In the Laboratory
ATR–FTIR Spectroscopy in the Undergraduate Chemistry Laboratory Part II: A Physical Chemistry Laboratory Experiment on Surface Adsorption Jennifer D. Schuttlefield, Sarah C. Larsen, and Vicki H. Grassian* Department of Chemistry, University of Iowa, Iowa City, IA 52242; *
[email protected] In Part I, attenuated total reflectance-Fourier transform infrared (ATR–FTIR) spectroscopy was shown to be a versatile tool for measuring infrared spectra of solids and liquids (1). In addition to measuring the spectra of solids and liquids, ATR– FTIR spectroscopy can be used to investigate the solution-phase surface chemistry of particle thin films (2). Surface chemistry is an area of current interest with applications in environmental chemistry, heterogeneous catalysis, and materials chemistry (3–8). We have designed an experiment to investigate the adsorption of sulfate ions from solution onto TiO2 particles that can be easily incorporated into an upper-level chemistry laboratory course. This experiment introduces students to several important concepts including surface adsorption, metal oxide surface chemistry, and environmental interfacial processes. Experimental Procedure An experiment designed to investigate the adsorption of sulfate ions, SO42−, onto TiO2 particle surfaces was introduced into the undergraduate physical chemistry laboratory. This experiment introduces students to the technique of ATR–FTIR spectroscopy and the Langmuir model for surface adsorption. The experiment was conducted in two, four-hour laboratory periods. The first day consisted of making Na2SO4 solutions, adjusting the pH to 3.0, and measuring the absorbance of the following sulfate solutions to obtain a calibration curve: 10, 20, 40, 60, 80,
and 100 mM. Infrared spectra were obtained using a Nicolet spectrometer (Nexus model 670) equipped with a DTGS KBr detector, where 150 scans were acquired at an instrument resolution of 4 cm‒1 over the spectral range between 650 and 4000 cm‒1 owing to the frequency cutoff of the ATR–FTIR internal reflection element (IRE) used. The most intense band in the infrared spectrum, the ν3 band of the sulfate ion, which occurs at a peak maximum of 1102 cm‒1, was used to determine the calibration curve (7, 9). Figure 1A shows representative studentacquired infrared spectra in the ν3 absorption band region at 1102 cm‒1 as a function of increasing sulfate ion concentration. From these data, a calibration curve was obtained as shown in Figure 1B. The calibration curve was used to determine dp, the effective path length as described in Part I (1), and was also used to subtract out contributions from solution-phase sulfate absorptions from the spectra that contain contributions from both solution-phase and adsorbed sulfate absorptions. After the students obtained a calibration curve, a slurry of TiO2 and water (Fisher, Optima) was placed on the ZnSe (Pike Technologies, Trough Plate, part #022-2010-45) IRE and the water was allowed to evaporate overnight. Prepared this way, the TiO2 layer is compact and adheres to the ZnSe IRE allowing the sulfate solution to be introduced and flushed through the horizontal cell without loss of the TiO2 sample. On the second day, students examined the adsorption of sulfate ions onto the TiO2 surface. Sulfate solutions of 0.01, 0.02, 0.05, 1.0, 4.0, and 5.0 mM were used. ATR–FTIR spectra
B
A 0.01
Absorbance
1102
[SO42ź]
1200
1150
1100
1050
Wavenumber / cmź1
1000
Absorbance (1102 cmź1)
0.05
0.04
y = 0.00038x + 0.0005 0.03
0.02
0.01
0.00
0
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80
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Concentration / mM
Figure 1. (A) Representative ATR–FTIR spectra in the spectral range from 975 to 1225 cm−1 for sulfate solutions ranging in concentration from 10 to 100 mM (10, 20, 40, 60, 80, and 100 mM). The sulfate absorption band is observed at 1102 cm−1. (B) The calibration curve obtained from the data shown in (A).
