In the Laboratory
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An Easy and Versatile Experiment to Demonstrate Solvent Polarity Using Solvatochromic Dyes Clodoaldo Machado* Departamento de Química, Universidade Regional de Blumenau, FURB, CP 1507, Blumenau, SC 89010-971, Brazil;
[email protected] Vanderlei Gageiro Machado Departamento de Química, Universidade Federal do Paraná, UFPR, CP 19081, Curitiba, PR 81531-990, Brazil
Physicochemical studies of many chemical processes have revealed that the solvents may exert strong influence on chemical equilibria and reaction rates and on the position and intensity of the spectroscopic absorption bands (1). Attempts to explain these solvent effects in terms of the “polarity” of the medium have demonstrated the complexity of this concept, although it can be well understood qualitatively (remember the oldest rule in chemistry: “like dissolves like”). For undergraduate students, the concept of solvent polarity is always connected directly to solvent permanent dipole moment (µ) or solvent dielectric constant (ε). However, this simple approach to solvation, based only on physical parameters, cannot be applied to many solvent-dependent processes, because not only do these nonspecific interactions (related to µ, ε, and refractive index) take place in solute–solvent systems, but specific interactions also occur. These latter interactions, such as electronpair donor–electron-pair acceptor activity or solvophobic and hydrogen bonding interactions, contribute to the observed failure. In consideration of these facts, Reichardt proposed the use of the term polarity for the “overall solvation capability” in a given process, as a result of the sum of both specific and nonspecific interactions, disregarding the interactions that provoke definite chemical alterations caused by a reaction between the solute and the solvent (2). Other parameters of solute–solvent interactions were proposed to circumvent these difficulties in defining solvent polarity based on a microscopic view. Many compounds show a significant change in the position of a UV–vis absorption band when the solvent is changed. This phenomenon, known
Figure 1. (a) Color strip associated with wavelength; (b) complementary color wheel; (c) dye 1 in different solvents; left to right: methanol, ethanol, 1-propanol, 2-propanol, DMSO, acetone, chloroform. (This figure appears in color in the table of contents and in JCE Online).
as solvatochromism, made many of these compounds popular in the study of solvent effects. Among them, Reichardt’s betaine 2,6-diphenyl-4-(2,4,6-triphenyl-1-pyridinio)-1-phenoxide (1) (see structures), also named betaine ET(30), is the most popular, and it has been employed to quantify the polarity of different media. When used for pure solvents, the well-established ET(30) polarity scale is derived (3). Another classical solvatochromic dye is Brooker’s merocyanine 4-[(1-methyl4(1H)-pyridinylidene)ethylidene]-2,5-cyclohexadien-1-one (2) (4 ). These compounds are very sensitive to small changes in the polarity of the medium, as can be observed visually as changes in the color of the solution. Thus, solutions of dye 1 are green in acetone, blue in 2-propanol, violet in ethanol, and red in methanol (Fig. 1). N
N
O 1
Betaine E T(30)
O 2
Brooker’s merocyanine
A quick classroom demonstration of solvent polarities using an overhead projector and solutions of dye 1 has been described (5). Here, we describe a physical chemistry experiment in which the student explores the solute–solvent interactions by means of solvatochromic probes 1 and 2. The discussion is based on the color correlation of the solutions with the media polarity values. The facile synthesis (6 ) of dye 1 and the minimal amount required to perform experiments justify the use of this compound in the experiment. The convergent synthesis of this dye has been described for an advanced undergraduate project (7). Dye 2 is included because of its very simple synthesis (4, 8) and for the purpose of comparing the two polarity scales to be obtained in the experimentation. Dyes 1 and 2 are also commercially available. Initially, visual observations are used to put the solvents into order of increasing polarity. These observations are compared with the UV–vis measurements. Therefore, by using this procedure, it is possible not only to obtain a quantitative polarity scale, but also to show the correlation of transition energy with observed colors. The influence of the addition
JChemEd.chem.wisc.edu • Vol. 78 No. 5 May 2001 • Journal of Chemical Education
649
In the Laboratory
of a more polar solvent or an electrolyte to a pure solvent is also shown. Finally, this experiment is interesting because it not only furnishes complementary theory on the subject of solvent polarity but it also opens up an interesting debate on the complexity of solute–solvent interactions. Experimental Method Equipment The following equipment is needed for these experiments: UV–vis spectrophotometer, analytical balance, microsyringe, assay tubes, and 5-mL volumetric flasks.
