Spectrophotometric Spectrophotometric Determination of Iron in Aqueous Solutions as a Complex of 1,10Phenanthroline J.P. Loja Institute of Chemistry, University of the Philippines, Diliman, Quezon City
August 30, 2013 September 6, 2013 I.
Methodology
The materials used in the experiment were measuring pipettes, beaker and volumetric pipettes. The first part of the experiment was the preparation of solutions listed below. 50 mL 25 ppm working standard Fe (II) solution 50 mL 10% (w/v) hydroxylamine hydrochloride solution 250 mL 0.20% (w/v) 1,10-phenanthroline solution 100 mL 1.0 M acetic acid 100 mL 1.0 M sodium acetate 100 mL 0.10 M acetate buffer with pH 4 from 1.0 M acetic acid and sodium acetate The solution preparation for the standardization was prepared by getting six clean volumetric flasks and labelling it from 0 to 5. These were added with 0.00, 1.00, 2.00, 3.00, 4.00 and 5.00 mL of 25 ppm working standard Fe (II) solution starting from 0. The flasks were added with 1.00 mL 10% hydroxylamine hydrochloride solution, 5.00 mL 0.20% 1,10-phenanthroline solution and 1.00 mL of acetate buffer pH 4 successively. The solutions were then swirled and diluted to mark. The solution preparation for the unknown was prepared by getting a 100 mL volumetric flask for unknown Fe (II) sample solution. The unknown Fe (II) sample solution was diluted to mark and was labelled as
“Stock Sample.”
Three clean 50 mL volumetric flask were labelled with “Unknown 1-3” then added with 10 mL of stock sample. The flasks were added with 1.00 mL 10% hydroxylamine hydrochloride solution, 5.00 mL 0.20% 1,10-phenanthroline solution and 1.00 mL of acetate buffer pH 4 successively. The solutions were then swirled and diluted to mark. The last part of the experiment was spectrophotometry to obtain the absorbance spectrum Fe (II) solution from 350 to 600 nm. The optimum wavelength, λmax, can be computed using the resulting absorbance spectrum using the zero flask as reference or blank solution. All the absorbance readings were recorded and was process repeated for the unknown solution.
II.
Results and Discussion
Spectrophotometry was the method in which the absorption of light by the sample was measured. Different compounds absorb different amount of visible light. Naturally, the complimentary color of the solution was used. This absorbance can be measured using a spectrophotometer. The machine would pass a series of wavelengths of light through a solution of a sample substance and also through an identical container which only has the solvent. Light passing through the reference cell was measured for each wavelength passing through the spectrometer. The light passing through the sample was then measured.[1] There were a certain requirements for a species to be analyzed by spectrophotometer. The solution being examined should be colored and diluted before being subjected to the machine. The amount of light absorbed by a solution can be used to compute an unknown concentration of an analyte by getting the absorbance and using BeerLambert’s Law. The Beer-Lambert’s law shows the linear relationship of absorbance and concentration as shown in Equation. 1.[2] A = kbc Equation 1. Beer-Lambert’s Law A = absorbance k = proportionality constant b = path length c = concentration of absorbing species
The law was expressed in terms of absorbance instead of transmittance in order to obtain a linear relationship between the absorbance and concentration making the computation both simple and straightforward since it was directly proportional to the other parameters. This was easier compared to the logarithmic graph produced when transmittance was used instead of absorbance..[2] However, there were several limitations in using Beer-Lambert’ s Law. The law was limited by chemical
and instrumental factors which cause nonlinearity. Some of these were the electrostatic interactions between molecules causing a deviation of absorptivity at high concentration, scattering of light due to particulates in the sample, fluoresecence of the solution, non-monochromatic radiation and stray light. Spectral scanning was the determination of the wavelength of light absorbed maximally by a particular solution commonly known as λmax. The only difference between absorbance and λmax was that absorbance
the ferrous complex and the color not as intense if 1,10Phenanthroline was added first. Hence, a mild reducing agent, namely hydroxylamine hydrochloride, was added before the color was developed to measure the total Fe content of the solution as shown in Equation 3.
Equation 3. Reaction between 1,10-phenanthroline and Fe (II)
[4]
doesn’t necessarily use maximal.
For accurate results, the λmax, maximum wavelength, was taken since the measurement of concentration was most sensitive at the maximum wavelength. The computed λmax for the experiment was 509.2. In addition, the baseline was determined by setting the absorbance of the reference substance as baseline value. Hence, all the resulting absorbance for other substances was recorded relative to the initial zeroed substance. The significance of obtaining the maximum absorbed light was that it determines the real peak wavelengths of the spectra. In addition, it allows the study in spectral shift in accordance to changes in the environment, monitoring the stability of reagent and checking for the purity of the sample. In the experiment, 1,10-phenanthroline, a tricyclic nitrogen heterocyclic compound that reacts with metals as illustrated in Figure 1 was used to produce the complex.
Then the 1,10-phenanthroline solution was added to produce the ferrous complex of [Fe(C12H8N2)3]2+. The phenanthroline was added in excess to ensure that all the remaining iron were complexed and to make sure that the reaction was complete.[6] The absorbing species in the experiment was the [Fe(C12H8N2)3]2+ complex where it absorbs strongly at the maximum wavelength, λ max. Calibration was done before doing the spectrophotometry for the unknown. A critical part of the experiment, spectrophotometer calibration was done to determine if the spectrophotometer was functioning properly and if the measurements were correct. The instrument readings should correlate with the a standard The absorbance was measured from each sample for standardization using a spectrophotometer with double beam. The data obtained were listed in Table I. Obtained Absorbance for Standardization and their Corresponding Concentration Absorbance Concen. A B Average 0.1170 0.1040 0.1105 0.5 0.2050 0.2130 0.2090 1.0 0.3040 0.3050 0.3045 1.5 0.4090 0.4070 0.4080 2.0 0.5050 0.5090 0.5070 2.5
Table I.
