In the Laboratory
Micelles in the Physical Chemistry Laboratory. Kinetics of Hydrolysis of 5,5'-Dithiobis-(2-nitrobenzoate) in Basic Solution
W
Kathryn R. Williams Department of Chemistry, University of Florida, Gainesville, FL 32611-7200;
[email protected] Amphiphilic compounds, commonly known as surfactants, have distinct polar and nonpolar regions. In aqueous solution, these molecules form aggregates, called micelles, when the amount exceeds the critical micelle concentration (cmc). Micelles are important in a wide variety of applications, and it is therefore important for students to have some appreciation of the special properties of these systems. A previous publication described a physical chemistry experiment on the diffusion coefficient of ferrocene in water and in three different micellar systems (1). The experiment presented here demonstrates the effects of micelles on the kinetics of a hydrolysis reaction. Background Incorporation of micelles into a reactant mixture can produce striking changes in the reaction rate, and both acceleration and inhibition effects have been observed (2). Rate enhancement is often called micellar catalysis; and indeed, in some cases, the rate increase is due to electrostatic effects, which destabilize the reactant and/or stabilize the transition state (3–6 ). However, these factors cannot account for the large changes in rates of second- and higher-order reactions (3, 7 ). The generally accepted explanation is based on the localization of the reacting species in or near the surface of the micelles, which have a relatively small volume compared to the bulk solution (3). This leads to a large increase in the effective concentration, and the observed rate (in terms of moles per unit time per liter of the entire solution) increases accordingly. The mathematical theories of micellar catalysis were proposed in the 1970s. The work of Berezin et al. (8) focused on the partitioning of the reactants between the aqueous and micellar phases. This approach worked well for organic substrates, but was inadequate for reactions involving small ions. Romsted (7) and later Quina and Chaimovich (9) proposed an ion exchange model, in which the small reacting ions must compete for surface sites with the surfactant counterions. The rates of many second-order reactions, including the hydrolysis described below, pass through a maximum at a certain micellar concentration (well above the cmc) (3). According to the partitioning model, there is a reversal in the local concentration effect. Between the cmc and the concentration at the maximum rate, the effective concentrations of the reactants increase owing to localization in the micellar phase. Eventually, the fractional volume due to the micelles becomes large enough to cause the effective concentrations and the observed rate to decrease. The ion exchange model takes into consideration that each increment of surfactant is accompanied by an equal number of counterions, which compete with the reacting ions for the surface sites. This eventually results in a decrease in the effective concentration of the reacting ions. 626
Hydrolysis Reaction The literature contains many examples of kinetics studies involving micelles, including several reports of student experiments and demonstrations (4–6, 10–13). The investigation presented here is a study of the hydrolysis of 5,5′dithiobis-(2-nitrobenzoate) (DTNB2᎑; see Fig. 1 for structures of all compounds) in basic solution both without and with positively charged micelles of cetyltrimethylammonium bromide (C16TAB). The DTNB2᎑ hydrolysis was studied in detail by Fendler and Hinze (14 ), whose results supported a mechanism involving initial rate-determining attack of hydroxide ion on the sulfur linkage to produce 5-thio-2-nitrobenzoate (TNB2᎑) plus 5-sulfeno-2-nitrobenzoate (SENB᎑): DTNB2᎑ + OH᎑ → TNB2᎑ + SENB᎑ (slow)
(1)
The sulfeno compound rapidly disproportionates to give another TNB2᎑ and 5-sulfinyl-2-nitrobenzoate (SINB2᎑): SENB᎑ + OH᎑ → 1⁄2TNB2᎑ + 1⁄2SINB2᎑ + H2O
(2)
2᎑
The overall reaction of one mole of DTNB consumes 2 mol of OH᎑ and yields 3⁄2 mol of TNB2᎑ plus 1⁄2 mol of SINB2᎑: DTNB2᎑ + 2OH᎑ → 3⁄2TNB2᎑ + 1⁄2SINB2᎑ + H2O (3) The reaction follows an overall second-order rate law, Rate = k2[OH᎑][DTNB2᎑]
(4)
but, owing to the large excess of hydroxide (0.01–0.03 M) over DTNB2᎑ (ca. 5 × 10᎑5 M), pseudo-first-order kinetics are observed: Rate = kOH[DTNB2᎑]
Figure 1. Chemical structures and abbreviations.