Chemicals The following chemicals are needed for these experiments: water, methanol, absolute ethanol, 1-propanol, 2-propanol, acetone, acetonitrile, chloroform, dimethyl sulfoxide (DMSO), N,N′-dimethylformamide (DMF), ethyl acetate, lithium perchlorate, sodium iodide, and dyes 1 and 2. Procedure Solutions of the dyes 1 and 2 with concentrations around 1 × 10᎑4 M are prepared in each solvent (5 mL). After the homogenization of the solutions, the solvents can be ordered by use of a complementary color wheel and a color strip (Fig. 1). Afterward, the maxima in the visible region wavelength (λmax) are determined for each solution with a UV–vis spectrophotometer. Transition energy values (ET, in units of kcal mol ᎑1) for the solvatochromic band of the dyes in each solvent are calculated by means of eq 1, E T = h c NA/ λmax = 28,591/ λmax
ET(30)
N
N
(1)
in which h is Planck’s constant, c is the speed of the light in a vacuum, and NA is Avogadro’s constant. The values obtained with the solutions of dyes 1 and 2 allow the construction of two semiquantitative polarity scales, since the solvents can be ranked in a determined order (e. g., from the most polar [water] to the most apolar). Afterward, these data are compared with the visual scale. The experiment is substantially enriched by two other assays: the addition of very small amounts of water to the solutions of dyes 1 and 2 in acetone, and the addition of small volumes of LiClO4 or NaI to the acetone (10᎑3 M). Hazards Considering the low concentration of dyes and the small volumes of solvent required, there are no significant hazards involved in this experiment. Students should follow customary safe laboratory procedures. Discussion The maxima in the wavelengths for dyes 1 and 2 were measured for various solvents and they were transformed into ET values via eq 1. These data are presented in Table 1, which shows the increasing polarities from ethyl acetate to water for the solvents studied. The collected transition energies for dye 1 constitute part of the polarity scale E T(30) (1). First, it is important to compare the transition energy values with the physical parameters ε and µ (Table 1) that are responsible for the nonspecific interactions between the solute and the media. 650
There is no correlation between E T values and ε or µ because empirical parameters such as E T(30) reflect a combination of specific and nonspecific interactions. We therefore have a plausible explanation for why these spectroscopic empirical parameters are very convenient in the establishing of relationships with spectral, thermodynamic, and kinetic data in many cases where constants such as ε and µ fail to correlate (9). An example of this is the interesting relationship between E T(30) values and the constructed scale for Brooker’s merocyanine. If a plot of E T(30) values as a function of the transition energies for dye 2 in the same solvents is made, a linear relationship between the two scales (r > .99) may be observed. The transition energies in the dyes studied involve the passage from a more dipolar zwitterionic ground state to a less dipolar transition state. In other words, the ground state dipolar moments (µg) for the dyes are greater than their excited state dipolar moments (µe) (Scheme I) (1). As a consequence, the stabilization of the ground state in the dye by more polar solvents is greater than for the transition state, increasing the transition energies. This is the basis for the negative solvatochromism registered for the two dyes studied, and it is opposite to the trends observed with dyes that display positive solvatochromism (µg < µe) (1).
O
O
µg > µ e
ET(30)
nonpolar solvent
ET(30)
polar solvent
solvent polarity increases
Scheme I
Even in ideal solvent mixtures, ET values for solvatochromic dyes are often nonlinear with the molar fraction of the more polar solvent. This lack of linearity results from the dye–solvent nonspecific and specific interactions previously described. These interactions are responsible for local inhomogeneities induced by the dye at microscopic-to-molecular level, a phenomenon known as preferential solvation (1, 10). Investigations involving mixed solvents show that the addition of water to acetone strongly diminishes the λmax values for the dyes studied. This is shown in Table 2 for dye 1. The observed deviations from linearity reveal that the dyes are preferentially solvated by water, the most polar solvent, in the acetone-rich mixtures.
Journal of Chemical Education • Vol. 78 No. 5 May 2001 • JChemEd.chem.wisc.edu
In the Laboratory Table 1. Comparison of Polarity Scales Obtained by UV–Vis Spectroscopy and Values of and ⑀ µ/ εb Solvent ET(30)a ET(Brooker)a (1030 C ᎑1 m᎑1) b
different subjects. A variety of dyes can be synthesized in experimental organic courses, and their solvatochromic properties can be studied. Inasmuch as many inorganic complexes are solvatochromic (12, 13), they can be prepared and studied physicochemically in inorganic chemistry courses. Dye 1 was proposed as a simple and precise tool for quantifying the presence of water in organic solvents, being adequate for use in analytical chemistry experiments (14). Biochemical experiments can be devised with solvatochromic dyes that are able to demonstrate enzymatic dielectric properties (15).
Water
63.1
64.7
5.9
78.30
Methanol
55.4
59.2
5.7
32.66
Ethanol
51.9
56.6
5.8
24.55
1-Propanol
50.7
54.6
5.5
20.45
2-Propanol
48.4
52.5
5.5
19.92
Acetonitrile
46.6
50.2
11.8
35.94
DMSO
45.1
49.8
13.5
46.45
Acknowledgments
DMF
43.2
49.1
10.8
36.71
Acetone
42.2
48.9
9.0
20.56
Chloroform
39.1
46.1
3.8
4.81
We are grateful to FURB and CNPq for financial support and to the reviewers for their comments and valuable suggestions.