Figure 1. Structure of 1,10-Phenanthroline
The reaction yielded a red-colored complex product shown in Equation.2.
Equation 2. Reaction between 1,10-phenanthroline and Fe (II)
In order to determine the total iron in the sample, the reagents was added in sequence. The iron should be in the ferrous state since the o-Phen can form a colored complex with Fe3+ which was not the desired reaction for the experiment. The resulting product would have a different spectrum from
1 2 3 4 5
As shown above, there were two sets of data from Trial A to B. The absorbance obtained from these were averaged and their concentration taken for the plotting of the calibration curve as shown in Figure 3.
0.6 2.5, 0.507
0.5
e c n 0.4 a b 0.3 r o s b 0.2 A
2, 0.408
0.5, 0.1105
0.1 0 0
1
1.5, 0.3045 1, 0.209 y = 0.1984x + 0.0102 R² = 0.9999 2 3
Concentration (ppm)
Figure 3. Plot of Absorbance vs Concentration
The calibration curve was plotted in order to determine the concentration of the unknown. Through excel, the equation for best fit line was given by the Equation 4. y = mx + b Equation 4. Equation for Best Fit Line
the concentration of Fe (II) in the stock sample in ppm listed in Table IV. Calculated Concentration of Fe (II) in the Stock Sample Concentration (ppm) 14.43 1 2 14.45 14.26 3 Average 14.38
Table IV.
The computed average of concentration for Fe (II) in stock sample was 14.38 ppm. This was then used to calculate for the molarity which was 2.60 x 10-4 M. The calculated range was 0.17, relative standard deviation was 0.59% and confident limit was 14.35 ± 0.211. Throughout the experiment, there might be errors made which were listed in Table V. Possible Errors Errors Effect in Concentration There were air bubbles inside Decrease the cuvette Not all Fe (II) in the solution Decrease had complexed The reagents were not added Indeterminate in the correct order Table V.
y = Absorbance m = slope = molar absorptivity x path length x = concentration b = error
The obtained equation was y = 0.1984x + 0.0102 and since the obtained R2 was very close to 1, it suggests that the absorbance and concentration increases in the same proportion. The spectrophotometry for unknown was then carried out and the absorbance listed in Table II. Obtained Absorbance for Unknown Absorbance 1 0.1820 0.1810 2 3 0.1800
Table II.
Using the equation for best fit line, the concentration of Fe (II) solution was calculated as shown in Table III. Calculated Concentration of Unknown Fe (II) Solution Concentration (ppm) 1 0.8659 0.8609 2 0.8558 3 Average 0.8609
Table III.
The average concentration of Fe (II) solution was 0.8609 ppm. This value was then used to calculate
III.
Conclusion
The experiment proved that the concentration of an unknown Fe (II) solution can be obtained using spectrophotometry. By plotting the calibration curve using the measured absorbance and concentration, the equation for best fit line, y = 0.1984x + 0.0102, was obtained. Applying the Beer-Lambert’s Law that relates absorbance as directly proportional to concentration, the linear equation was used to get the average concentration of the unknown Fe (II) solution which was 0.8609 ppm. The concentration of Fe (II) in stock sample was then calculated yielding the answer of 14.38 in ppm and 2.60 x 10-4 in molarity. It can be concluded that the experiment was a success since the calculated R 2 is very close to 1. This means that the absorbance and concentration increases in the same proportion which was the expected trend in the experiment.
IV.
References
[1] Clark, Jim. The Beer-Lambert Law. http://www.chemguide.co.uk/analysis/uvvisible/be erlambert.html. (accessed September 3, 2013) [2] Sheffield Hallam University. Beer’s Law. http://teaching.shu.ac.uk/hwb/chemistry/tutorials/ molspec/beers1.htm (accessed September 3, 2013) [3] The University of Adelaide Australia. Department of Chemistry. Beer-Lambert Law. http://www.chemistry.adelaide.edu.au/external/so c-rel/content/beerslaw.htm (accessed September 3, 2013) [4] Held, Paul. Ph.D. Spectral Scanning Capabilities. http://www.biotek.com/resources/articles/spectro photometer-spectral-scanning-capabilities.html (accessed September 3, 2013) [5] Thermo Fisher Scientific. The Importance of Spectral Scanning and Spectral Analysis for Achieving the Optimal Assay Performance. http://www.thermoscientific.jp/labproducts/microplate-reader/docs/varioskan-flash2.pdf (accessed September 4, 2013) [6] University of Kentucky. Department of Chemistry. Experiment 5: Molecular Absorption Spectroscopy: Determination of Iron With 1,10-Phenanthroline. http://www.chem.uky.edu/courses/che226/labs/05 0-fe_absorption.pdf (accessed September 5, 2013)
V.
Appendices
Sample Calculations
Best Fit Line = y = 0.1984x + 0.0102
Concentration of Unknown Fe (II) Solution U1 = U2 = U3 =
ppm
ppm
ppm
Concentration of Fe (II) in the Stock Sample
U1 = ppm U2 = ppm U3 = ppm M = (14.35 ppm)(
)(
) = 2.60 x 10-4 M
