Spectrophotometric Determination of Iron

Transcription

1 Spectrophotometric Determination of Iron Introduction: For chemical species that appear to have color, it is a logical assumption that the intensity of the color is proportional to the concentration of the species in solution. We see color as a complement of the visible wavelength being absorbed by the sample. Things that appear red absorb blue visible light and reflect other visible colors to our eyes. Conversely, things that appear blue are absorbing red light. Absorption of light or more precisely electromagnetic radiation is related to available energy levels in the molecule or ion. A molecule in its "ground state" or lowest energy level can absorb energy to jump to an "excited state" or higher energy state. The amount of energy and therefore the wavelength of radiation involved in this transition is a function of the electronic structure of the molecule or ion. The eye can only see a limited range of electromagnetic radiation, from approximately 400 to 700 nm. However, molecules, atom, and ions are capable of absorbing many different energies of radiation ranging from ultraviolet (UV) to microwaves depending on the specific energy levels being excited. For some types of energy changes, the wavelengths of light are very specific for certain types of chemical structure resulting in a method of qualitatively identifying chemical species. Other types of energy absorption may be less qualitative since it may relate only to bond types. In both cases however, our initial premise that intensity of absorption is related to concentration can be used for quantitative analysis. A spectrometer is an instrument that measures the amount of light through transmitted through a substance. The drawing below illustrates an Ocean Optics spectrometer, which we will use in this experiment. hυ (1) The light transmitted through a sample cuvette enters the spectrometer. (2) A small aperture, or slit, controls the amount of light that enters the spectrometer. (3) A lens, (4) a mirror, (5) and another mirror control the path that the light takes on its way to (6) the diffraction grating, which separates (or disperses) light into its wavelengths (like a prism). (7) Each wavelength is detected separately with a CCD detector. cuvette hυ Page 1 of 5

2 detector I Light source I 0 When light is absorbed by a sample, the radiant power or intensity of the light transmitted decreases. Radiant power, I, refers to the energy per second per unit area of the beam. In the figure, light passes through a diffraction grating and aperture that selects one wavelength. Light of this wavelength, with radiant power I 0, passes through a sample of pathlength b. The radiant power of the beam emerging from the other side of the sample is I. Mathematically, the amount of light that is absorbed (A) is given by Note that if no light is absorbed, A = 0. The amount of light absorbed by the sample should be proportional to the molar absorptivity (ε, units are L/mol cm), the concentration (C, in mol/l), and the cuvette pathlength (b, in cm). This relationship is quantified in the Beer-Lambert (or Beer's) Law which is A = εbc A has no units, although our instrument uses o.d. (optical density) in place of units. Note that this equation can take the form of a straight line equation, y = mx + b where the intercept, b, is zero. If we measure a series of solutions of known C at a given wavelength in a cuvette with a pathlength b we can then determine the slope of the line, m, which is ε x b. This procedure generates a "calibration curve" which allows the determination of an unknown concentration, C unk, from the measurement of the absorbance of the unknown, A unk. Page 2 of 5

3 Many of the transition metal ions such as copper, nickel, cobalt, and chromium exhibit color in solution. However, this color can be made more intense by reacting the metal ion with a molecule that increases the absorbance of the metal ion. Iron(II), Fe 2+, exhibits little color in solution. When Fe 2+ complexes with the ligand o-phenanthroline (or 1,10-phenathroline), a stable, intensely colored red complex is formed that can be used to determine iron. The intensity of the color varies over the ph range of 2 to 9. In this procedure an ammonium acetate buffer will adjust the ph to between 6 and 9. Fe 2+ will form a 1:3 complex with the 1,10-phenanthroline ligand. That is, 3 ligands will bind to each Fe 2+ ion. 1,10-phenanthroline is bidentate. The iron must be in the +2 oxidation state, requiring a pre-reduction step before formation of the colored complex. Hydroxylamine is used as a reducing agent. 2 Fe NH 2 OH + 2 OH - 2 Fe 2+ + N H 2 O Before Lab you should: Read sections 18-1 through 18-4 in your textbook. Write the reaction of the reduction of Fe 3+ by hydroxylamine in your lab notebook. Draw the Fe 2+ :1,10-phenanthroline complex in your lab notebook. Record the molar mass of iron (II) ammonium sulfate hexahydrate in your notebook. Reagents: Hydroxylamine solution 9 M H 2 SO 4 Sodium acetate solution 1,10-phenanthroline solution Iron (II) (ferrous) ammonium sulfate hexahydrate EXPERIMENTAL PROCEDURE Preparation of Standards You will be divided into groups to prepare the following solutions: 1. Prepare a stock Fe solution by accurately weighing to the nearest 0.1 mg approximately 0.07 g of pure iron (II) ammonium sulfate hexahydrate and quantitatively transferring to a 1 L volumetric flask. 2. Add 200 ml water and shake to dissolve any remaining solid. 3. Add 5 ml of 9 M sulfuric acid. 4. Dilute to the mark with distilled water and homogenize thoroughly. 5. Calculate the concentration of the solution in mg Fe/L. Convert this concentration into mol/l. Prepare a series of standards by pipetting into each of five ml volumetric flasks, 1.00, 5.00, 10.00, 25.00, and ml aliquots of the stock Fe 2+ solution (a buret can be used for this addition). Calculate the concentration of each dilution in mol/l and label each solution with its concentration. 6. Into a sixth ml volumetric flask pipet 50 ml of distilled water to serve as a blank. Page 3 of 5

