Calorimetric and NMR Investigations of the Micellar Properties of

Langmuir 1999, 15, 429-436

429

Calorimetric and NMR Investigations of the Micellar Properties of Sodium Dodecyl Sulfate in Aqueous Mixtures of Isomeric Butanediols C. A. McMahon, Brent Hawrylak, D. Gerrard Marangoni,* and R. Palepu* Department of Chemistry, St Francis Xavier University, P.O. Box 5000, Antigonish, Nova Scotia, B2G 2W5, Canada Received October 7, 1997. In Final Form: October 26, 1998

The enthalpies (∆micH°), entropies (∆micS°), and Gibbs energies (∆micG°) of micelle formation have been determined for a series of sodium dodecyl sulfate (SDS)/butanediol isomer micelles using isoperibol solution calorimetry. Nuclear magnetic resonance (NMR) paramagnetic relaxation enhancement experiments were used to determine the degree of solubilization of the isomeric butanediols (BTDs) in a 0.100 M solution of SDS micelles. In addition, the manner in which the butanediol cosolvents interact with SDS micelles, as a function of the position of substitution of the hydroxyl groups on the four-carbon backbone, has been examined closely via 13C NMR spectroscopy. All of these results are interpreted in terms of the differences in the interaction of the isomeric butanediols with SDS micelles because of the addition of the second hydroxyl headgroup on the four-carbon backbone of the alcohol, when compared to the well-studied SDS/ 1-butanol (C4OH) mixed micellar system.

Introduction Over the past number of years, there has been a growing interest in the micellar properties of ionic surfactants dissolved in aqueous/nonaqueous solvent mixtures.1-7 It is well-known that the presence of nonelectrolytes (e.g., n-alcohols) in aqueous solution has a significant impact on the properties of the micelles formed by ionic surfactants in the solvent mixture, when compared to the micelles formed by the surfactant in water. In fact, as the percentage of the nonelectrolyte in solution is increased, micelle formation has still been reported for ionic surfactants in mixed-solvent systems such as ethylene glycol, hydrazine, formamide, n-alcohol, ethoxylated alcohols, and diethylene glycol.2-7 The presence of the organic cosolvent may alter dramatically the manner in which the surfactant self-assembles to form micelles. This, in turn, may lead to significant changes in the properties of the micelles formed by ionic surfactants (i.e., differences in the critical micelle concentrations, or cmc values, aggregation numbers, and degrees of counterion binding) that are a function of the amount of cosolvent in the mixed-solvent system. The thermodynamic properties of micelle formation (the Gibbs energy, enthalpy, and entropy of micelle formation) are sensitive parameters for discussing the interactions that lead to micelle formation in aqueous and nonaqueous solvent systems. There have been a few literature reports of the thermodynamic parameters of micelle formation in water/nonelectrolyte mixtures and in nonelectrolyte * To whom correspondence should be addressed. Phone: (902) 867-3886 (R.P.) and (902) 867-2324 (D.G.M.). Fax: (902) 867-2414. E-mail: [email protected] and [email protected]. (1) Onori, G.; Santucci, A. Chem. Phys. Lett. 1992, 189, 598. (2) Bakshi, M. S. Bull. Chem. Soc. Jpn. 1996, 69, 2723. (3) Callaghan, A.; Doyle, R.; Alexander, E.; Palepu, R. Langmuir 1993, 9, 3422. (4) Misra, P. K.; Mishra, B. K.; Behera, G. B. Colloids Surf. 1991, 57, 1. (5) Gracie, K.; Turner, D.; Palepu, R. Can. J. Chem. 1996, 74, 1616. (6) Beesley, A.; Evans, D. F.; Laughlin, R. G. J. Phys. Chem. 1988, 92, 791. (7) Evans, D. F.; Kaler, E. W.; Benton, W. J. J. Phys. Chem. 1987, 91, 5944.

