Spectrophotometric Iron Fall 2010.doc

Revised Fall, 2010 P. W. Crawford CH271 SPECTROPHOTOMETRIC DETERMINATION OF IRON WITH ORTHOPHENANTHROLINE Introduction Spectrophotometric methods exist for the quantitative determination of practically every element on the periodic chart as well as a large number of compounds. Many standard methods of analysis approved by the FDA, EPA, and clinical diagnostic work depend on the selective absorption of visible radiation by any ionic or molecular species which exhibits a characteristic color in solution. This analytical technique dates back to 1838 when it was called “colorimetry-or measuring color.” In most methods of absorption spectrophotometry the absorption spectrum provides a “finger print” of the chemical species for qualitative purposes. Beer’s Law developed in 1852 establishes the quantitative relationship between the amount of radiation absorbed and the concentration of the absorbing chemical species. Beer’s law may be expressed as A = bc where A is absorbance,  is the molar absorptivity (M-1cm-1), b is pathlength (cm), and c is molar concentration. Fundamentally, applying Beer’s Law for quantitative analysis requires only two absorbance readings, one for a standard solution and one for the unknown solution. The ratio of the 2 absorbance readings can then be used to calculate the concentration of the unknown. Assuming the same pathlength for each measurement, Aunk/Ast = cunk/cst Although an unknown solution’s concentration may be determined very quickly using this method, this is not necessarily considered to be the most reliable way to measure unknown concentrations. (Why not?) A preferred method involves the use of a Beer’s Law plot, i.e. a plot of absorbance versus concentration. Such a plot is prepared by measuring the absorbances for a series of standard solutions. The slope equals b. Upon measuring the absorbance of the unknown solution, the plot may be used to determine the concentration of an unknown solution. An excellent and sensitive method for the determination of iron is based on the formation of the orange-red Fe (II) - orthophenanthroline complex. The orthophenanthroline molecule is a weak base: in acidic solution the principal species is the phenanthrolium ion PhH+. Thus, the complex formation reaction is best described by Fe2+ + 3PhH+  Fe(Ph)32+ + 3H+

Kstab = 1 X 1018 at 25C at pH = 4.0

Clearly, the position of the equilibrium is pH dependent. Below pH of 2, the reaction is incomplete. Quantitative formation is obtained between a pH of 3 to 9. A pH of 4.0 is best to prevent the precipitation of various iron salts such as phosphates which might cause problems with some samples.

Excess reducing agent is added to keep the iron in the +2 oxidation state. The complex itself is stable for long periods of time once it is formed. As with most colorimetric methods, many other ions may also form complexes with orthophenanthroline and thus cause erroneous results unless properly dealt with. The purpose of the experiment is to illustrate the procedures typically involved in the analytical application of visible spectrophotometry. Key Terms Used in Absorption Spectroscopy Absorption Spectrum: A plot of the absorption characteristics of an analyte as a function of wavelength. Typically one plots absorbance vs wavelength,  (in nanometers). This plot provides a fingerprint for qualitative identification of the analyte. The plot can be used to determine max (the wavelength at which the analyte absorbs the most light) and the molar absorptivity constant for the analyte, , or the absorptivity constant for the analyte, a. Beer's Law: Developed in 1852. States that the absorbance of light is directly proportional to the concentration (c) of the absorbing species when the path length (b) of the radiation in the absorbing medium is fixed. A = abc A = bc

where a = absorptivity constant where  = molar absorptivity constant (c given in mol/L or mmol/mL).

Beer's Law Plot: Calibration curve that is a plot of absorbance at max vs. concentration for a series of standard solutions of known concentration. This plot should generate a straight line. Thus, the concentration of an unknown can be determined using the plot or calculated from the mathematical equation which describes a straight line (mx + b = y). The latter is the better way to do this. Chemicals and Equipment 1. 2. 3. 4. 5. 6. 7. 8. 9.

Spectrophotometer Several 100 mL and 1-500 mL volumetric flasks with corks or lids 10.00 and 5.00 volumetric pipets 1 % Hydroxylamine HCl solution buffered with sodium acetate to pH = 4.0 Sodium acetate - acetic acid buffer at pH = 4.0 0.2 % 1,10-phenanthroline solution 3M H2SO4 0.0200 mg/mL (20.0 ppm) standard iron solution Unknown: Determine % Fe in solid ferrous ammonium sulfate sample

A.

Procedure Determination of the Absorption Spectrum, max and the Molar Absorptivity Constant, , for the Fe-orthophenanthroline complex.