Ethyl acetate
38.1
—c
6.1
6.02
a
Transition energy for the solvatochromic band of the dye, in kcal mol , calculated according to eq 1. b Values taken from ref 1a. Dielectric constants are given at 25 °C except for chloroform (20 °C). c Insoluble.
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Supplemental Material
᎑1
Table 2. Effect of Addition of Water on Solvatochromic Band of Dye 1 in Acetone Water (Vol %)
Color
λmax/nm
ET(30)/(kcal mol )
᎑1
0
blue-green
678
42.2
5
blue-green
653
43.8
10
blue
633
45.2
20
blue
599
47.7
30
violet
575
49.7
40
violet
560
51.1
50
purple
550
52.0
The hypsochromic shift provoked by addition of salts to the solutions of dyes 1 and 2 in acetone or other solvents, known as negative halochromism, is well documented (11, 12). Charge-transfer absorptions in these dyes depend on the electronic affinity of the pyridinium electron acceptor and the ionization energy of the phenoxide electron donor. Thus, when salts such as LiClO4 and NaI are added to solutions of 1 and 2 in acetone, the ion-pair association between the phenoxide donor in the dyes and the cation of the salts raises the ionization energy of the electron-donor moiety (2). In other words, the cations stabilize the ground state of these dyes better than their excited state. As a result, the energy transitions between the two states increase and salt-induced hypsochromic shifts are observed. The dye–Li+ association is stronger than the same interaction with Na+, and this result may be rationalized by the fact that Li+ presents the highest effective charge. The experiment described here can be adapted into a demonstration to enrich students’ concept of polarity in upper-level chemistry courses, and consequently to open the discussion of the solute–solvent interactions. Comprehension of solute–solvent interactions is required for the understanding of all processes occurring in solutions. This experiment can also form one of a group of experimental modules involving
A more detailed description of this experiment with examples of questions for discussion and sample student data sheets is available in this issue of JCE Online. Literature Cited 1. (a) Reichardt, C. Solvents and Solvent Effects in Organic Chemistry, 2nd ed; VCH: Weinheim, 1988; pp 285–405. (b) Reichardt, C. Chem. Rev. 1994, 94, 2319–2358. (c) Suppan, P.; Ghoneim, N. Solvatochromism; Royal Society of Chemistry: Cambridge, 1997; Chapter 3. 2. Reichardt, C. Chem. Soc. Rev. 1992, 147–153. 3. Dimroth, K.; Reichardt, C.; Siepmann, T.; Bohlmann, F. J. Liebigs Ann. Chem. 1963, 661, 1–37. 4. Brooker, L. G. S.; Keyes, G. H.; Heseltine, D. N. J. Am. Chem. Soc. 1951, 73, 5350–5356. 5. Johnson, D. A.; Shaw, R.; Silversmith, E. F. J. Chem. Educ. 1994, 71, 517. 6. Rezende, M. C.; Radetski, C. M. Quím. Nova 1988, 11, 353– 354 (Chem. Abst. 1989, 111, 8876 w). Kessler, M. A.; Wolfbeis, O. S. Synthesis 1988, 635–636. 7. Osterby, B. R.; McKelvey, R. D. J. Chem. Educ. 1996, 73, 260–261. 8. Minch, M. J.; Shah, S. S. J. Chem. Educ. 1977, 54, 709. 9. Reichardt, C. In Organic Liquids: Structure, Dynamics, and Chemical Properties; Buckingham, A. D.; Lippert, E.; Bratos, S., Eds.; Wiley: Belfast, 1978; pp 269–281. 10. Novaki, L. P.; El Seoud, O. A. Ber. Bunsenges. Phys. Chem. 1997, 101, 902–909. Machado, V. G.; Machado, C.; Nascimento, M. G.; Rezende, M. C. J. Phys. Org. Chem. 1997, 10, 731–736. 11. Gageiro, V.; Aillon, M.; Rezende, M. C. J. Chem. Soc. Faraday Trans. 1992, 88, 201–204. Reichardt, C.; Asharin-Fard, S.; Schäfer, G. Chem. Ber. 1993, 126, 143–147. 12. Zanotto, S. P.; Scremin, M.; Machado, C.; Rezende, M. C. J. Phys. Org. Chem. 1993, 6, 637–641. 13. Soukup, R. W.; Schmid, R. J. Chem. Educ. 1985, 62, 459462. Scremin, M.; Zanotto, S. P.; Machado, V. G.; Rezende, M. C. J. Chem. Soc., Faraday Trans. 1994, 90, 865–868. 14. Johnson, B. P.; Gabrielson, B.; Matulenko, M.; Dorsey, J. G.; Reichardt, C. Anal. Lett. 1986, 19, 939–962. Langhals, H. Anal. Lett. 1990, 23, 2243–2258. 15. Kanski, R.; Murray, C. J. Tetrahedron Lett. 1993, 34, 2263–2266.
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