solvents;1-5 however, the thermodynamic properties of micelle formation in these systems have mostly been determined via the use of the temperature derivatives of the ln(Xcmc) values. There has been a significant amount of work in the literature that suggests that the use of the temperature derivatives of the ln Xcmc values leads to ∆micH° values (and hence ∆micS° values) that are significantly different from those of direct measurements via calorimetry.8-10 Despite this fact, there are very few direct (calorimetric) measurements of the thermodynamic properties of micelle formation of ionic surfactants in mixedsolvent systems, such as the n-alcohols in water; hence, very little is known about how the presence of the organic cosolvent influences the thermodynamic properties of micelle formation.8-13 In water/nonaqueous solvent systems, a significant factor in the determination of the fundamental micellar properties (i.e., cmc values, thermodynamic parameters, aggregation numbers, degrees of counterion binding), as a function of increasing concentration of nonelectrolyte in the mixed solvent, is the degree to which the nonelectrolyte penetrates into the micellar phase.14-21 The mole fraction (8) Van Os, N. M.; Daane, G. J.; Handrikkman, G. J. Colloid Interface Sci. 1991, 141, 109. (9) White, P.; Benson, G. C. Trans. Faraday Soc. 1959, 54, 1025. (10) Burrows, J. C.; Flynn, D. J.; Leriche, T. G.; Kutay, S. M.; Marangoni, D. G. Langmuir 1995, 11, 3388. (11) Mukherjee, K.; Mukherjee, D. C.; Moulik, S. P. J. Phys. Chem. 1994, 98, 4713. (12) Mukherjee, K.; Mukherjee, D. C.; Moulik, S. P. Langmuir 1993, 9, 1727. (13) Johnson, I.; Olofsson, G.; Jo¨nnson, B. J. Chem. Soc., Faraday Trans. 1 1987, 83, 3331. (14) Marangoni, D. G.; Kwak, J. C. T. In Surfactant Science Series: Solubilization; Christian, S. A., Scamehorn, J. F., Eds.; Marcel Dekker: New York, 1995; Chapter 14. (15) Stilbs, P. J. Colloid Interface Sci. 1982, 87, 385. (16) Stilbs, P. J. Colloid Interface Sci. 1982, 89, 547. (17) Carlfors, J.; Stilbs, P. J. Colloid Interface Sci. 1985, 104, 489. (18) Stilbs, P. Prog. NMR Spectrosc. 1987, 19, 1. (19) So¨derman, O.; Olsson, U. Curr. Opin. Colloid Interface Sci. 1997, 2, 131. (20) Gao, Z.; Wasylishen, R. E.; Kwak, J. C. T. J. Phys. Chem. 1989, 93, 2190.

10.1021/la971093+ CCC: $18.00 © 1999 American Chemical Society Published on Web 12/30/1998

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of solubilizate in the micellar phase, p, is defined as

p ) cs,mic/cs,t

(1)

