2011_UV-Vis Iron(III)-salicylate complex.pdf

Instrumental Analysis Laboratory College of Charleston UV-VIS SPECTROMETRY

UV-VIS Spectrometry Multivariate Linear Regression Page 1 of 11

SPECTROPHOTOMETRIC MEASUREMENTS Adaptedfrom: SalicylateDetectionbyComplexationwithIron(III)andOpticalAbsorbanceSpectroscopyAnUndergraduateQuantitativeAnalysisExperiment,J.Chem.Ed.,2008,85,16581659. J. T. Mitchell-Koch, Department of Chemistry, Emporia State University, Emporia, KS, K. R. Reid and M. E. Meyerhoff, Department of Chemistry, University of Michigan, Ann Arbor, MI 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. Since our vision if not quantitatively calibrated, an electronic instrument called a spectrophotometer is used to precisely measure light intensities at given energy (wavelength) settings. A spectrophotometer is an instrument that measures the amount of transmission of light through a substance. The drawing below illustrates a simple spectrophotometer system consisting of a light (energy) source, a monochromator to select a given energy range, a sample, and a light intensity detector.



When light is absorbed by a sample, the radiant power or intensity of the light beam decreases. Radiant power, I, refers to the energy per second per unit area of the beam. In the figure, light passes through a monochromator that selects one wavelength. Light of this wavelength, with radiant power I0, 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

0IA lnI

Note that if no light is absorbed, A = 0 and if all the light is absorbed ( I = 0) then A = . The amount of light absorbed by the sample should be proportional to the probability that the molecule or ion will absorb the electromagnetic radiation (a), the number of absorbing molecules or ions per unit volume that the light beam passes through (C), and the length of the light path (b). This relationship is quantified in the Beer-Lambert (or Beer's) Law which is

A = a b C Note that this equation is in the form of Y = m X + b where the intercept, b, is zero when X or the concentration, C, is zero. If we measure a series of solutions of known C at a given wavelength in a cuvet or sample cell with a constant pathlength, b, then we can determine the proportionality constant, m, which is a b. This procedure generates a "calibration curve" which allows the determination of an unknown concentration, Cunk, from the measurement of the absorbance of the unknown, Aunk. Determination of the slope, m, and intercept, b, of the calibration curve gives

unkunk

ACm b

However, many systems involve more than one colored component. If these components act independently, then Beers Law is still applicable but more than one wavelength must be used for the analysis. The figure below illustrates a two component spectrum.



The maximum absorbance of component 1 occurs at wavelength 1 while the maximum absorbance of component 2 occurs at wavelength 2. The dotted line represents a solution that is a mixture of components 1 and 2. In this experiment, we will use a fiber optic diode array spectrophotometer. A schematic diagram of this instrument is shown below.

The advantage of this instrument is that all wavelengths are recorded at once. Therefore we can signal-average to reduce noise and apply other digital spectral smoothing techniques. The spectrometer uses a high-pressure deuterium lamp to produce ultraviolet radiation but the instrument is less sensitive in the UV region as compared to the visible region where the spectral source is a incandescent tungsten lamp. In this experiment, we will only use spectral data between 450 and 650 nm although every spectrum records data from 187 to 900 nm. We will analyze the quantitative data using linear regression which in this case is applying Beers Law the component in our sample.

VisibleSpectrophotometry:DeterminationofSalicylateviaReactionwithFe(III)

Background Spectroscopic analysis is a critical tool in the identification and quantitation of different molecules. This experiment introduces you to the use of electronic absorption spectroscopy in the visible region of the spectrum for the determination of salicylate. There are several uses for salicylate and it is therefore included in many everyday products. Salicylic acid is the major metabolite of aspirin and is commonly found in medications that treat acne, warts and other similar ailments. When acetylsalicylic acid (aspirin) is taken for a headache or inflammation, it is rapidly hydrolyzed in the stomach. The products of this reaction are salicylic acid and acetic acid. The former is readily absorbed into the blood stream and is then able to act as an analgesic agent. In acne treatment, the salicylic acid decreases the shedding of skin cells from hair follicles. These cells are typically responsible for clogging pores and causing pimples. Salicylic acid also has a keratolytic (peeling) effect, which causes dead cells to be shed more easily. This facilitates in the removal of a thin layer of skin and promotes the unclogging of pores. More concentrated