1. Prepare a 0.2 % solution of 1,10-phenanthroline in a 100.0 mL volumetric flask. Weigh out 0.2 grams of 1,10-phenanthroline and transfer it to a volumetric flask. Add 75.0 mL of distilled water to the flask, and then use a stir bar and magnetic stirrer to dissolve the compound. Remove the stir bar, and then dilute the solution up to the mark on the volumetric flask. You will need to stir this solution. 2. Preparation of a standard Fe(Ph)32+ solution: Measure 10.00 mL of the standard Fe solution (0.0200 mg/mL) into a 100 mL volumetric flask, add 10.0 mL of pH = 4 buffer, 5 mL of hydroxylamine HCl, and 10.00 mL of phenanthroline solution. Dilute the solution in the flask with distilled H2O to exactly 100.0 mL, and mix throughly. Allow 15 minutes after adding all the reagents so that the color of the complex can fully develop before making any absorption measurements. Calculate the concentration (in ppm, M, and mg/mL) of the standard Fe solution prepared. Note that for dilutions, you need to use the following equation to calculate the new concentration: C1V1 = C2V2. 3. Absorption spectrum for the Fe(Ph)32+ standard solution: Using the spectronic D, measure and record the absorbance of the standard from 400 to 600 nm. Take readings at 25 nm intervals except near the vicinity of the absorption maximum where you should take readings at 4-5 nm intervals. CAUTION: For a single beam instrument you MUST reset 100% T (or 0.00 absorbance) each time you change the wavelength using the H2O solvent blank. Plot the absorbance spectrum, i.e. A on the y-axis vs.  on the x-axis, using Microsoft Excel. Be sure to clearly label your axes. Using the Beckman photodiode array UV-Vis spectrophotometer in REDI scan mode, obtain the absorbance spectrum of the standard from 400 to 600 nm. (Ask the instructor for instructions.) Be sure to scan the blank containing distilled H2O before running your spectrum. Record the max value and the absorbance at this wavelength for Fe(Ph)32+. max values as well as absorbance values at any wavelength can be obtained using the Trace and Auto Scale functions of the instrument. 4. Determination of Molar Absorptivity Constant, , for Fe(Ph)32+: Use the molar concentration of the standard Fe solution and its absorbance at max to calculate the molar absorptivity constant, , for Fe(Ph)32+. (Do this for both instrument readings.) Assume a 1.00 cm cell path length. The literature value for  is 1 x 104 cm-1 M-1. What is the relative % error for your experimental value of ? Which instrument gives a more accurate value for ? B.

Quantitative Analysis 1. Preparation of Unknown: Obtain a soluble iron sample unknown. Do not oven dry, but do store in a dessicator. Weigh accurately about 0.20 g of the unknown into a 500 mL volumetric flask. Add about 100 mL H2O to dissolve the sample, add 5 mL of 3 M H2SO4 and dilute with H2O to exactly 500 mL. Mix the solution thoroughly. 2. Prepare 4 additional standard solutions of Fe(Ph)32+: You already have a 0.0020 mg/mL standard Fe(Ph)32+ solution using 10.00 mL of 0.020 mg/mL Fe solution and other reagents specified by the instructions in A.2. Make 4 additional standard solutions exactly as was done in procedure A.2. except using 5.00 mL, 15.00 mL, 20.00 mL, and 25.00 mL of the standard 0.020 mg/mL Fe solution. The amounts of the other reagents should be

exactly the same as used in A.2. Calculate the final concentration of each of the above standard solutions in mg/mL of Fe. 3. Prepare 4 unknown Fe sample solutions: Prepare 4 unknown Fe(Ph)32+ solutions exactly as done in A.2. except use 5.00, 7.50, 10.00, and 12.50 mL of your unknown Fe sample solution. The amounts of the other reagents should be exactly the same as used in A.2. Use a plastic 10 mL Mohr pipet (CAUTION: This pipet operates as a handheld/controlled buret. Do not drain below the 10.00 mL calibration mark.) The color intensity of all of these solutions should be greater than that of the least concentrated standard Fe(Ph)32+ solution, but less than that of the most concentrated standard solution. 4. Measure and record the absorbance of all five standard Fe(Ph)32+ solutions and the 4 unknown Fe(Ph)32+ solutions at max using a distilled H2O blank. 5. Quantitative Data Treatment and Calculations a. Beer's Law Plot: Using Microsoft Excel, plot the absorbance (Y axis) versus concentration (mg/mL) of the standard Fe(Ph)32+ solutions (X-axis). You should obtain a straight line. Use your Beer's law plot to determine the concentration (mg/mL of Fe) of each of the four unknown Fe(Ph)32+ solutions. Keep in mind that the test unknown solutions were prepared via dilution of your stock solution. Thus, you will need to use the equation: C1V1 = C2V2 to get the original concentration of your unknown solution. b. Determine the % Fe, standard deviation, and ppt relative error in your soluble Fe salt unknown based on your four measured solutions. Note: About 10 ppt RSD error is considered very good for most instrumental methods of analysis.