Here, cs,mic is the concentration of solubilizate in the micellar phase and cs,t is the total concentration of solubilizate. The p values of solubilizates have been measured via a number of different techniques.14 NMR determinations of the p values of solubilizates are accomplished via the application of the Fourier transform pulsed-gradient spin-echo (FT-PGSE) experiment by Stilbs and co-workers15-19 or by the nuclear magnetic resonance (NMR) paramagnetic relaxation enhancement experiment (NMR-PRE).20,21 This method is based on the changes in the relaxation rate of the solubilizate in surfactant and surfactant-free solutions, in the presence of paramagnetic ions, and has been successfully applied in a number of surfactant/solubilizate systems.20-22 13C NMR spectroscopy can provide semiquantitative information on conformational changes that occur upon micelle formation.23-27 In the case of mixed micellar systems, the 13C chemical shifts can provide information on the packing changes that occur when a second component (a cosurfactant) is dissolved in the micellar interior. Jansson et al.27 examined the chemical shift changes incurred by the zwitterionic surfactant (decyldimethylammonio)propanesulfonate (DAPS) upon micellar inclusion of 1-hexanol, 1-octanol, and 1-decanol. Their results indicated that the 13C chemical shift changes for the surfactant upon the incorporation of alcohols into DAPS micelles were dependent on the chain length of the alcohol and, hence, the degree of solubilization of alcohol (the p value) into DAPS micelles. These chemical shift changes were explained in terms of the conformational changes incurred by both the alcohol and the surfactant upon transfer from the aqueous phase to the micellar phase. There has been a significant amount of work in the literature dealing with the properties of n-alcohol with anionic and cationic micelles. However, little research has been done on the interaction of other families of alcohols (i.e., alkanediols) with anionic and cationic micelles.28-32 The presence of the second, hydrophilic headgroup has been shown to have a significant impact on the interaction of the solubilizate with the micellar interior, when compared with their monohydroxy counterparts.28-31 The position of the second headgroup on the hydrophobic chain may influence the interaction of the diol with the micelles. In this paper, we report our results on the measurement of the thermodynamic properties of micelle formation of sodium dodecyl sulfate (SDS) in a series of butanediol/ water mixtures. Isoperibol titration calorimetry is used to obtain estimates of the micellization enthalpy of SDS (21) Gao, Z.; Marangoni, D. G.; Labonte´, R.; Wasylishen, R. E.; Kwak, J. C. T. Colloids Surf. 1990, 45, 269. (22) Yoshino, A.; Murate, K.; Yoshida, T.; Okabayashi, H.; Krishna, P. R.; Kamaya, H.; Ueda, I. J. Colloid Interface Sci. 1994, 166, 375. (23) Brycki, B.; Szafran, M. Magn. Reson. Chem. 1992, 30, 535. (24) Perrson, B. D.; Drakenberg, T.; Lindman, B. J. Phys. Chem. 1979, 83, 3011. (25) Mazumdar, S. J. Phys. Chem. 1990, 94, 5947. (26) Chevalier, Y.; Chachaty, C. Colloid Polym. Sci. 1984, 262, 489. (27) Jannson, M.; Li, P.; Stilbs, P. J. Phys. Chem. 1987, 91, 5279. (28) Blokhus, A. M.; Høiland, H.; Backlund, S. J. Colloid Interface Sci. 1986, 114, 9. (29) Høiland, H.; Blokhus, A. M. In The Structure, Dynamics, and Equilibrium Properties of Colloidal Systems; Bloor, D. M., Wyn-Jones, E., Eds.; Kluwer Academic Publishers: Boston 1990. (30) Can˜adas, O.; Valiente, M.; Rodenas, E. J. Colloid Interface Sci. 1998, 203, 294. (31) Palepu, R., unpublished results. (32) So¨derman, O.; Geuring, P. Colloid Polym. Sci. 1987, 265, 76.

Figure 1. 13C NMR spectrum and peak assignments for the 0.100 M SDS/1% 1-butanol (C4OH) system.

in the mixed-solvent systems as a function of the position of the hydroxyl groups on the C4 backbone, the composition of the mixed solvent, and the temperature. The fraction of the butanediol solubilizate dissolved in the mixed micelle has been determined via the NMR-PRE experiment. Finally, we have determined the 13C chemical shift changes that occurred when SDS is dissolved in the isomeric butanediol/water and C4OH/water mixtures. All of these results will be interpreted in terms of the difference in the interaction of the two-headed alcohols (the isomeric butanediols) with SDS micelles compared to the wellstudied SDS/1-butanol mixed micellar system. Experimental Section Solution Preparation and Calorimetry. SDS was purchased from BDH (specially pure). The surfactant was purified by repeated recrystallizations from methanol and Soxhlet extraction from diethyl ether. The isomeric butanediols and the 1-butanol (purity >99%) were obtained from Aldrich and were used as received. Conductivity grade water (1.0 × 10-6 S cm-1) was obtained by passing previously deionized water through a Millipore 4000 S purification system. The alcohol/water mixtures were prepared on a mass percent basis. The SDS concentrations are reported as the moles of surfactant per liter of the mixed solvent. Densities of the titrant solutions were measured using an Anton-Parr DMA 45 densitometer. The heat capacity data for the titrant solutions and the enthalpies of micellization were obtained using a Hart model 4285 isoperibol solution calorimeter. The operation of the calorimeter and the measurement of the heat capacity data were described previously.10 NMR Experiments. The NMR spectra (1H and protondecoupled 13C) for the SDS/mixed-solvent systems were obtained on a Bruker AC-200 operating at 200.13 MHz for protons and 50.32 MHz for 13C. For the 13C spectra, a pulse length of 5 µs, corresponding to a flip angle of ca. 60°, was used. A sweep width of 12 500 Hz and 32K of data points were used. The number of acquisitions was usually 1024. A typical 13C spectrum for the system SDS/1.0% 1-butanol is given in Figure 1. D2O was used as the solvent for all of the NMR experiments. The 1H spectra were referenced to the HOD peak (δ ) 4.81 ppm) while the 13C chemical shifts were referenced to the deuterium lock signal using the method of So¨derman.32 All spin-lattice relaxation times (T1 values) were measured on freshly prepared solutions. A solution of the paramagnetic spin probe, the sodium salt of 3-carboxyproxyl, was prepared in D2O (99.9% Aldrich). SDS solutions of 100 mM were prepared directly in the NMR tubes or in small glass vials with either D2O or the paramagnetic ion solution as the solvent. The isomeric butanediols (BTDs) were added to the NMR tubes or the glass vials with a Hamilton syringe. The T1 measurements were done