solutions of salicylic acid are used in wart treatment to help soften the wart and to stimulate an immune response toward the human papillomavirus, responsible for causing wart formation. Due to the many medical applications of salicylic acid, the development of analytical techniques for its quantification is important. Indeed, there are a number of methods that have been employed, including, gas-liquid chromatography (GLC), ultraviolet spectroscopy, and fluorescence spectroscopy. The most widely used methods in clinical laboratories, however, use colorimetric or visible spectrophotometry. A version of this method will be applied throughout the experimental procedure to first quantitate salicylate in a commercial product (face wash), and also in an unknown solution that you will be given. The second part of the procedure uses spectrophotometry to investigate the chemical nature of the reaction that yields the colored product you analyze. Measurement Principles Beer's Law states that the absorbance of a compound is directly proportional to its concentration (A=abc). This linear relationship allows us to first construct a calibration curve by collecting the absorbance values for samples of known concentration at a given wavelength, preferably the max, the wavelength where maximum absorption occurs. The resulting equation for the linear regression then lets us determine the concentration of an unknown sample by determining its absorbance at the same wavelength. Salicylate and salicylic acid do not absorb visible light, creating an experimental challenge. Upon reaction with iron (III) ions, however, a highly colored species results:

Salicylic acid (sal) iron(iii)-salicylate complex highly colored The complex can be easily detected with a simple spectrophotometer and thus, you will be able to quantify salicylate in unknown samples. Under the acidic experimental conditions all salicylate will be protonated as shown in the chemical equation above. The chemical equation shown above contains the coefficients and subscripts x and y. In the second portion of this experiment, you will use the method of continuous variation (also called Job's method) to determine these quantities for the predominant complex. For this procedure, several solutions containing different quantities of salicylate and Fe3+ will be prepared. While the amount of each reactant is varied, the total moles of both reagents will remain constant. The solution that yields the greatest absorbance at max indicates the predominant stoichiometry of the iron-salicylate complex.

OH

O

OH+X Y Fe 3+ (Fe )3+ y(Sal)x



Safety Hazards General laboratory safety rules should be followed. Nitric acid is corrosive, and spills should be cleaned up immediately.

INSTRUCTIONS UV-VIS SPECTROPHOTOMETRY EXPERIMENT

Be sure to clean up the area when finished Run the SpectraSuite program from the Desktop. Make sure the detector is hooked up to the USB port and the lamp module is on. The instructor will show how to run the dark current (electronic diode noise) and the use water or the 10 mM Fe(NO3)3 solutions in a cuvet as the reference spectrum. You may have to adjust the integration time, number of runs, and boxcar smoothing to obtain the optimum spectra. Check with the instructor on how this can be done. Save each spectrum as a tab-separated variable *.txt file and save. You will need a flash drive to copy the files. With the spectrometer on, block the light beam with a plug and save the dark current by pressing the gray bulb.

Then place your blank in the beam and adjust scans to average to 3 and the boxcar integrator to 5. Adjust the integration time until the high point of the spectrum stays at or near 4000 counts all across the spectral region that you are interested in.



The block the source again and record the dark current under the new conditions. Never select anything from the top level commands. To save the spectrum, always select the file disk circled in the picture. The A or absorbance button should be lit and scale set to absorbance. You can select the little magnifying glass highlighted in the last picture to set the viewing wavelength and absorbance units. To save the file, select the file disk, choose Processed Spectrum, Tab Delimited-No Header, and then choose the Browse button.

Create a file with the first spectrum to be saved, open the folder and then type in the file name and select Save.



Select Save again.

The Save command will gray out and the choose Close. You are now ready to insert the cuvet with your next sample.