Journal of Chemical Education • Vol. 77 No. 5 May 2000 • JChemEd.chem.wisc.edu
(5)
In the Laboratory
The hydrolysis experiment fulfills several instructional goals. Students become very familiar with first-order plots and second-order rate constants. Also, the experimental conditions specified in the student instructions clearly demonstrate the effect of C16TAB micelles on the rate. The compound is readily available from commercial sources, and the reaction rate is easily monitored by measuring either the decrease in the DTNB2᎑ absorbance or the corresponding increase in the absorbance of the TNB2᎑ product. With the conditions specified in the student instructions, the reactions go to completion in the time range of 15 seconds (ca. 0.02 M NaOH and 2 × 10᎑3 M C16TAB) to 15 minutes (ca. 0.01 M NaOH without C16TAB), so that students can obtain several data sets in a typical laboratory period. The physical chemistry laboratory at the University of Florida is equipped with two Spectral Instruments SI 440 fiber optic spectrophotometers, which are especially useful, because the reaction mixture can be contained in a simple vessel surrounded by thermostated water. However, with suitable adjustment of reaction volumes, the experiment can easily be conducted using a conventional UV instrument with 1-cm cells.
Table 1. Typical Kinetics Results for DTNB2 - Hydrolysis Without C1 6 TAB; [DTNB2 ᎑] = 5.07 × 10᎑5 M ᎑
[OH ]/mol L᎑1
kO H ,3 2 4 n m /s᎑1
kO H ,4 1 0 n m /s᎑1
0.01046
0.00376
0.00375
0.01568
0.00618
0.00593
0.02091
0.00848
0.00810
0.01354
0.01327
0.03137
k2 = (0.462 ± 0.009) M᎑1 s᎑1 With C1 6 TAB; [DTNB2 ᎑] = 5.07 × 10᎑5 M; [OH᎑] = 0.02091 M [C1 6 TAB]/mol L᎑1
kO H ,3 1 2 n m /s᎑1
kO H ,4 3 5 n m /s᎑1
5.42 × 10᎑4
0.121
0.118
2.17 × 10᎑3
0.155
0.156
᎑3
4.34 × 10
0.152
0.147
8.68 × 10᎑3
0.110
0.107
1
0
-1
-2
-3
-4 0
5
10
15
20
25
30
Time / s Figure 2. Plots of ln(A∞ – A) vs time for TNB2᎑ in 0.02091 M NaOH. Squares: 5.42 × 10 ᎑4 M C16TAB, slope = ᎑0.1176 s ᎑1. Circles: No C16TAB, slope = ᎑0.00810 s᎑1, data linear over a total time of 400 s. 0.011 0.010 0.009 0.008
k(s - 1)
Solutions of known [OH᎑] (and [C16TAB] for pertinent runs) are prepared by mixing suitable volumes of water and 0.2 M NaOH (and 0.02 M C16TAB) to give a total volume of 20 mL in the reaction vessel. After temperature equilibration at 25 °C and spectrophotometer setup, the reaction is initiated by rapidly adding 100 µL DTNB2᎑ from a pipetting device. (The DTNB2᎑, provided as a ca. 10᎑2 M solution in ethanol, is stable for several weeks.) Data are collected at the λmax values for both DTNB2᎑ (324 nm in H2O or 312 nm in C16TAB) and TNB2᎑ (410 nm in water or 435 nm in C16TAB), until the DTNB2᎑ has completely reacted.1 Table 1 and Figures 2 and 3 present typical results. As shown in Figure 2, students verify that the reaction is first order with respect to DTNB2᎑. Values of kOH are obtained from the slopes of the first-order plots for a series of reactions using nominal [OH᎑] values of 0.010, 0.015, 0.020, and 0.030 M in water. The second-order rate constant, k2, is obtained from the slope of a kOH versus [OH᎑] plot, as shown in Figure 3. Plots of log kOH versus log [OH᎑] are also very linear, with slopes in reasonable agreement with the theoretical value of unity (1.15 ± 0.02 for the data in Table 1). The effect of added micelles is investigated using several C16TAB concentrations at a constant [OH᎑] of ca. 0.020 M. Referring to Figure 2 and Table 1, the rate enhancement in the micellar environment is obvious. By using surfactant concentrations spanning the range of ca. 5 × 10᎑4 to 9 × 10᎑3 M, students can also observe that the maximum in rate enhancement occurs in the range of 2 × 10᎑3 to 4 × 10᎑3 M C16TAB, in reasonable agreement with the results of Fendler and Hinze (ca. 1 × 10᎑3 M C16TAB in 0.0214 M OH᎑ at 26.4 °C [14 ]). As part of the experiment, students obtain full spectra of DTNB2᎑ (at pH 8.0, where hydrolysis is very slow) and TNB2᎑ in both water and C16TAB solution. The absorbances are used with the molar absorptivity data provided by Fendler and Hinze (14 ) to verify that the overall stoichiometry given by eq 3 is indeed observed.