Micellar Properties of Sodium Dodecyl Sulfate

Langmuir, Vol. 15, No. 2, 1999 431 Table 1. cmc Values and Micelle Formation Enthalpies, Entropies, and Gibbs Energies of SDS in 1,2-Butanediol/ Water Mixtures at 298 and 308 Ka 298 K

308 K

cmc cmc wt % value value diol (mM) ∆micH° ∆micG° ∆micS° (mM) ∆micH° ∆micG° ∆micS° 0.00 1.00 2.00 3.00 5.00 10.0

8.10 7.10 6.60 6.20 5.80 5.90

-0.22 0.00 -0.48 -0.73 -4.90 -8.21

-35.9 -36.4 -36.3 -36.5 -36.6 -36.3

120 122 120 120 106 94

8.30 7.70 7.40 6.70 6.20 6.40

-5.10 -4.59 -6.53 -7.28 -9.21 -13.6

-36.6 -36.6 -36.8 -37.2 -37.3 -36.7

102 104 98 97 91 75

a ∆ micH° ( 0.50 kJ/mol. ∆micG° ( 0.50 kJ/mol. ∆micS° ( 3 J/(K mol).

Figure 2. Typical q vs V titration curve for the SDS/1,2butanediol system. on solutions of ≈0.060 M of the diol, because it is well-known that large quantities of additives may make significant perturbations to the micellar structure.14 The spin-lattice relaxation times (T1 values) were obtained using the standard inversionrecovery pulse program in the Bruker software library. The T1’s were calculated from the peak heights obtained from at least 12 variable delays, using a three-parameter, nonlinear least-squares fitting procedure. All T1 measurements were carried out at 24 ( 1 °C.

1. Micellization Enthalpies for SDS Dissolved in Isomeric Butanediol/Water Mixtures. For an alcohol/ surfactant mixed micelle, the ∆micH° value refers to the enthalpy change that occurs when n moles of surfactant S (charge a), m moles of counterion C (charge b), and l moles of additive A aggregate in aqueous solution to form a mixed micelle, Mic, of charge (z:

∆micH°

(2)

The ∆micH° values for SDS/isomeric BTD systems were measured directly via isoperibol solution calorimetry. The isoperibol solution calorimeter measures the heat changes that occur when small quantities of a concentrated surfactant solution are titrated into a fixed quantity of solvent or when a small volume of a concentrated surfactant solution is diluted with the addition of solvent. A typical plot of qrxn vs the volume of added titrant is given in Figure 2. It can be clearly seen from Figure 2 that there are two linear regions above and below the break point (the cmc region). From the slopes of the qrxn vs Vtitrant plots, the enthalpy of micelle formation is calculated as follows.

