In the Laboratory
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An Experiment on Photochromism and Kinetics for the Undergraduate Laboratory Hernán E. Prypsztejn Escuela Técnica No. 1, Otto Krause, Av. Paseo Colón 650, 1063 Buenos Aires, Argentina R. Martín Negri* Departamento de Química Inorgánica, Analítica y Química Física, Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires, Ciudad Universitaria, Pabellón 2, 1428 Buenos Aires, Argentina;
[email protected] The Kinetic Process The thermodynamically more stable isomer of 6-NO2BIPS is the closed-ring isomer, referred to as the normal form, N. It has an absorption band at the near UV region, around 350 nm, and does not absorb in the visible. The open-ring isomer has the molecular structure of a cyanine dye and is referred to as the merocyanine form, MC (see structures below). O−
H3C
CH3 h ν (UV)
N
O
CH3
NO2
H3C
CH3
∆
N+ CH3
N isomer
MC isomer
NO2
1.2 1.0
Absorbance
Photochromism is the process of inducing color changes in a medium by action of incident electromagnetic radiation (1). Although many compounds exhibit photochromic behavior (2), the family of spiropyrans is one of the most frequently used in research and technological applications. The spiropyrans, such as 1′,3′-dihydro-1′,3′,3′-trimethyl-6-nitrospiro(2H-1-benzopyran-2,2′-2H-indole) (referred to as 6-NO2BIPS, Fig. 1), have been used as photochromic glasses (3) and optical memory devices (4 ) and in studies of matrix relaxation in polymers (5). The synthesis (3, 6 ), thermochromism (7 ), and some applications (8) of related spiropyrans have been reported in this Journal. This work presents an experiment on photochromism and photochemical kinetics for undergraduate physical chemistry laboratory courses. It allows students to determine parameters for the decoloration process including relaxation times, activation energy, and entropy, and also ∆H ° for the thermal equilibrium between the two isomers involved in the photochromic process. This laboratory experiment has been performed since October 1998 as part of the course Chemical Physics II for undergraduate chemistry students at the Faculty of Sciences (FCEyN), University of Buenos Aires, Argentina. The experiment is also performed at a lower level by students of two technical high schools of Argentina, which integrates the Chemistry, Humankind, and Our Environment project (9). The simplicity, speed, and low cost of an experiment dealing with relatively complex chemical concepts is the rationale for adopting it.
0.8 0.6
a
0.4
c
0.2 0 250
b 300
350
400
450
500
550
600
650
700
Wavelength / nm Figure 1. Absorption spectra of 6-NO2-BIPS in ethanol and toluene. (a) In ethanol immediately after irradiation; the unlabeled spectra were recorded at several times afterwards. (b) In ethanol before irradiation. (c) In toluene after irradiation. The concentrations in ethanol and toluene are not matched.
MC also absorbs in the UV, but unlike N, it shows a strong and characteristic absorption band in the visible spectral region between 500 and 600 nm (Fig. 1). In the absence of UV radiation, N is the more abundant isomer in solution, although MC is also present in thermal chemical equilibrium with N. When a solution of 6-NO2-BIPS is excited with UV light at the absorption band of N, the thermal equilibrium between N and MC is displaced, and the MC concentration and color intensity of the solution increase. For this reason MC is generally called the photo isomer. After irradiation MC returns thermally to N, a process that can easily be followed by measuring the visible absorption of MC as a function of time. In this way the kinetics of decoloration can be completely characterized. In addition, ∆H ° for the thermal equilibrium can be determined if a UV–vis spectrophotometer with a constant-temperature sample compartment is used. Materials and Methods
Chemicals 1′,3′-Dihydro-1′,3′,3′-trimethyl-6-nitrospiro(2H-1benzopyran-2,2′-2H-indole) was purchased from SigmaAldrich and used as received. Analytical grade ethanol, toluene, and methylcyclohexane were from Merck. UV Source The UV excitation source is a commercial photographic flash, used as an external flash unit.
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In the Laboratory
Absorption Measurements A Shimadzu UV-3101PC spectrophotometer with a constant-temperature sample compartment was used. A single-beam spectrophotometer, such as a Spectronic 20, can be used to follow the kinetics when the absorbance decay is relatively slow (on the order of minutes), as in the case of solutions in ethanol. Plastic cuvettes of 10-mm path length can be used for the kinetics measurements. Quartz cuvettes must be used when the absorption of the UV band of N is recorded for calculations of ∆H °. Temperature Measurements Temperature is controlled by external water circulation and measured by a calibrated thermistor. Temperatures between 25 and 70 °C were assayed. Time Measurements Time is measured with a laboratory chronometer. Time zero is arbitrary defined by the operator.
0 -0.5
ln(A(t ) - A i)
-1.0 -1.5 -2.0 -2.5 -3.0 -3.5
0
10
20
30
40
50
60
Time / min Figure 2. Ln (A(t) – Ai) vs time in ethanol at 25 ºC. Ai = 0.022. Absorbances were measured at 540 nm.