Example Calculations There are several ways to treat the data to get the % Fe in the sample. The following is perhaps the easiest treatment. A student weighed 300.0 mg of unknown sample into exactly 500.0 mL of solution using a volumetric flask. A 10.00 mL aliquot of this solution was treated with the appropriate reagents to form Fe(Ph)32+ and diluted to exactly 100.0 mL. This solution had an absorbance of 1.00 and was found to contain 3.60 x 10-3 mg/mL of Fe from the student's Beer's Law Plot. Original Unknown Sample Solution: A 300.0 mg unknown sample in 500.0 mL total volume contains 0.600 mg/mL sample Measured solution: Contains 3.60 x 10-3 mg/mL Fe in 100.0 mL volume. 100.0 mL total x 3.6 x 10-3 mg/mL Fe = 0.360 mg Fe in 100.0 mL volumetric flask, which came from 10.00 mL of the original sample solution. original sample soln. contains 0.360 mg Fe in 10.00 mL = 3.60 x 10-2 mg/mL Fe In other words, using C1V1 = C2V2: C1 = (100.0 mL x 0.00360 mg/mL)/10.00 mL = 0.0360 mg/mL Percent Fe in sample: % Fe = (0.0360 mg/mL Fe x 100)/0.600 mg/mL sample = 6.00% Important Notes NOTE 1. You may work in pairs, but each will have an individual unknown. Each will turn in their own complete lab report, including data, graphs, and calculations for parts A and B. Note 2. Beer's Law data can be used to determine the mathematical equation of a straight line from your plot (graph of A vs. Conc.), or from a manual linear least squares analysis of the data or a computer linear least squares analysis of your data. All of the above reduce the error associated with obtaining the concentration of the unknown iron solutions directly from the graph. SEE LAST PRE-LAB PROBLEM. Note 3. Save your standard solutions for the next part of the experiment that follows: Determination of Soluble Iron in a Commercial Iron Supplement Tablet/Pill/Capsule.

CH271 Pre-Lab Problems Spectrophotometric Detn. of Iron 1. Graph absorbance vs  from the data provided to determine max of the intensely yellow FeCl63complex. The concentration of Fe (FW = 55.85) in the solution measured was 5.6 x 10-3 mg/mL. Calculate the molar absorptivity constant for the complex. Assume that b, pathlength, is 1.00 cm. Note: The molar concentration of Fe equals the molar concentration of the complex since all the iron is bound to chloride. nm 400 420 440 450 460 470 480 500 520 540 560 580 600

Absorbance 0.09 0.25 0.48 0.55 0.60 0.62 0.55 0.37 0.28 0.18 0.05 0.00 0.00

Ans.:  max is about 470 nm,  is about 6200 cm-1 M-1 2. 12.50 mL of a standard Fe2+ solution of 0.020 mg/mL was converted to the colored iron orthophenantroline complex in a 100.0 mL volumetric flask as directed in the experiment. This solution gave an absorbance of 0.90 a. What is the concentration in mg/mL of Fe solution prepared? Ans. 2.50 x 10-3 mg/mL b. What absorbance would you expect to get from a second Fe2+ solution prepared as directed with a concentration of 4.00 x 10-3 mg/mL. Ans. 1.44 c. If 7.50 mLs of a 0.020 mg/mL Fe2+ solution was prepared as directed, what absorbance would you expect to get? Ans. 0.54

3. Using the 0.020 mg/mL standard solution of Fe2+ as directed in the experiment, a student got the following calibration curve data. mLs of standard used - diluted to 100.0 mL 0.00 5.00 10.00 15.00 20.00 25.00

Concentration (mg/mL) 0.0 1.0 x 10-3 2.0 x 10-3 3.0 x 10-3 4.0 x 10-3 5.0 x 10-3

Absorbance 0.0 0.36 0.80 1.16 1.62 1.96

Make a graph of absorbance vs. concentration of Fe (mg/mL). The student dissolved 300.0 mg of his sample in 500.0 mLs. He/she took 3 different volumes of the sample solution, added the appropriate reagents to a 100.0 mL volumetric flask (as was done with all the standards) and got the following data: mLs of unknown 10.00 8.00 6.00

Absorbance 1.14 0.90 0.66

Calculate the average % Fe in the sample plus the standard deviation and ppt error for the data. Answers will depend somewhat on method of data treatment Data obtained directly from the graph % Fe = 4.83 ± 0.08 with 17 ppt Data obtained from y = mx + b using graph. m = slope = 396, b = 0 % Fe = 4.72 ± 0.09 with 19 ppt Data obtained from y = mx + b using computer Excel program. m = slope = 398, b = -0.01 % Fe = 4.76 ± 0.07 with 15 ppt