In this experiment, the concentration of salicylate present in an over the counter acne medication/face wash and in an unknown sample will be determined by spectrophotometry. Salicylate itself absorbs ultra-violet radiation and is therefore difficult to measure directly with simple instrumentation. One method adopted for the measurement of salicylate in clinical situations involves mixing samples containing salicylate with an excess of ferric ions, Fe(III) under acidic conditions. The resulting complex absorbs strongly in the visible region of the spectrum and can be easily determined spectrophotometrically. The first section of the experiment involves using this salicylate-iron complex for the determination of salicylate concentration in an acne medication and an unknown sample. This will be possible by first generating a calibration curve for salicylate from several standard solutions of different concentration. In the second section the nature of the salicylate-iron complex will be investigated by using the method of continuous variation. This procedure involves varying the amount of each reagent added, salicylate and Fe(III), while keeping the total number of moles constant. The mixture yielding the maximum absorbance corresponds to the predominant stoichiometry of the complex formation.



Solutions Needed for the Experiment: Solution Composition Notes

1. 0.1 M (100 mM) sodium salicylate) Weigh out 16.01 of sodium salicylate (MM - 160.11

g/mole) and dilute to 1.0 L with distilled water.

0.010 M (10 mM) sodium salicylate Dilute the 100 mM sodium salicylate 1:10.

10 mM Fe3+ Dilute 5.987 mL of the stock 404 gm/L Fe(NO3)3 stock to

1.0 L with 0.060 M HNO3

60 mM HNO3 Dilute 3.797 mL of concentrated HNO3 (15.8 M) to 1.0 L. PartA:DeterminationofReactionStoichiometry:AnApplicationofthe

MethodofContinuousVariation1. Obtain ~20 mL 10 mM acidic ferric nitrate as well as ~50 mL of dilute nitric acid (60mM) in small beakers. You will also need the 10 mM sodium salicylate solution. Return any unused solution to the bottle.

2. Prepare solutions for spectrophotometric analysis by pipetting the appropriate amount of each solution into a small test tube. Use the amounts from the Table 1 below. Label all your vials.

3. Add 3.00 mL of 60 mM HNO3 to make each test tube to a total volume of 4.00 mL. Mix thoroughly and add to a plastic cuvet.

Table1:SolutionCompositionforMethodofContinuousVariation

SolutionVolume10mMsalicylate(mL)

Volume10mMferricnitrate(mL)

Volume60mMnitricacid

MoleRatio

Fe(NO3)3:

salicylate

MoleFractionFe(NO3)3

1 0.100 0.900 3.000 9.00 0.90 2 0.200 0.800 3.000 4.00 0.80 3 0.250 0.750 3.000 3.00 0.75 4 0.330 0.670 3.000 1.99 0.67 5 0.400 0.600 3.000 1.50 0.60 6 0.500 0.500 3.000 1.00 0.50 7 0.600 0.400 3.000 0.67 0.40 8 0.670 0.330 3.000 0.50 0.34 9 0.750 0.250 3.000 0.33 0.25 10 0.800 0.200 3.000 0.25 0.20 11 0.900 0.100 3.000 0.11 0.10



4. Using the 60 mM HNO3 as your blank, collect the absorbance values each solution. Data Analysis Part A: The data collected in this part of this experiment will allow you to examine the nature of the reaction between Fe(III) and salicylate. While the mole ratios of reagents were varied in each mixture, the total number of moles remained the same. Therefore, the mixture that yields the greatest absorbance represents the predominant reaction stoichiometry. In order to find which stoichiometry is favored by this complex, plot the absorbance value of each solution (at the max) versus the mole fraction of iron. PartB:SpectrophotometricDeterminationofSalicylateinAcneMedication 1. Prepare five standard solutions of sodium salicylate in deionized water. For this task, we use an initial stock solution of 100 mM (0.1M) of sodium salicylate. By dilution in appropriate a 10 mL volumetric flask; prepare standards of 20 mM, 40 mM, 60 mM and 80 mM (10 mL volumetric flasks will give enough of these standards for this experiment). Sodium salicylate has a formula weight of 160.11 g/mole. It is important to record the accurate mass of sodium salicylate that you used to prepare the stock solution, so that you will know the exact molarity of these standards. Dilute them to volume in the 10 mL volumetric flask with distilled water.

