CHM 161 Spectrophotometry: Analysis of Iron(II) in an Aqueous Solution

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1 CHM 161 Spectrophotometry: Analysis of Iron(II) in an Aqueous Solution Introduction Many compounds exhibit colors in aqueous solution due to the absorption of certain wavelengths of light. The intensity of the color of a solution is proportional to the concentration of the absorbing species. In more concentrated solutions more light is absorbed; however, the wavelength of the absorbed light does not change. A comparison of the intensity of the color of solutions of known concentration with the intensity of an unknown permits identification of the concentration of the unknown solution. In this experiment you will determine the mass of iron in a commercial iron supplement pill and will compare your results with the manufacturer's stated value. You will analyze for iron by allowing iron(ii) to react with an organic compound (1,10-phenanthroline) to form an orange-red complex ion. To ensure that all of the iron in the solution is in the +2 state, the solution is treated with hydroxylamine hydrochloride (NH 2OH. HCl(aq) > or NH 3OH +, Cl ) to reduce any iron(iii) present to iron(ii). The reaction depends on whether the solution is acidic or basic. acid 2 Fe 3+ (aq) + 2 NH 3OH + (aq) 2 Fe 2+ (aq) + N 2(g) + 4 H + (aq) + 2 H 2O (l) 2 Fe 3+ (aq) + 2 NH 2OH (aq) + 2 OH (aq) 2 Fe 2+ (aq) + N 2(g) + 4 H 2O (l) base Sodium acetate is added to the solution to adjust the acidity to a level at which the iron(ii)- phenanthroline complex is especially stable. You will make a definite volume of solution having a known mass of the iron supplement pill, but an unknown concentration of the iron(ii) ion. To determine the concentration of Fe(II), you will first measure the absorbance of a fixed wavelength of light by standard solutions containing known concentrations of the iron(ii) ion. You will use these data to prepare a calibration graph showing absorbance as a function of concentration. You will then measure the absorbance at the same wavelength of light by the "unknown" solution. Using the calibration curve, you will be able to determine the concentration of iron(ii) in the unknown. Knowing the volume of the solution in which the pill was dissolved and the Fe(II) ion concentration, you will be able to calculate the mass of iron in the original pill. Principles (Spectronic 20) The amount of light that a solution absorbs depends on its properties, on the concentrations of absorbing species, and on the inside diameter of the tube holding the sample. The Beer-Lambert Law gives the relation among these variables: A = abc where A is absorbance, b is the solution path length, c is the concentration of the absorbing substance, and a is the molar absorptivity of the solute. The molar absorptivity is a constant for a given solute at a particular wavelength. With proper selection of cuvettes, the path length b can also be held constant. As seen in Figure 1, the concentration is directly proportional to the absorbance. The absorption of light by a solution is measured in a spectrophotometer as shown in Figure 2. schematic representation of the instrument is given in Figure 3.

2 2- Figure 1. Beer's Law Plot for Aqueous CrO A E E E E E-05 Conc. (M) Figure 1. Relationship Between Absorbance and Concentration Figure 2. Spectronic 20

3 Figure 1. Schematic Diagram of a Spectrophotometer Radiation (light) passes from a source through a slit to a wavelength selector which allows radiation of a specific wavelength,, and intensity, I o, to pass through the sample. The sample may absorb some of the light. The light which passes through the sample has the same wavelength,, with an intensity, I. If any light has been absorbed by the sample, I is less than I o. The light then passes to a detector. Its intensity is converted to an electronic signal which is displayed on a meter on the front of the spectrophotometer. The Beer-Lambert law can be written to include the intensities of light before and after passing through the sample: A = log(i o/i) = log(i/i o) = abc The ratio (I/I o) is the transmittance and (I/I o) x 100 is the percent transmittance.. Transmittance of light in the spectrophotometer depends on factors other than the concentration of the absorbing species. When light passes through the walls of the container, interaction between the light and the walls is inevitable. A loss of power will occur. The incoming light beam may also suffer a loss of power due to scattering of light by solvent molecules or by the absorption of light by solution particles other than the solute being analyzed. A simple technique is employed to take these factors into account. A "blank reference" containing all the components of the solution except the species being analyzed is placed in a cuvette that matches the one holding the compound being studied. The spectrophotometer is set so that 100% of the light of the chosen wavelength is transmitted by the "blank reference." Thus, any absorption of light by the test solution will be due to the presence of the substance being analyzed. The spectrophotometer scale is given in both percent transmittance and absorbance. In practice, it is easier to read the percent transmittance than the absorbance scale because the former is linear whereas the latter is logarithmic. The percent transmittance is easily converted to absorbance. A = log(i o/i) = log(i/i o) = log(t) = log(%t/100) To determine the unknown concentration of a certain chemical substance, a calibration curve is first constructed. The absorbance of the substance is measured at several knownn concentrations. The absorbance is plotted as a function of concentration and the best straight line is drawn through the experimental points. The absorbance of a solution of unknown concentration of the substance is then measured, and its concentration is read from the calibration curve.