Revised Fall, 2010 P. W. Crawford CH271 Experiment 6B Determination of Soluble Iron in a Commercial Iron Supplement Tablet/Pill/Capsule The total amount of iron in an adult is approximately 4 to 5 grams. About 70 to 75% of this relatively small amount of iron has an active and vital physiological role, and the remaining 25 to 30% is present in various storage forms that can be readily mobilized if needed. The physiologically active iron is mainly present in the form of oxygen carrying pigments such as hemoglobin (about 65%) and myoglobin (3 to 5%), as well as in the form of a number of enzymes involved in electron transfer and reactions (less than 1%). Normally the iron needs of the body are met by a dietary intake of approximately 5 to 20 mg/day. The term tired blood” is a mild case of anemia caused by a shortage of Fe2+. Severe anemia is a condition in which the red corpuscles of the blood are reduced in number or deficient in hemoglobin, causing pallor, shortness of breath, and palpitation of the heart. Anemia, tired blood, is most often treated by increasing ones iron intake with commercially available iron supplement tablets. This experiment will use spectrophotometric measurements to verify the labeled iron content of these supplements. You may bring your own or use an iron supplement tablet available in the laboratory. Procedure A.

Sample Preparation 1. Record the labeled iron content per tablet to be analyzed. 2. Place one tablet in a clean 250 mL Erlenmeyer flask. Add 50 ml of distilled water and one mL of 3 M H2SO4 and gently heat on a hot plate or with a Bunsen burner. Most of the tablet should dissolve leaving some insoluble binder as residue. 3. Transfer the solution to a 100.0 ml volumetric flask, dilute to the mark, mix, and allow any remaining solid residue to settle to the bottom. 4. Calculate the expected iron concentration in mg/ml of the tablet stock solution. Calculate the volume of the tablet stock solution which when diluted to exactly 100.0 mls should have an iron concentration of about 0.020 mg/mL. Round this volume off to the nearest mL. Using this volume of tablet stock solution, prepare a diluted stock solution in a 100.0 mL volumetric flask, i.e. dilute the calculated mL of Fe tablet stock solution to 100.0 mL with distilled water. This diluted tablet solution is now ready to be treated and analyzed spectrophotometrically for Fe as done previously in the first part of this experiment, Experiment 6A.

B.

Color Development and Measurements 1. Pipet exactly 10.00 mL of the diluted stock tablet solution prepared in step A4 above into a clean 100.0 ml volumetric flask. Add 5 mL of pH = 4 buffer, 5 mls of hyroxylamine-HCl, and 5 mL of phenanthroline. Dilute with distilled H2O to exactly 100.0 ml and mix thoroughly. This is the Fe tablet test solution. 2. Measure the absorbance of the 0.0020 mg/ml Fe(Ph)32+ standard at max as was done in Experiment #6A. Now measure the absorbance of the Fe tablet test solution. Using the equation Aunk/Ast = cunk/cst, calculate the concentration of Fe in the Fe tablet test solution. 3. Measure the absorbance of the 0.0020 mg/ml Fe(Ph)32+ standard at max again. Once you have measured the absorbance, change to the concentration mode of the Spectronic 20 and enter the concentration of the standard as 2.00. (Follow the directions in the instructions to the instrument on how to do this.) Now measure the concentration of the Fe tablet test solution directly and record the concentration in mg/ml x 10-3. 4. Measure and record the absorbance of three standard Fe(Ph)32+ solutions from experiment 6A and the Fe tablet test solution at max using a distilled H2O blank. Prepare a Beer's Law Plot as done in experiment 6A. Use your Beer's law plot to determine the concentration (mg/mL) of Fe in the Fe tablet test solution.

Data Treatment: Calculate the total mg of Fe in the commercial tablet preparation using the concentration values obtained in steps B2, B3, and B4 above. Do these values agree with each other? Calculate the absolute and % relative error in your experimental result compared to the labeled content in each case. Which of these three values do you have the most confidence in? Explain your answer briefly!

CH271 Pre-Lab Problems Determination of Soluble Iron in a Commercial Iron Supplement Tablet/Pill/Capsule 1. An iron supplement tablet labeled to contain 70 mg of iron was dissolved in exactly 100.0 mL of H2O. What volume of this solution is needed to give a 0.020 mg/ml Fe solution when diluted to 100 mL. (2.86 mL) 2. Exactly 3.00 mL of the tablet solution in #1 was diluted to 100.0 mL. 10.00 mL of this solution was treated with the appropriate reagents to form Fe(Phen)3 and diluted to 100.0 mL. The concentration of Fe of this solution was found to be 0.0022 mg/ml. Calculate the total mg of Fe per tablet. (73.3 mg) 3. Calculate the % relative error in the experimental value for the mg of Fe/tablet compared to the labeled value. % relative error = (Experimental value - true value) x 100/true value = 4.7% high