4 3 2
ln(τ / min)
Sample Preparation Solutions were prepared by taking a few milligrams of 6-NO2-BIPS with a capillary and dissolving it in about 15 mL of solvent (ethanol or toluene). No volumetric glassware is required. Solutions in ethanol must be prepared one day in advance because 6-NO2-BIPS dissolves slowly in it. Solutions can be stored at room temperature in the dark for a week.
1 0 -1 -2
Hazards
-3 2.9
Because the plastic cover of the UV excitation source, which normally filters the UV light, must be removed, it is necessary to use plastic protection glasses and laboratory gloves. Results and Discussion
Kinetic Measurements The molecular rearrangement leading from the excited state of N to MC is in the nanosecond or picosecond time scale (10), too fast to be observed with a commercial spectrophotometer. Therefore the formation of MC can be considered to be an instantaneous process for the time resolution of this experiment. After irradiation the visible band is observed to be at its maximum absorbance. The spectral changes in ethanol at 25 °C before irradiation and at different times afterward are shown in Figure 1. Excellent fits of the data were obtained using single exponential decay functions, as shown in Figure 2, confirming a first-order decay. Similar behavior was observed in ethanol at all temperatures assayed. Therefore, the time evolution of the absorbance after irradiation, A(t), can be described by A(t) = (A0 – Ai )e᎑t /τ + Ai
(1)
where τ represents the relaxation time of the process, that is, the reciprocal of the rate constant for the thermal back-isomerization process. A0 is the absorbance of MC at time zero (after irradiation) and Ai is the absorbance of MC before irradiation, measured at the same wavelength (for example, at 540 nm).
646
3.0
3.1
3.2
3.3
T -1 / (10-3 K-1) Figure 3. Ln τ vs 1000/T in ethanol. The calculated activation energy is 112 ± 2 kJ/mol.
The values of τ decrease very rapidly when the temperature increases (from 31 min at 25 °C to 6.2 s at 70 °C), indicating a relatively high activation energy (Ea) for the thermal backisomerization. Linear Arrhenius plots were obtained (Fig. 3) when data at different temperatures were fitted by τ = A exp(E a/RT )
(2)
From Figure 3 the values Ea = (112 ± 2) kJ/mol and ln(A/min) = (᎑ 41.9 ± 0.8) in ethanol were obtained. The activation enthalpy and entropy, ∆H # and ∆S #, can be estimated using the activated complex theory for unimolecular reactions in solution: ∆H # = E a – RT
(3)
∆S = ᎑R ln(A e k T /h)
(4)
#
where the symbols R, h, and k represent the gas, Planck, and Boltzmann constants, respectively. The values obtained were ∆H # = (110 ± 2) kJ/mol and ∆S # = (61 ± 7) J K ᎑1 mol ᎑1 in ethanol at 25 °C .
Equilibrium Measurements In absence of UV radiation, a thermal equilibrium between N and MC is established in ethanol. The value of ∆H ° for
Journal of Chemical Education • Vol. 78 No. 5 May 2001 • JChemEd.chem.wisc.edu
In the Laboratory
the thermal equilibrium can be estimated by measuring the absorbance in the UV and visible regions at different temperatures, if some assumptions are made: 1. ∆H ° does not change with temperature in the range considered. 2. The absorbance at 337 nm (A 337) is assumed to be exclusively due to N. The contribution of MC to the absorbance is negligible at the shoulder of the UV band (337 nm). 3. Molar absorption coefficients do not change with temperature, as is the case for related compounds such as merocyanine 540 and symmetric carbocyanine dyes (11). 4. Unitary activity factors are assumed.
Under these assumptions, the van’t Hoff equation
d ln K eq dT
=
∆H ° RT 2
(5)
predicts a linear relationship for ln(A540/A337) vs 1/T when Keq is expressed in terms of the ratio of absorbances A540/A337 using the Beer–Lambert law (A540 and A337 are the absorbances at 540 nm and 337 nm, respectively, in the absence of UV irradiation). The slope of ln(A540/A337) vs 1/T is equal to (᎑∆H °/R). The straight line obtained in ethanol in the range from 25 to 70 °C confirms the reasonability of the assumptions (figure not shown). ∆H ° = (9.8 ± 0.2) kJ/mol was obtained.
Solvent Effects The photochromism of 6-NO2-BIPS gives an excellent opportunity to show very strong solvent effects. The spectral position of the visible absorbance band, the color of the solution, and the measured relaxation time (τ) are all seen to depend on the solvent used. The 6-NO2-BIPS solutions after irradiation are blue in toluene, pink in ethanol, and violet in cyclohexane. The relaxation times at 25 °C calculated from our data are 11 s in toluene and 31 min in ethanol. The different relaxation times observed in ethanol and toluene can be interpreted in terms of the effect of polarity on the activation barrier. It has been proposed that the colored form (MC) is a zwitterion and is more polar than N and the transition state (12). In this case the more polar solvent should stabilize MC relative to both the transition state and N, increasing the activation barrier. The slower relaxation time in ethanol than in toluene is in complete agreement with this model. This observation could be the basis of an expansion of our experiment to include solvent effects more specifically. For example, the effect of solvation stabilization on the energy of ground and excited states of a different merocyanine dye has been reported in this Journal (13). Changes in the following properties can be analyzed using at least two different solvents: color of the solution, difference of energy between the ground and excited states, Keq, and τ at room temperature. Other solvents can also be used to demonstrate solvent effects. For example, solutions of 6-NO2-BIPS in methylcyclohexane are violet after irradiation. However, chlorinated solvents cannot be used to demonstrate solvent effects. Solutions in chloroform become yellow after irradiation, probably owing to a photochemical reaction with the solvent.