4 The color of a solution may be due to an anion such as the yellow chromate ion, CrO 4 2 (aq), the purple permanganate ion, MnO 4 (aq), etc.; or a cation such as the blue copper(ii) ion, the green nickel(ii) ion, etc. Some solutions such as aqueous iron(ii) nitrate do not have sufficient color to give good colorimetric analyses. In such a case, a small amount of a reagent that forms a sufficiently colored complex ion can be added to the solution containing the compound to be analyzed. In this experiment, a small amount of 1,10-phenanthroline, C 12H 8N 2, is mixed with a solution containing iron(ii) to form an orange-red complex ion. Fe 2+ (aq) + 3 C 12H 8N 2(aq) [Fe(C 12H 8N 2) 3] 2+ (aq) Analysis of Iron(II) in an Aqueous Solution Procedure: Students should work in groups of 4 on this experiment. You will use volumetric glassware to make solutions of accurately determined concentrations and then use a Spectronic 20 spectrophotometer to determine the absorbance of light by their samples. Each student in the group must participate in all procedures. Part 1.a. Determination of Iron in an Iron Supplement Pill - Sample Preparation* Select one of the "iron supplement" brands. Record in your notebook the brand name and the stated amount of iron in each tablet. One tablet will be shared among the four students in the group. Record the mass of the pill in your notebooks. * Grind the pill to a fine powder with a mortar and pestle and divide it into 3 or 4 equal portions. Transfer your sample of the powder to a 125 ml flask and add about 10 ml of 6 M HCl to the powder in your flask and stir to help the powder dissolve. Stopper the flask and let this suspension sit in the acid solution until the next lab period. The acid will "digest" the tablet. Some of the tablet's binder will not dissolve and the solution will remain cloudy. *This step will be done for you. You should record the total pill weight and the ¼ pill weight into your lab notebooks. Analysis of Iron(II) in an Aqueous Solution Procedure Part 1.b. Sample Preparation By means of gravity filtration, filter the solution and collect the solution in a clean 200 ml or 250 ml volumetric flask. (See the experiment on the synthesis of calcium carbonate for details on how to perform a gravity filtration.) To ensure that all of the iron is transferred into the volumetric flask, wash the solid portion remaining in the filter paper two times with 10 ml of distilled water from your wash bottle. Continue to collect the solution in the volumetric flask. At this point the solution in the volumetric flask should be clear. If it is cloudy, you should filter again using new filter paper. If the solution is clear, add distilled water until the flask is about ½ full. Swirl it to mix the contents. Add more distilled water until the bottom part of the flask is almost full and swirl again. Using your wash bottle, fill the neck of the flask with distilled water almost up to the reference line. Use an eye dropper to fill the neck until the bottom of the meniscus rests exactly on the line. Stopper the flask and shake well. Let the flask sit for at least five minutes continuing.