∆micH° ) L2(above cmc) - L2(below cmc) )

{

dq 1000 dq (above cmc) (before cmc) M dV dV

298 K

308 K


8.1 8.1 7.8 7.4 6.7 6.2

-0.22 0.15 -0.44 -0.51 -2.88 -7.80

-35.9 -35.9 -35.6 -35.8 -36.0 -36.1

120 121 118 118 111 95

8.3 8.1 8.1 7.8 7.4 6.7

-5.10 -4.39 -4.84 -5.80 -6.55 -13.37

-36.6 -36.2 -36.0 -36.4 -36.1 -36.8

102 103 101 99 96 76


Results and Discussion

nSa + mCb + lA h [Mic](z

Table 2. cmc Values and Micelle Formation Enthalpies, Entropies, and Gibbs Energies of SDS in 2,3-Butanediol/ Water Mixtures at 298 and 308 Ka

}

(3)

Here, L2(above/below cmc) represents the partial molar enthalpy of the surfactant in the concentration range just above and below the cmc value, respectively; (dq/dV)(before cmc) is the observed slope of the qrxn vs V curve when the surfactant solution is first diluted into the solvent, (dq/ dV)(above cmc) is the observed slope just above the cmc, and M is the molarity of the stock surfactant solution in the calorimeter buret. Because the cmc may also be determined from the qrxn vs V curves, the Gibbs energy

of micelle formation (∆micG°) may be obtained from the application of the charged pseudophase model of micelle formation to the cmc values.

∆micG° ) (2 - R)RT ln Xcmc

(4)

Here Xcmc is the mole fraction of surfactant at the cmc, and R is defined as the degree of counterion dissociation. Therefore, once ∆micG° and ∆micH° have been obtained from the calorimetric experiment, the micelle formation entropy, ∆micS°, may be found from the following equation:

∆micS° ) (∆micH° - ∆micG°)/T

(5)

Hence, in principle, the Gibbs energy and the enthalpy of micelle formation may also be obtained directly from the calorimetric experiment. It should be noted that the thermodynamic properties of micelle formation might also be estimated by differentiating from the temperature dependence of the ln Xcmc values. However, it has been demonstrated in the literature that the agreement between the enthalpies of micelle formation from direct measurements vs temperature differentiation of the ln Xcmc values is poor.8-10 Hence, we will discuss only the calorimetrically determined micellar thermodynamic properties in the following paragraphs. The results for the calorimetric determination of the ∆micH° of SDS in isomeric butanediol/water mixtures are given in Tables 1-4 and plotted in Figures 3 and 4 at 298 K, respectively. There are a number of trends apparent in Tables 1-4. The first of these is that for the formation of micelles in aqueous butanediol solutions the enthalpy of micellization remains essentially constant at the lower weight percent of the diols and decreases significantly thereafter. This indicates that the interactions between the diols and the surfactants (and each of the components

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Table 3. cmc Values and Micelle Formation Enthalpies, Entropies, and Gibbs Energies of SDS in 1,3-Butanediol/ Water Mixtures at 298 and 308 Ka 298 K

308 K


8.1 8.2 8.3 8.4 8.4 8.6

-0.22 -0.07 -0.50 -0.73 -2.56 -6.07

-35.9 -35.9 -35.6 -35.5 -35.1 -35.2

120 120 118 117 109 98

8.3 8.1 8.3 8.3 8.4 9.0

-5.10 -10.9 -11.3 -12.9 -25.0 -40.0

-36.7 -36.4 -36.4 -36.4 -35.9 -36.3

102 83 91 76 35 -12


Table 4. cmc Values and Micelle Formation Enthalpies, Entropies, and Gibbs Energies of SDS in 1,4-Butanediol Water Mixtures at 298 and 308 Ka 298 K

308 K


8.1 8.2 8.2 8.3 8.7 8.5

-0.2 -0.4 -0.5 -1.1 -4.9 -8.2

-35.9 -35.9 -35.9 -35.6 -35.4 -35.5

120 119 119 116 102 91

8.3 8.1 8.3 8.3 8.4 9.0

-5.1 -5.9 -6.5 -6.3 -7.1 -13.6

-36.6 -36.7 -36.6 -36.3 -36.1 -35.8

102 100 97 97 94 72


Figure 3. ∆micH° values as a function of the alcohol concentration at 298 K for ([) SDS/1,2-BTD, (9) SDS/2,3-BTD, and (2) SDS/C4OH.10