4 7. To all of the solutions add, in this order, 1 ml of hydroxylamine hydrochloride solution, 10 ml of 1,10-phenanthroline, and 8 ml of sodium acetate buffer. 8. Dilute to the mark, mix thoroughly, and allow to stand for 10 minutes. Preparation of Unknown In order to prepare your unknown in the desired concentration range for the spectrophotometric measurement it will be necessary to do a serial dilution. Each student will individually prepare their own unknown. Be sure to record your unknown number in your lab notebook. 1. Accurately weight g of your unknown to the nearest 0.1 mg. 2. Quantitatively transfer the solid to a 1 L volumetric flask. 3. Add 200 ml of water, 50 ml of 9 M H 2 SO 4 and dissolve. 4. Dilute to the mark and mix thoroughly 5. Pipet a ml aliquot into a ml volumetric flask. 6. Dilute to the mark and mix thoroughly Finally prepare the actual sample for analysis. 7. Pipet 20.0 ml of this second solution into a ml volumetric flask. 8. Treat this as you did your standards by adding, in this order, 1 ml of the hydroxylamine solution, 10 ml of the 1,10- phenanthroline solution, and 8 ml of sodium acetate. 9. Dilute to the mark, mix thoroughly and allow to stand for 10 minutes. Preparation of cuvettes Using a ruler with mm gradations, measure the pathlength (not including the thickness of the plastic walls) of your cuvettes in cm to the correct number of decimal places. (All cuvettes have the same pathlength). Taking care to remember which solution is in which cuvette (don t write or put tape on them), fill one cuvette with each solution (blank, standards, and unknowns). It is important that the cuvettes are not marked, scratched, or smudged with finger prints use Kimwipes to keep them clean during experiment. Spectrophotometric Analysis of Solutions Spectrasuite software for Ocean Optics Absorbance Measurements 1. Spectrasuite software is open and recognizing the USB650 spectrometer, showing a signal trace on the graph. 2. Near the top of the screen is a check box labeled Strobe/Lamp Enable. Click the box so that it is checked. The trace should now look very different, a signal so large that its top is cut off. 3. In the top right corner, change the integration time from 100 ms to 25 ms. The signal should not be cut off now. 4. Put your blank into the cuvette holder such that the arrows line up. Make sure it is pushed down into the holder. 5. Push the button that looks like a yellow light bulb, this sets your blank measurement. 6. Now take out the blank and put in the black cuvette. 7. Push the button that looks like a gray light bulb, this sets your dark measurement. Page 4 of 5

5 8. Take out the black cuvette and click on the blue A button. This takes you to the graph where you can do absorbance measurements. 9. Put your most concentrated sample into the cuvette holder. Observe the graph. You should see an absorbance peak. 10. Click anywhere in the graph to get a cursor to appear. Now click the cursor to the part of the peak where you see it is the highest most absorbance. This is your λ max. Look just below the x-axis, on the left, and you will see the wavelength (nm) for the location of your cursor. 11. In your lab notebook, sketch this graph, including the labels on the axes, the range of wavelengths that the absorbance peak covers, and the λ max. You will keep the cursor at λ max for the rest of the experiment. 12. Just to the left of the wavelength value, below the x-axis, there is a red number that is changing slightly every second. That is the absorbance value. Go to the top of the screen and set the scans to average to 5, and the boxcar width to 3. This should slow down the rate that the absorbance is changing. (If someone used the spectrometer before you today, this may have already been done.) 13. Starting with your least concentrated solution, put the cuvette into the holder and wait 5 seconds. Write down the absorbance for each solution. If the number is still fluctuating, watch it for a few seconds and take the number that you think is most representative. 14. Obtain a table of concentrations (mg/l) and absorbances. Plot these in Excel as concentration vs. absorbance. Obtain the trendline and R2 value and have them appear on the graph. Your R 2 value should be no less than to obtain good results. If this is not the case, you and your partner can consider re-doing the graph. If the new R 2 is not better, it is because your standards are not completely accurate talk to your instructor. Calculations Using the line equation (in the format y = mx + b), and the Beer s Law Equation (A = εbc), calculate ε for your complex ion (in L/mol cm). Calculate the percentage of iron in your unknown at this point: Using the calculated concentration of the complex ion in your solution, calculate the amount of iron (in mg) in the unknown stock solution. Report the result in terms of the % Fe in your unknown. The range of unknown values should be 5 % to 15 %. Print your absorbance spectrum and Beer s Law Plot and submit them with your ε and % Fe. Page 5 of 5