Spectroscopic Determination of Iron in a Vitamin Pill Inthisexperimentyouwillanalyzetheamountofironinavitamintablet.Youwillextract
the iron from the vitamin pill and then react it with phenanthroline to form an iron‐tris‐
phenanthrolinecomplex,Fe(phen)3.Thecommonnameforthiscompoundisferroin,andit
is the same compound, which is used as an indicator in many redox titrations. (See, for
example the titration of Fe by Ce in this lab manual.) Ferroin is a red orange compound
whosecolordependsstronglyontheconcentrationofFe2+inthesolution.Thecolorofthe
solution is related to the absorbance of light passing through the solution, and the
absorbanceofthesolutiondependslinearlyontheFe(phen)3concentrationsaccordingto
Beer’s law. The concentration of Fe in the solution can be determined by measuring the
absorbanceofthetestsolutionandcomparingittothemeasuredabsorbancesofaseriesof
standards.
Thestepsinthisanalysisareasfollows:
1. Select a representative sample of multi‐vitamin pills for analysis by first weighing two pills, and then crushing them and selecting a weighed portion of the resulting powder for analysis. This step must be done quantitatively and carefully because it is prone to sampling error. 2. Extract the iron from the powdered sample by dissolving it in hot hydrochloric acid. The dissolved iron will exist in two oxidation states, Fe2+ and Fe3+ . 3. Select a known amount of the sample and convert the Fe to the Fe(phen)3 complex. To do this you must first make sure all of the iron is in the Fe2+ oxidation state by adding a reducing agent. You will use the reaction (hydroquinone + 2Fe3+  quinone + 2Fe2+ + 2H+) to reduce any Fe3+ to Fe2+. Finally you will add phenanthroline to the solution to form the complex. 4. Prepare a series of Fe(phen)3 solutions having known Fe2+ concentrations using ammonium iron sulfate as a primary standard. 5. Measure the absorbance of the sample and standards. 6. Prepare a calibration curve from the standards and use this to determine the concentration of iron in your vitamin pill. Besuretorecordallmassesandvolumesusedinthepreparationofthesesolutionssothat
youcancalculatetheamountofironthatwaspresentintheoriginalvitamintablet!
SpectroscopicDeterminationofIroninaVitaminPill
Procedure (Students will prepare standards in groups of three for this
experiment, but each student will prepare a vitamin pill solution. The
followingprocedureisadequateforvitaminscontainingabout50mgofiron.)
ExtractaSampleofFefromtheVitaminPill
1. Weigh two vitamin tablets (check with your TA for the number for the number of tablet to be analyzed) and record the mass in your notebook. 2. Crush the vitamin tablet with a mortar and pestle to a fine powder. 3. Weigh the mass of the powder for sample preparation directly into a small, dry Erlenmeyer flask. (The mass delivered to the flask should be within a few milligrams of the total pill mass.) Try to avoid getting powder on the side of the flask. Record the mass of the crushed sample in your notebook. Add 25 mL of concentrated hydrochloric acid to the flask and heat it in the hood to near boiling for about 10 minutes. Any powder on the side of the flask should be washed down with a small amount of HCl. Clean the mortar and pestle before proceeding further! 4. Allow the vitamin solution to cool and filter it into a 250 mL volumetric flask. For accuracy you will need this transfer to be quantitative. Make sure to rinse the flask and filter paper well. Rinse the Erlenmeyer flask several times with water and pour the rinse water through the filter into the volumetric flask. When about 50 mL of water has been added, shake the flask to mix the solution. (Prepare your standards while the solution is cooling. Do not dilute to the mark until the solution is at room temperature!) When the solution has cooled to room temperature, fill the volumetric flask to the mark. Mix it by capping the flask and inverting it several times. Label this vitamin solution #1. 5. Turn on the spectrometer light source now so it can warm up and stabilize for your measurements. Prepare Standard Solutions in Groups of three from the (NH4)Fe(SO4)
StockSolution
1. Standards will be prepared from a 40 ppm stock solution of Fe2+ that will be provided. 2. Each group will prepare 5 standards and a blank in six 100 mL volumetric flasks. Using a graduated pipet, deliver 10, 7, 5, 2, 1, and 0 mL of the 40 ppm stock solution into six 100 mL volumetric flasks. Label the flasks 4, 2.8, 2, 0.8, 0.4 and 0 ppm, respectively. 3. Add 8 mL of sodium citrate solution to each flask with the graduated cylinder. This will adjust the pH of the solution to the proper level. 4. Add 2 mL of the hydroquinone solution to each flask with the graduated cylinder. 5. Add 3 mL of the phenanthroline solution to each flask with the graduated cylinder. When you add the phenanthroline the solution should turn orange. SpectroscopicDeterminationofIroninaVitaminPill
6. Dilute each flask to the mark and mix well by capping and inverting the flask. 7. Allow the solution to sit for at least 10 minutes. Noticethatthetreatmentofthestandardsandblankisidenticaltothetreatmentof
thesampleunderanalysis!