Conclusions This experiment allows students to deal with theoretical and experimental aspects of photochemical processes, which are illustrated in a very simple manner. The preparation and manipulation of the samples is also very quick and simple. The experiment is visually appealing owing to the color change upon irradiation, which is strongly dependent on the solvent used. In our experience, these factors increase students’ interest in excited-state photophysical processes and photochemical reactions. The experiment also provides an excellent opportunity to make links with related subjects such as solvent effects, reactivity of organic compounds, double-bond conjugation, and color theories, which are discussed with the students. The experiment is similar to one reported through Project SERAPHIM (14 ), which uses mercury(II) dithizonate, the thermochromism of which was discussed in this Journal (2d ). We feel that our experiment has two advantages: it does not have the disposal problems of Hg salts and it offers a more natural extension into solvent effects. The determination of τ at room temperature was also performed by high school students at two technical schools in Buenos Aires (Escuela Técnica No. 1 Otto Krause and Escuela Técnica No. 27 Hipólito Irigoyen). Their results are in excellent agreement with those presented here. The experiment allows the introduction and demonstration of basic concepts of photochemistry at the high-school level in a simple way. Exposing a sample to sunlight can induce the photochromism, which is visually attractive for the students. Acknowledgments The experiment was developed within the context of the project Chemistry, Humankind, and Our Environment (9), a collaboration of four high schools and the University of Buenos Aires with the purpose of improving chemistry teaching at the high school level. The Chemistry, Humankind, and Our Environment project is supported by Fundación Bunge & Born and Fundación YPF (Argentine), which are acknowledged. The University of Buenos Aires (grant JX-48) also supported the work. We thank Horacio Corti, director of the project, for his constant support and encouragement. We thank one of the reviewers for the suggestions to improve the manuscript and for the reference to the SERAPHIM material. W
Supplemental Material
Supplemental material for this article is available in this issue of JCE Online. Experimental procedures are presented in detail, including costs and CAS registry numbers of chemicals used. A demonstration of the change of color induced by UV light is also available. Literature Cited 1. Dürr, H.; Bouas-Laurent, H. Photochromism. Molecules and Systems; Studies in Organic Chemistry 40; Elsevier: Amsterdam, 1990; p 5. Bertelson, R. C. In Techniques in Chemistry: Vol. 3, Photochromism; Brown, C. H., Ed.; Wiley-Interscience: New York, 1971; Chapter 1.
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In the Laboratory 2. (a) Ault, A.; Kouba, C. J. Chem. Educ. 1974, 51, 395. (b) Zaczek, N. M.; Levy, W. D.; Jordan, M. L.; Niemyer, J. A. J. Chem. Educ. 1982, 59, 705. (c) Petersen, R. L.; Harris, G. L. J. Chem. Educ. 1985, 62, 802. (d) Hutton, A. T. J. Chem. Educ. 1986, 63, 888. (e) Pickering, M. J. Chem. Educ. 1980, 57, 833. 3. Osterby, B.; McKelvey, R. D.; Hill, L. J. Chem. Educ. 1991, 68, 424. 4. Dvornikov, A. S.; Malkin, J.; Rentezepis, P. M. J. Phys. Chem. 1994, 98, 6746. 5. Smets, G. Adv. Polym. Sci. 1982, 50, 576. Richert, R.; Bässler, M. Chem. Phys. Lett. 1985, 118, 534. Levitus, M.; Talhavini, M.; Negri, R. M.; Zambon Atvars, T. D.; Aramendía, P. F. J. Phys. Chem. B 1997, 101, 7680. 6. Guglielmetti, R.; Meyer, R.; Dupuy, C. J. Chem. Educ. 1973,
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50, 413. 7. Byrne, J. P. J. Chem. Educ. 1978, 55, 267. 8. Chambers, K. W.; Smith, I. M. J. Chem. Educ. 1974, 51, 354. Jacques, P. J. Chem. Educ. 1991, 68, 347. 9. Announcements: Chemistry, Humankind, and Our Environment ; J. Chem. Educ. 1998, 75, 1068. 10. Zhang, J. Z.; Schwartz, B. J.; King, J. C.; Harris, C. B. J. Am. Chem. Soc. 1992, 114, 10921. 11. Aramendía, P. F.; Negri, R. M.; San Román, E. J. Phys. Chem. 1994, 98, 3165. 12. Flannery, J. B. Jr. J. Am. Chem. Soc. 1968, 90, 5660. 13. Shah, S. S.; Henscheid, L. G. J. Chem. Educ. 1983, 60, 685. 14. Project SERAPHIM is described at http://jchemed.chem. wisc.edu/ChemEd/Orgs/org26.html (accessed Jan 2001).
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