5 The concentration of iron(ii) in this flask is too high to be used without additional dilution. (The absorbance is too high for the Spec 20 to analyze.) Use wash and good beakers (see pipetting instructions) to rinse a 2 ml Mohr pipet and transfer an exactly 2.00 ml of this solution into a 100 ml volumetric flask. Label the flask, Diluted pill sample. Add reagents to the solution as for the blank in Part 2. below: 10.0 ml of 10% hydroxylamine hydrochloride, 10.0 ml of 0.2% 1,10-phenanthroline, and 8.0 ml of 10% sodium acetate solution. Swirl and dilute the solution using the same procedure as above; cap and shake well. Part 2. Preparation of the Five Known Iron(II) Solutions Each team will need five 100 ml volumetric flasks with stoppers. The flasks, initially filled with distilled water, should be emptied and must be chemically clean. (How can you tell?) Five solutions will be made, each with a different, but precisely known volume of a standard solution of iron(ii) ammonium sulfate, Fe(NH 4) 2(SO 4) 2. 6 H 2O. The standard solution contains grams/l of Fe(NH 4) 2(SO 4) 2. 6 H 2O. This concentration corresponds to 1.00 x 10 1 mg/l of Fe(II). Into each flask, put: 10.0 ml of 10% hydroxylamine hydrochloride solution, 10.0 ml of 0.20% 1,10-phenanthroline solution, and 8.0 ml of 10% sodium acetate solution. (A 10 ml graduated cylinder is adequate for this purpose since all the reagents are in excess.) Label the flasks "0", "5", "10", "25", and "50" corresponding to the amount of iron(ii) ammonium sulfate to be added to the flask. "0" Do not add any of the iron(ii) ammonium sulfate solution to the flask. This solution will be your "blank reference." "5" Pipet 5.00 ml of the iron(ii) ammonium sulfate solution to the flask. "10" Pipet ml of the iron(ii) ammonium sulfate solution to the flask. "25" Pipet ml of the iron(ii) ammonium sulfate solution to the flask. "50" Pipet ml of the iron(ii) ammonium sulfate solution to the flask. After adding the indicated reagents to each flask, add distilled water until the flask is about ½ full. Swirl it to mix the contents. Add more distilled water until the bottom part of the flask is almost full and swirl again. Using your wash bottle, fill the neck of the flask with distilled water almost up to the reference line. Use an eye dropper to fill the neck until the bottom of the meniscus rests exactly on the line. Stopper the flask and shake well. Let the flasks sit for at least five minutes before using the contents. The wavelength at which the colored solution absorbs most strongly, max, is determined by measuring the absorbance of a solution on the Spec 20 over a large range of wavelengths. Estimate the value of max from Figure 4. below and set the wavelength control on the Spec. 20 to this value. Record the value of max you used in your notebook. Figure 4. Graphical Determination of max

6 Part 3. %T and Absorbance Measurements on Known Solutions Make a table in your notebook with rows labeled 5, 10, 15, and 20 and with columns labeled %T, A, T calc, A calc, and Conc., (mg/l). Analysis of samples is done in small tubes, called cuvettes, which are placed in the sample compartment. The cuvettes look much like ordinary test tubes. However, they are specially fabricated and selected to have a constant diameter. These tubes are expensive and should handled carefully and not be mixed with other test tubes in the laboratory. The cuvette should be chemically clean. (How can you tell?) If it is, rinse it threee times with small amounts of the solution whose absorbance is to be measured. Next, fill the cuvette approximately 3/4 full with the solution. Each cuvette has a white line or reference symbol painted on it. Wipe the outside of the cuvette with a tissue before inserting it into the sample compartment in the instrument. When the cuvette is placed in the Spec 20, the line or symbol should face directly toward the reference line on the lip of the sample compartment. The Spec 20 will have been turned on before the beginning of the laboratory period. Before making the measurements, perform the following calibration steps. 1. Turn the Power Switch/Zero Control counterclockwise to the first resistance. If you were to continue to turn the knob, the instrument would be turned off. 2. Turn the Wavelength Control knob to the desired wavelength. 3. With no sample in the Sample Compartment and with the cover closed, turn the Power Switch/Zero Control (left knob) to bring the meter needle to zero on the "percent transmittance" scale of the meter. Look at the needle with one eye closed and line up the needlee with its mirror image and with the 0.0 % T line.