PreparetheTris‐phenanthrolineComplexoftheFeintheSample
1. Pipet 1.0 mL of vitamin solution #1 into a second 100 mL volumetric flask. Add about 10 mL of water. Label this as vitamin solution #2. 2. Add 8 mL of sodium citrate to vitamin solution #2 with the graduated cylinder. This will adjust the pH of the solution to the proper level. 3. Add 2 mL of the hydroquinone solution to vitamin solution #2 with the graduated cylinder. 4. Add 3 mL of the phenanthroline solution to vitamin solution #2 with the graduated cylinder. When you add the phenanthroline the solution should turn orange. 5. Dilute to the mark and mix well by capping and inverting the flask. 6. Allow the solution to sit for at least 10 minutes. Measure the Absorbance Spectra of Standards (Including Blank) and
Sample
You will use a diode array spectrophotometer for this measurement. The instrument is
operated using menu driven software. Your lab instructor will assist you with this
measurement.Tomeasureabsorbanceofeachsolution,transferabout3mLofthesolution
toacuvette.
1. Begin by closing shutter and collecting a dark spectrum by clicking on the Dark Spectrum Menu item in the file menu. 2. Next record the reference spectrum. Fill the cuvette with the blank solution and place it in the cuvette holder. Click on the Reference Spectrum menu item in the File menu. 3. Now you can record the spectrum of each standard. Begin with the lowest concentration solution. Empty the cuvette and rinse it once or twice with the solution to be measured next. Fill the cuvette, place it in the cuvette holder, and collect the spectrum. Find the absorbance of the sample at 508 nm for the sample and record it in your lab notebook. Be sure to save the spectrum file and give it an identifiable name. SpectroscopicDeterminationofIroninaVitaminPill
4. Plot standard concentration (x‐axis) versus absorbance (y‐axis) and perform a linear regression analysis on these results. The resulting equation will be Absorbance = mreg x concentration of Fe in ppm + breg (1) where mreg and breg are the regression values for the slope and y‐intercept, respectively. Make sure to record the standard errors of the slope and y‐intercept of the regression line! (See Appendix 3 and 4 for additional details.) You will need to invert this equation to solve for the concentration of your unknown. 5. Each student will need to measure the absorbance of his/her vitamin pill test solution as well. Using the absorbance and equation (1), calculate the concentration of Fe in the solution in units of ppm. Be sure to include an estimate of error in your final result. Propagate error according to instructions in the linear regression handout. Calculations
1. Vitamin solution #1 is 100 times more concentrated than vitamin solution #2. Calculate the mass of Fe in vitamin solution #1 in milligrams using this fact and the measured concentration of vitamin solution #2 and the total volume of vitamin solution #1. 2. Recall that vitamin solution #1 is prepared from a fraction of the original whole vitamin pill, so to determine the mass of Fe in the original vitamin pill you must multiply by the total pill mass and divide by the mass of sample used to make vitamin solution #1. Also report the Fe content of the pill as a percentage of the mass specified on the vitamin bottle. SpectroscopicDeterminationofIroninaVitaminPill
StudentName:
Chemistry 3200 Spectroscopic Determination of Iron in a Vitamin Pill Date: _________
LabInstructor: _________________
_________
Section:
Massofvitaminpill: ___________________
Massusedtopreparevitaminsolution#1:____________________
Iron StandardConcentration
(ppm)
Absorbance
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
RegressionLineValues
Slope:__________
Slopestandarderror:________
y‐Intercept:__________ y‐Interceptstandarderror:________
Vitaminsolution#2absorbance:_________________
Concentrationofironinvitaminsolution#2(ppm):____________________
Massofironinthevitaminpill(mg):____________________
Percentageofthespecifiedironcontentinthevitaminpill(%):_______
AttachthecalibrationplotgeneratedonanExcelspreadsheet.
Indicatetheabsorbanceoftheunknownontheplot.
AttachaprintoutoftheExcelspreadsheetcontainingthelinearregressiondata.
SpectroscopicDeterminationofIroninaVitaminPill
StudentName:
CalculationofthemassofFevitaminsolution#2:
CalculationofthemassofFeinthevitaminpill:
Calculationofthepercentageofthespecifiedironcontentinthe
vitaminpill:
Calculationoftheerrorinthemeasurement.
SpectroscopicDeterminationofIroninaVitaminPill