Micellar Effects on the Spontaneous Hydrolysis of Phenyl


Bruce H. Lipshutz , Zarko Bošković , Christopher S. Crowe , Victoria K. Davis , Hannah C. Whittemore , David A. Vosburg , and Anna G. Wenzel. Journa...
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In the Laboratory

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Micellar Effects on the Spontaneous Hydrolysis of Phenyl Chloroformate A Kinetic Experiment for the Physical Chemistry Laboratory

Juan Crugeiras,* J. Ramón Leis, and Ana Ríos Departamento de Química Física, Facultad de Química, Universidad de Santiago de Compostela, 15706 Santiago de Compostela, Spain; *[email protected]

The structure and properties of micelles and other association colloids have been extensively investigated owing in part to their importance as simple models of biological membranes. These molecular assemblies are known to alter the rates of chemical reactions, and a large number of papers have been published in recent years describing the ability of aqueous micelles to control reaction rates (1, 2). In spite of the growing interest in microheterogeneous systems, relatively few kinetic experiments have been designed to help students understand the principles of chemical reactivity in membrane mimetic systems (3–10). The laboratory experiment described in this paper was developed to familiarize students with the chemistry of reactions occurring in the microenvironment provided by a particular type of molecular aggregate, the micelle. The spontaneous hydrolysis of phenyl chloroformate offers numerous advantages over other reactions for this purpose. Its mechanism in aqueous solution has been investigated in detail (11–13), the reaction can be monitored by methods easily accessible to undergraduate students (14, 15), and total reaction times are convenient. In addition, micellar effects on unimolecular reactions are the simplest to analyze quantitatively and provide information on the nature of the micelle as a reaction medium. Background Ionic surfactants are amphiphilic molecules that contain ionic head groups and long hydrocarbon tails. In aqueous solution above a certain concentration, known as the critical micelle concentration (cmc), they form roughly spherical aggregates of 50–100 molecules called micelles. These organized structures have apolar interiors formed by hydrophobic tails and ionic surfaces consisting of surfactant headgroups, a fraction of the counterions, and solvating water molecules. Many organic reactions involving relatively hydrophobic substrates occur at the micellar surface (also known as the Stern layer). Therefore, the nature of this micelle–water interface is of key importance because it largely determines the observed micellar effects on reaction rates. The dielectric constant of this region has been estimated to be ~35, showing that its polarity is significantly lower than that of water and comparable to that of ethanol (16 ). Also, the local concentration of charged groups in the Stern layer lies in the range of 3 to 5 M (2). The Pseudophase Model The quantitative rate analysis of chemical reactions in micelles does generally involve the use of the well-known pseudophase model proposed by Menger and Portnoy in 1967 (17 ). 1538

This model assumes that, although micellar solutions are macroscopically homogeneous, the micelles can be considered as a separate phase from water. The reaction may take place in both the aqueous and micellar pseudophases and the overall reaction rate is the sum of the rates in each pseudophase. The application of this model to unimolecular reactions is very simple, since the observed micellar effects can be accounted for in terms of the distribution of only one substrate, S, between water and the micelles (Scheme I). nD

Dn

Sw + Dn

Ks

kw

Sm km

products Scheme I

In Scheme I, Dn is the micelle, KS is the binding constant of the substrate to the micelle, and kw and km are the firstorder rate constants for reaction of the substrate in the water and micellar pseudophases, respectively. Assuming that the concentration of monomers remains constant above the cmc, the concentration of micelles, [Dn], is given by eq 1, where [D] is the total concentration of surfactant and N is the micelle aggregation number.

Dn =

D – cmc N

(1)

The rate law derived for this scheme is shown in eq 2 and it leads to an expression for the observed first-order rate constant, eq 3, which is analogous to the Michaelis–Menten equation for enzyme-catalyzed reactions. rate = kw[Sw ] + km[Sm] = kobsd[ST]

k obsd =

k w + k mK S Dn 1 + K S Dn

(2) (3)

The value of k w can be easily evaluated from kinetic measurements in the absence of surfactant. Equation 3 rearranges to eq 4, which allows the calculation of km and KS by plotting 1/(kw – kobsd) versus 1/[Dn].

1 1 1 = + k w – k obsd k w – k m k w – k m KS Dn

(4)

This equation gives a good description of the kinetics of both micellar inhibited and catalyzed unimolecular reactions.

Journal of Chemical Education • Vol. 78 No. 11 November 2001 • JChemEd.chem.wisc.edu

In the Laboratory

Results and Discussion The spontaneous hydrolysis of phenyl chloroformate is a well-known reaction (13) that involves the addition of a water molecule to the carbonyl group in the slow step (Scheme II). OCOCl

H 2O slow

OCOOH + HCl

OH + CO2

OCOOH

Scheme II

The experiment described in this paper involves a study of this reaction in the presence of anionic micelles of sodium dodecyl sulfate (SDS). The reaction can be easily followed by monitoring the formation of phenol in the UV region. However, very often student labs are equipped with simple spectrophotometers, such as the Spectronic 20 (Milton Roy), without a UV source. Alternatively, the first-order rate constant can be obtained by recording the increase in electrical conductivity due to the formation of HCl. Figure 1 shows typical values of kobsd as a function of the total surfactant concentration. Below the cmc, the surfactant monomers do not have a significant effect on the observed rate constant, whereas above the cmc the presence of SDS micelles inhibits the reaction. These results are qualitatively explained assuming that the substrate is taken up by the micelles and that its hydrolysis at the micellar surface is slower than in water. A quantitative analysis of the observed micellar effects can be made in terms of the pseudophase model according to Scheme I. A value of the cmc can be determined as the concentration of surfactant corresponding to the break point in the kobsd versus [SDS] plot (Fig. 1, inset). Additionally, if the rate constants are obtained by conductivity measurements, the cmc of SDS can be calculated from the change in the