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Journal of Chemical Education • Vol. 85 No. 2 February 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory A 1161
Langmuir Adsorption Constant
observed absorbance of total absorbance of solution phase adsorbed sulfate absorbance
R
observed absorbance of adsorbed sulfate
1200
1150
1100
1050
1000
950
900
Wavenumber / cm B
1.2
1.0
0.8
0.6
0.4
0.2
0.0 0
1
2
3
4
5
6
Concentration / mM C
6
y = (4.46 × 10ź5)x + 1.0275 R2 = 0.9993
5
(1)
Before plotting an isotherm, it is necessary for students to convert adsorbed-sulfate absorbance into coverage. This is done by determining where the adsorbed-sulfate absorbance levels off, indicating the presence of one monolayer. Once this is determined the fractional coverage, θ, is calculated by
1250
ź1
4
1/R
1070
[SO42ź]
1300
Coverage
The students then analyzed their data using the Langmuir adsorption model (11) to determine the Langmuir adsorption constant, Kads, for sulfate adsorption on TiO2. Students were given a handout that included a discussion of the assumptions behind the derivation of the Langmuir adsorption model: (i) there are distinct sites on the surface; (ii) all of the sites on the surface are equivalent; (iii) there are no interactions between adsorbed molecules; and (iv) only one molecule can adsorb per surface site (11, 12). Thus, the value of the coverage, θ, is such that it can vary from 0 to a maximum value of 1 and a maximum of one monolayer can adsorb on the surface (vide infra). Once a value of θ = 1 is obtained, no more sulfate can adsorb and there is a plateau in the adsorption isotherm. The students were shown a typical adsorption isotherm (i.e., coverage, θ, versus concentration) and then were instructed to determine an isotherm from their data. Figure 2B is an example of an adsorption isotherm obtained by the students in their experiment. To create an isotherm, students had to use the calibration curve that was obtained on the first day of the experiment. Students were instructed to create a spreadsheet with the following columns: concentration, total absorbance, solution-phase sulfate absorbance, and adsorbed-sulfate absorbance. The total absorbance was taken as the absorbance at the peak maximum from the sulfate ion spectra in the presence of a TiO2 thin film. The peak maximum was used to remain consistent with day 1 experiments in which the maximum peak absorbance for the solution phase was used to calculate the effective path length. The solution-phase absorbance was determined from the calibration curve (shown in Figure 1B) from the first day of the experiment. This analysis assumes that the absorbance is a linear sum of two components, adsorbed sulfate and solution-phase sulfate:
1111 0.02
Absorbance
as a function of increasing concentration were collected by the students. Representative spectra are shown in Figure 2A where the sulfate ν3 band is shifted to 1111 cm‒1 upon adsorption. In addition, the absorption band has shoulders at 1070 and 1161 cm‒1 that are due to a reduction in symmetry that occurs upon sulfate adsorption on the TiO2 particle surfaces (10).
3
2
1
0 0
2
4
6
8
10
12
(1/[SO42ź]) / (104 Mź1) (2)
maximum absorbance of adsorbed sulfate
After the students plotted concentration versus θ, they were asked to determine Kads. The adsorption constant, Kads, is defined
Figure 2. (A) Infrared spectra of the sulfate ion adsorbed on TiO2 as a function of sulfate ion concentration (0.01, 0.02, 0.05, 1.0, 4.0, and 5.0 mM). The spectral region from 900 to 1300 cm−1 is shown. (B) An adsorption isotherm as a function of sulfate ion concentration is constructed from the data shown in Figures 1A and 2A. (C) From the linearized form of the Langmuir adsorption equation, where 1/θ is plotted versus 1/[SO42−], Kads can be determined from the slope of the line.
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 2 February 2008 • Journal of Chemical Education
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In the Laboratory
as the equilibrium constant according to K ads
R SO 4
2
1
R
Conclusions
(3)
and as discussed above, θ is defined as the fractional coverage and is equal to the number of adsorbed molecules, Nads, divided by the total number of adsorption sites, Nsites, available on the surface: N R ads N sites
(4)
We have successfully implemented a laboratory experiment that allows students to use ATR–FTIR spectroscopy to investigate the adsorption of ions at the aqueous–solid interface. In this experiment, students determine Kads for the adsorption of sulfate ions onto the surface of TiO2 particles. This experiment can also be expanded to investigate other inorganic anions such as the phosphate ion or the adsorption of organics on to the surface of TiO2. Acknowledgments
The maximum number of molecules that can adsorb, Nmax, is equal to the total number of adsorption sites available on the surface, thus the coverage, θ, can vary from 0 to 1. The term (1 − θ) in the denominator of eq 3 of gives the fractional number of sites that remain open on the surface (11). Equation 3 can be rearranged in terms of the coverage, θ, which shows the coverage dependence of adsorbed sulfate as a function of sulfate concentration in solution. This dependence is
This material is based upon work supported by the National Science Foundation under Grant No. 0503854. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the author(s) and do not necessarily reflect the views of the National Science Foundation. The authors would like to thank teaching assistant Michael Nydegger, and all students in the Physical Measurements class during the fall of 2004 and 2006 that were involved in testing this experiment over the course of its development.