7 4. Wipe the outside of the cuvette filled with the "0" solution with a dry Kimwipe tissue, insert it in the Sample Compartment, and close the cover. Adjust the Transmittance/Absorbance Control (right knob) until the meter reads 100 percent transmittance. Do not touch any of the knobs until after you are finished with all the measurements. 5. Remove the "0" solution and put the cuvette into a beaker. 6. Wipe the outside of the cuvette filled with the "5" solution, insert it in the Sample Compartment, close the cover and read the percent transmittance, estimating to 1 decimal place, from the meter. Record the value in your notebook. Also record an estimated value of A as a check for your later calculations. 7. Repeat Step 6 for each standard solution. 8. Fill in the next two columns in your table: T = %T/100; A = log(t). Calculations for Beer s Law Plot Each of your five volumetric flasks contains a solution whose Fe 2+ concentration, C final, can be easily calculated: C initial x V initial = C final x V final (Follow the rules of significant figures!) C initial = 1.00 x 10 1 mg/l C final =?? V initial = Volume iron(ii) ammonium sulfate added V final = ml C final = (1.00 x 10 1 mg/l) x (Volume iron(ii) ammonium sulfate added) /100.0 ml Put these calculated concentrations in the last column in your table. Draw a rough graph of A vs. the concentration of iron(ii) (mg/l). The four points should fall on a straight line. If they do not, consult your instructor. This line is your calibration curve and will be made more accurately using Excel and used in Part 4 to determine the concentration of the iron in your unknown sample. This result can then be used to calculate the mass of iron in the iron supplement pill. Part 4. %T and Absorbance Measurements on Unknown Iron Solution Measure the % T of your unknown solution as you did for the standards. Zero the instrument with the blank as before and record the percent transmittance to 1 decimal place. Calculate an accurate absorbance from T. Use the slope and y intercept of the calibration curve and the calculated value of the absorbance of your solution to determine the iron concentration in the 100 ml volumetric flask: C unk = (A unk y int.)/slope This is the concentration of the diluted solution in the 100 ml volumetric flask. Calculations Sample Calculations A student performed the experiment following the procedures outlined above. The original pill had a mass of g. A sample mass of g was transferred to the 250 ml volumetric flask.

8 A volume of 2.00 ml was withdrawn from the flask, added to a 100 ml volumetric flask and diluted to a volume of 100 ml. A sample from this flask was analyzed in a Spec 20 and a concentration of 2.5 mg/l was determined from the calibration curve. How many mg iron were contained in the original iron supplement pill? 1. Calculation of Concentration of the Original Unknown Solution in the ml Flask. C f x V f = C i x V i (1.28 mg/l)* x (100 ml) = (C i) x (2.00 ml) C i = 64.0 mg/l *From Beer s Law plot: C i = C unk = (A unk y int.)/slope 2. Calculation of Mass of Iron in the Original 250 ml Flask (64.0 mg/l) x (0.250 L) = 16.0 mg Fe in flask 3. Mass of Iron in Original Pill 16.0 mg Fe in sample x (0.346 g of pill)/0.085 g of sample) = 65 mg Fe in pill Compare the amount of iron you determined to exist in the tablet with the amount of iron listed on the manufacturer's label. After completing the experiment, empty all cuvettes, rinse each one three times with distilled water, and return them. Empty all volumetric flasks and rinse each twice with distilled water. Fill the flasks with distilled water up into the neck and stopper them. Return the flasks to where they were found.

9 CHM 161: Fe Spectrophotometric POST-LABORATORY ASSIGNMENT 1. Construct a data table (as shown below) for the standard solutions. This table must include percent transmittance, absorbance, and the concentration for each of the four standards. Standard Solution %T Absorbance [Fe] (mg/l) Show a sample conversion of percent transmittance to absorbance using the 10 standard solution. 3. Show a sample calculation for [Fe] mg/l using the 10 standard solution. 4. Construct a graph of A (absorbance) as a function of concentration of Fe(II) in mg/l for the four standard solutions. On this graph, show clearly how you determined the concentration of Fe(II) in the solution made from the iron supplement pill. (See separate handout on Beer s Law Plots.)

10 5. Report the mg iron in the pill and calculate your percent error using the manufacturer's claim (65 mg Fe) as the accepted value.

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