Table2:SolutionPreparationforCalibrationCurveSolution

mL100mMsodiumsalicylate/10mL

volumetricConcentrationofsodiumsalicylate

1 2.00 20mM2 4.00 40mM3 6.00 60mM4 8.00 80mM

2. Obtain a sample of acne face wash solution as well as an unknown salicylate sample (solution). Be sure to record the code of your unknown for your lab report. 3. In separate test tubes, pipet exactly 50 L of each standard, including the 100 mM sodium salicylate, the acne face wash, and your unknown. It is important to label these test tubes for identification of each solution. Add 5.00 mL the acidic 10 mM ferric nitrate solution (stock in lab) to each test tube. Be sure to mix these solutions well! 4. Using the acidic 10 mM ferric nitrate solution as your blank (100%T), collect optical absorbance spectra for all of the solutions Data Analysis Part B: 1. The objective for this section of the experiment is to determine the concentration of salicylate in unknown samples. Construction of a calibration curve from the absorbance data collected



from the salicylate standard solutions is the first necessary step towards this goal. In Excel, plot the absorbance value of each standard (at the versus the standard's concentration. The data should show a linear relationship. Generate a linear regression line and equation for this line. 2. Determine the concentration of salicylate for each of your samples of acne medication and unknowns from the linear regression line. Convert your face wash data to units of weight percent assuming the density of the solution to be 1.00 gm/mL. Data Processing in Excel: After the spectra are run the data files can be imported into EXCEL. The data will be in the form of wavelength and absorbance data for each spectrum. Copy this data into EXCEL using the text file input command and tab-delimited data. Column A is always constant contains the wavelength of the spectrum and column B should be the absorbance. Move the cursor to column C and import the next set of data. Make sure that the data is aligned correctly according to wavelength and then copy only the absorbance data column B). After importing the spectral data, your spreadsheet should now contain the wavelengths in column A and columns of absorbances. Be sure to label the data in each column. You can simplify the spreadsheet by deleting all wavelengths from 187 to 300 and 650 to 900 nm. Plot the absorbances versus wavelength for your method of continuous variation and analytical data. This is simplified in Excel by finding the wavelength and absorbances that you want to copy and the using the TRANSPOSE option in the Paste Special command.



Your report should contain the following:

1.ProvideaplotofthesalicylateFe(III)complexspectrafromthemethodofcontinuousvariations.

2.Showthe"JobPlot"ofabsorbance(atmax)versusmolefractionofiron(III)that

youobtainedinpartAoftheexperiment.Fromthisdata,indicatewhatyoubelieveisthestoichiometryofthereaction.

3.ProvideaplotofthesalicylateFe(III)complexspectraforyourcalibrationcurves

(includingyourunknownspectra.)4.Showtheplotofthecalibrationcurveforsalicylate(includingequationoflinewith

R2value).5.Reportthemeanconcentrationofsalicylateinacnefacewash(inunitsofweight

percent)thatyoufound.6.Reporttheaverageconcentrationofsalicylateinyourunknownsample(inunitsof

molarity,ormillimolarity.)Makesureyouindicateyourunknownsamplecode.. REFERENCES: 1. Lab Handout. 2. Ferguson, G. K. J. Chem. Educ. 1998, 75, 467469. 3. Hein, J.; Jeannot, M. J. Chem. Educ. 2001, 78, 224225. 4. Simonson, L. A. J. Chem. Educ. 2001, 78, 1387. 5. Yang, S.-P.; Tsai, R.-Y. J. Chem. Educ. 2006, 83, 906909. 6. Cavanaugh, M. A.; Bambenek, M. A. J. Chem. Educ. 1978, 55, 464. 7. Lane, S. R.; Stewart, J. T. J. Chem. Educ. 1974, 51, 588589. 8. Battezzati, A.; Fiorillo, G.; Spadafranca, A.; Bertoli, S.; Testolin, G. Anal. Biochem. 2006,

354, 274278. 9. Rogic, D. J. Mol. Struc. 1993, 294, 255258. 10. Lange, W. E.; Bell, S. A. J. Pharm. Sci. 1966, 55, 386389. 11. Saltzman, A. J. Biol. Chem. 1948, 174, 399404. 12. Annino, J. S.; Giese, R. W. Clinical Chemistry: Principles and Procedures, 4th ed.; Little,

Brown and Co.: Boston, 1976; pp 355357.

2011_UV-Vis Iron(III)-salicylate complex.pdf

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