Literature Cited
R
SO 4 2 Kads 1 SO 4 2 K ads
(5)
The students determined Kads for sulfate adsorption by using a linearized form of eq 5:
1 R
SO 4
1
2
K ads
1
(6)
A plot of 1/θ versus 1/[SO42−] yields a straight line with a slope of 1∙Kads. Figure 2C illustrates a plot taken from student data of 1/θ versus 1∙[SO42−]. From Figure 2C, a Kads value of 2.2 × 104 M‒1 was determined. Typical values for Kads determined by the students were on the order of 104 to 105 M‒1. Hug and Sulzberger (7) used principal component analysis and a two-site model to describe the adsorption of sulfate on TiO2 particle surfaces that yielded values of Kads,1 = 3.1 × 103 M‒1 and Kads,2 = 1.8 × 105 M‒1. For simplification purposes, a one parameter fit with a solution-phase calibration curve was used here to analyze the data. By implementing this experiment, we have shown the use of ATR–FTIR spectroscopy in the undergraduate advanced laboratory as a technique to study surface adsorption at the aqueous–oxide particle interface. Hazards Hydrochloric acid is corrosive, causes severe burns to all body tissue, and may be fatal if swallowed or inhaled. Titanium dioxide may cause irritation to skin, eyes, and respiratory tract. Students are instructed to collect waste (i.e., dilute solutions at a pH of 3 of sodium sulfate) in specific containers that will be collected by the Health Protection Office at the University of Iowa or disposed of properly by their instructors. The ZnSe ATR crystal should be handled with care and a pH lower than 3 should not be used.
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1. Schuttlefield, J. D.; Grassian, V. H. J. Chem. Educ. 2008, 85, 279–281. 2. Hind, A. R.; Bhargava, S. K.; McKinnon, A. Adv. Colloid Interface Sci. 2001, 93, 91–114. 3. Lefevre, G. Adv. Colloid Interface Sci. 2004, 107, 109–123. 4. Stumm, W. Chemistry of the Solid-Water Interface: Processes at the Mineral-Water and Particle-Water Interface in Natural Systems; John Wiley & Sons, Inc.: New York, 1992. 5. Stumm, W.; Morgan, J. J. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, 3rd ed.; John Wiley & Sons, Inc.: New York, 1995. 6. Turchi, C. S.; Ollis, D. F. J. Catal. 1990, 122, 178–192. 7. Hug, S. J.; Sulzberger, B. Langmuir 1994, 10, 3587–3597. 8. Wang, X.; Yu, J. C.; Liu, P.; Wang, X.; Su, W.; Fu, X. J. Photochem. Photobiol., A 2006, 179, 339–347. 9. Ross, S. D. Inorganic Infrared and Raman Spectra; McGraw-Hill: London, 1972; pp 204–212. 10. Peak, D.; Ford, R. G.; Sparks, D. L. J. Colloid Interface Sci. 1999, 218, 289–299. 11. Atkins, P.; de Paula, J. Atkins’ Physical Chemistry, 7th ed.; W. H. Freeman and Company: New York, 2002; pp 989–1005. 12. Kolasinski, K. W. Surface Science: Foundations of Catalysis and Nanoscience; John Wiley & Sons, Ltd: New York, 2002; pp 168–202.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/Feb/abs282.html Abstract and keywords Full text (PDF) with links to cited JCE articles Supplement
Student handout
PowerPoint presentation used in the class lecture
Journal of Chemical Education • Vol. 85 No. 2 February 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
