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