Ln(Ainf -A)
Overview of the Procedure and Data AnalysisW
0.007 0.006 0.005 0.004 0.003 0.005
0.010
0.015
0.020
0.025
0.030
[NaOH] / (mol/L) Figure 3. Plot of kOH vs [OH᎑] for DTNB2᎑ hydrolysis without C16TAB. The two points for each [OH᎑] represent data using the disappearance of DTNB2᎑ and the appearance of TNB2᎑.
JChemEd.chem.wisc.edu • Vol. 77 No. 5 May 2000 • Journal of Chemical Education
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In the Laboratory
Safety Considerations Students should follow customary safe laboratory procedures. Considering the low concentrations involved, the reaction solutions may be flushed down the drain. Excess DTNB2᎑ stock solutions should be saved for incineration. Conclusion The hydrolysis of DTNB2᎑ is an excellent choice for a student kinetics experiment. The reagents and instrumentation are readily available, and the reaction times are convenient. Studies in both water and surfactant solutions allow students to observe the catalytic effect of C16TAB micelles. Acknowledgment I thank The Camille and Henry Dreyfus Foundation, Inc., for providing funds to purchase the fiber optic spectrophotometers. W
Supplemental Material
Supplemental material for this article is available in this issue of JCE Online. The laboratory procedure and data analysis are explained in detail in the student instructions. Note 1. Fendler and Hinze (14 ) comment on the opposite wavelength changes produced by C16TAB (blue shift for DTNB 2᎑;
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red shift for TNB2᎑). Interested faculty may want students to consider this behavior in their reports.
Literature Cited 1. Williams, K. R.; Bravo, R. J. Chem. Educ. 2000, 77, 392– 394. 2. Fendler, J. H.; Fendler, E. J. Catalysis in Micellar and Macromolecular Systems; Academic: New York, 1975. 3. Cordes, E. H. Pure Appl. Chem. 1978, 50, 617–625. 4. Corsaro, G. J. Chem. Educ. 1973, 50, 575–576. 5. Corsaro, G. J. Chem. Educ. 1976, 53, 589–590. 6. Corsaro, G. J. Chem. Educ. 1980, 57, 225–226. 7. Romsted, L. S. In Micellization, Solubilization, and Microemulsions; Mittal, K. L., Ed.; Plenum: New York, 1976; Vol. 2, pp 509–530. 8. Martinek, K.; Yatsimirski, A. K.; Levashov, A. V.; Berezin, I. V. In Micellization, Solubilization, and Microemulsions; Op. cit.; pp 489–508. 9. Quina, F. H.; Chaimovich, H. J. Phys. Chem. 1979, 83, 1844– 1850. 10. García-Mateos, I.; Herráez, M. A.; Rodrigo, M.; Rodríguez, L. J.; Velázguez, M. M. J. Chem. Educ. 1981, 58, 584–585. 11. Reinsborough, V. C.; Robinson, B. H. J. Chem. Educ. 1981, 58, 586–588. 12. Rodenas, E.; Vera, S. J. Chem. Educ. 1985, 62, 1120–1121. 13. Marzzacco, C. J. J. Chem. Educ. 1996, 73, 254–255. 14. Fendler, J. H.; Hinze, W. L. J. Am. Chem. Soc. 1981, 103, 5439–5447.
Journal of Chemical Education • Vol. 77 No. 5 May 2000 • JChemEd.chem.wisc.edu
