and each of the four standard solutions.
2. Convert the % Transmittance of each solution to absorbance. [A=2 – log(%T)]
3. Present the concentration of Iron(II) ions (Fe2+) in mg/L and absorbance of the blank and standards in a table for graphing.
4. Construct a calibration graph from the above values and draw the best-fitting straight line through the data points. Determine the concentration of iron with its unit in the diluted unknown solution on the calibration graph.
5. From the diluted unknown concentration, calculate the concentration of iron in the original unknown solution.
6. Calculate the mass of iron in the solid unknown sample.
7. Calculate the percentage of iron in the unknown sample where:
% iron in unknown sample = (mass of iron in unknown sample / total mass of unknown sample) X 100%
8. Calculate the theoretical % Fe in one mole of K3Fe(C2O4)3 • 3H2O. Is your result from Question 7 higher or lower than the theoretical value?

Added by Adam S.

Chemistry: Structure and Properties

Nivaldo Tro 2nd Edition

Instant Answer

Solved by Expert Sri K

11/14/2022

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Step 1: Calculate the concentration of Fe2+ ions in the blank and each of the four standard solutions using the given value for the concentration of iron in the stock solution and the volume used. Show more…

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B. Preparation of the unknown sample Accurately weigh by difference to 3 decimal places about 0.2 g of potassium iron (III) oxalate into a 100 mL beaker. Add about 50 mL of water, and stir to dissolve the solid. Transfer this solution to a clean 250.0 mL volumetric flask (rinsing the beaker and stirring rod as necessary with small portions of water), and make up to exactly 250.0 mL. This solution is your original unknown solution. After mixing, transfer the original unknown solution to a clean DRY beaker. Rinse out the volumetric flask. Using a pipet, deliver exactly 2.00 mL of the original unknown solution in the beaker into the rinsed 250.0 mL volumetric flask. Add the reagents as described in Part A, using the same quantities of ammonium acetate, hydroquinone and o-phenanthroline. Make to 250.0 mL and mix well. Allow 10 minutes for colour development. This is the diluted unknown solution. C. Preparation of the calibration graph Prepare a series of sample holders by cleaning out regular 18 mm x 150 mm test tubes. Rinse these twice with deionized water and twice with the solution that it will contain. Fill these test tubes about one-third full with the blank, the standards and the diluted unknown samples respectively. Once the colour has developed for 10 minutes, ask your laboratory instructor to calibrate the spectrophotometer using the blank solution, record this reading in your data table and then determine and record the absorbance values of the rest of the solutions and samples on the spectrophotometer at 490 nm. Note that the absorbance scale is the bottom scale, it is an exponential (logarithmic), NOT a linear, scale and that Absorbance has no units. Each of you will need to record the absorbance of the blank and standard solutions but only the absorbance of your own unknown sample. CALCULATIONS 1. Use the given value for the concentration of iron in the stock solution and the volume you used to calculate the concentration of Iron(II) ions (Fe2+) in mg/L in the blank and each of the four standard solutions. 2. Convert the % Transmittance of each solution to absorbance. [A=2 – log(%T)] 3. Present the concentration of Iron(II) ions (Fe2+) in mg/L and absorbance of the blank and standards in a table for graphing. 4. Construct a calibration graph from the above values and draw the best-fitting straight line through the data points. Determine the concentration of iron with its unit in the diluted unknown solution on the calibration graph. 5. From the diluted unknown concentration, calculate the concentration of iron in the original unknown solution. 6. Calculate the mass of iron in the solid unknown sample. 7. Calculate the percentage of iron in the unknown sample where: % iron in unknown sample = (mass of iron in unknown sample / total mass of unknown sample) X 100% 8. Calculate the theoretical % Fe in one mole of K3Fe(C2O4)3 • 3H2O. Is your result from Question 7 higher or lower than the theoretical value?