Figure 1. Influence of [SDS] on the observed first-order rate constant for spontaneous hydrolysis of phenyl chloroformate at 25 °C. Inset: Plot of kobsd vs [SDS] used to determine the cmc of the surfactant.

electrical conductivity of the SDS solutions with increasing surfactant concentration (18). The values obtained using both methods are in the range (8.2 ± 0.4) × 10᎑3 M and are consistent with published values (18–20). The concentration of micelles, [Dn], was calculated using cmc = 8 × 10᎑3 M and N = 60 (21). The hydrolysis rate constant in water, kw, equals the value of kobsd in the absence of surfactant (1.31 × 10᎑2 s᎑1). Figure 2 shows the good fit of the experimental data to eq 4. Values of the substrate association constant, KS = (4.6 ± 0.1) × 103 M᎑1, and the rate constant for hydrolysis inside the micelle, km = (9 ± 1) × 10᎑4 s᎑1, are easily calculated from the intercept and the slope of the linear plot. The solvolysis of phenyl chloroformate is believed to follow an addition–elimination pathway, the addition step being rate-determining (13). The observed micellar effect for this reaction is attributed to the lower polarity of the micellar interface and to the electrostatic destabilization of the transition state (see structure) by the negatively charged headgroups of the surfactant (22). δO O

C

Cl

O H

H

H O δ+ H

The laboratory experiment could easily be organized so that each group of students carries out the same study at a different temperature. After all the students have analyzed their rate data using the pseudophase model, the values of kw and km can be put together to obtain the activation parameters for the hydrolysis of phenyl chloroformate in both the aqueous and micellar pseudophases according to the Eyring equation.

Figure 2. Linear plot of 1/(kw – kobsd) vs 1/[Dn] used to calculate KS and km according to eq 4.

JChemEd.chem.wisc.edu • Vol. 78 No. 11 November 2001 • Journal of Chemical Education

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In the Laboratory

Conclusions This experiment is suitable for an advanced laboratory in physical chemistry for the following reasons. First, it exposes students to the common laboratory techniques used to obtain kinetic data. Most students are familiar with simple kinetics and are able to correctly calculate first-order rate constants from the original data. Second, in addition to introducing the main features of aqueous micelles and the factors determining micellar effects on chemical reactivity, this experiment illustrates some of the methods commonly used to determine the cmc of surfactants. Third, it provides numerous learning opportunities for students by clearly showing how aqueous micelles can alter the rates of chemical reactions and how the observed micellar effects can be quantitatively analyzed using the pseudophase model.

3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13.

Hazards

14.

Phenyl chloroformate, acetonitrile and sodium dodecyl sulfate are harmful if ingested or inhaled or come in contact with the skin. Protective clothing, appropriate gloves, and safety glasses should be worn when handling these materials.

15. 16. 17.

WSupplemental

18.

Material

Notes for the instructor and detailed instructions for students are available in this issue of JCE Online. Literature Cited 1. Bunton, C. A.; Savelli, G. Adv. Phys. Org. Chem. 1986, 22, 231–309 and references cited therein. 2. Bunton, C. A.; Nome, F.; Quina, F. H.; Romsted, L. S. Acc.

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19. 20. 21. 22.

Chem. Res. 1991, 24, 357–364 and references cited therein. Corsaro, G. J. Chem. Educ. 1973, 50, 575–576. Corsaro, G.; Smith, J. K. J. Chem. Educ. 1976, 53, 589–590. Corsaro, G. J. Chem. Educ. 1980, 57, 225–226. Garcia-Mateos, I.; Herraez, M. A.; Rodrigo, M.; Rodriguez, L. J.; Velazquez, M. M. J. Chem. Educ. 1981, 58, 584–585. Marzzacco, C. J. J. Chem. Educ. 1992, 69, 1024–1025. Reinsborough, V. C.; Robinson, B. H. J. Chem. Educ. 1981, 58, 586–588. Rodenas, E.; Vera, S. J. Chem. Educ. 1985, 62, 1120–1121. Williams, K. R. J. Chem. Educ. 2000, 77, 626–628. Queen, A. Can. J. Chem. 1967, 45, 1619–1629. Butler, A.; Roberson, I. H. J. Chem. Soc., Perkin Trans. 2 1974, 1733–1736. Kevill, D. N.; D’Souza, M. J. J. Chem. Soc., Perkin Trans. 2 1997, 1721. Seoud, O. A. E.; Seoud, M. I. E.; Pires, P. A. R.; Takashima, K. Educ. Chem. 1997, 34, 22–23. Seoud, O. A. E.; Takashima, K. J. Chem. Educ. 1998, 75, 1625–1627. Cordes, E. H. Pure Appl. Chem. 1978, 50, 617–625. Menger, F. M.; Portnoy, C. E. J. Am. Chem. Soc. 1967, 89, 4698–4703. Dominguez, A.; Fernandez, A.; Gonzalez, N.; Iglesias, E.; Montenegro, L. J. Chem. Educ. 1997, 74, 1227–1231. Goodling, K.; Johnson, K.; Lefkowitz, L.; Williams, B. W. J. Chem. Educ. 1994, 71, A8–A12. Stam, J. v.; Depaemelaere, S.; Schryver, F. C. D. J. Chem. Educ. 1998, 75, 93–98. Turro, N. J.; Yekta, A. J. Am. Chem. Soc. 1978, 100, 5951–5952. Al-Lohedan, H.; Bunton, C. A.; Mhala, M. M. J. Am. Chem. Soc. 1982, 104, 6654–6660.

Journal of Chemical Education • Vol. 78 No. 11 November 2001 • JChemEd.chem.wisc.edu