with water) that lead to micelle formation in these mixedsolvent systems are significantly different from those we have observed for the micellization of SDS in aqueous solutions of C4OH and 1-pentanol (C5OH) reported earlier.10 The ∆micH° values for the SDS/C4OH mixed micellar systems are also plotted in Figures 3 and 4 for comparison. In the paper by Burrows et al.,10 the micellization enthalpies were found to increase with increasing concentration of 1-butanol at 298 and 308 K. For the SDS/ BTD systems, there is a strong tendency toward decreasing micellization enthalpies with increasing diol concentration at both 298 and 308 K. It should be noted that micellization enthalpies for the SDS/BTD systems decrease from 298 to 308 K, at a constant diol concentration, which is in excellent agreement with the literature.8-10 To explain the trends in the micellization enthalpies, we must look at the interactions that occur between the

Figure 4. ∆micH° values as a function of the alcohol concentration at 298 K for ([) SDS/1,3-BTD, (9) SDS/1,4-BTD, and (2) SDS/C4OH.10

different components in solution. Normally, these include interactions that occur between the components in the micellar phase (i.e., SDS-SDS interactions and BTDSDS interactions) as well as interactions that occur between the diols/water and SDS/water. These interactions may be broken down into a hydrophobic portion as well as an electrostatic contribution. The interactions that occur between surfactant and alcohol alkyl chains and water lead to positive contributions to the ∆micH° values. A negative contribution to the micellization enthalpies is due to the interactions between the hydrophobic groups of the alcohols and the surfactant.12 Hence, the relative magnitude of the micellization enthalpy depends on the balance between the hydrophobic effects, electrostatic contributions, and interactions between the alkyl chains of both the surfactant and the alcohol. In the case of alcohol/ surfactant mixed micelles, the dominant contribution to the micelle formation enthalpy at temperatures near 298 K is that from the hydrophobic effects of the alkyl chains of the alcohol and the surfactant. This leads to positive enthalpies of micelle formation in the case of SDS/1butanol and SDS/1-pentanol mixed micelles at 298 K.10 For the SDS/BTD systems, the alcohol alkyl chain contains only four carbon atoms with two hydroxyl groups. We expect that the partition constants (p values) of the BTDs would be small (see below); the alkyl chain interactions in the micellar interior would be dominated by the interactions between the SDS alkyl chains (there would be few SDS/BTD interactions in the micellar interior). In terms of electrostatic contributions, because the BTDs contain two hydroxyl headgroups, they would be very efficient in decreasing the repulsions between the headgroups in the palisade layer. Therefore, we should also observe a large negative electrostatic contribution to the ∆micH° values. For the SDS/BTD systems, the relatively constant ∆micH° values at low weight percent diol are probably due to the additional contributions from the small alkyl chain of the butanediols to hydrophobic effects. We note again that the trend in the micellization enthalpies with increasing alcohol concentration is substantially different from that observed for small chain length alcohols mixed with SDS micelles such as C4OH and C5OH.10 This is primarily due to the presence of the second hydrophilic hydroxyl group on the carbon backbone, which anchors these alkanediols in the palisade region of the micellar interior. As the concentration of the BTDs in the mixed-


Langmuir, Vol. 15, No. 2, 1999 433

Table 5. Spin-Lattice Relaxation Rates (R1) and Partition Constants (p Values) for SDS/BTD Systemsa alcohol

R1(aq) (s-1)

R1(mic) (s-1)

R1p(aq) (s-1)

R1p(mic) (s-1)

p

Kx

∆trG° (kJ mol-1)

1,3-BTD 1,2-BTD 1,4-BTD 2,3-BTD

0.426 0.446 0.408 0.254

0.467 0.550 0.467 0.275

3.17 3.17 3.26 3.16

3.06 3.07 3.10 3.00

0.06 ( 0.04 0.08 ( 0.04 0.08 ( 0.04 0.06 ( 0.04

33 ( 23 43 ( 22 43 ( 22 35 ( 23

-8.5 ( 1.6 -9.3 ( 1.2 -9.3 ( 1.2 -8.8 ( 1.4

a

M(SDS) ) 100 mM; M(PROXYL) ) 10.1 mM; M(BTDs) ) 60 mM.

solvent systems is increased, the ∆micH° values decrease substantially. This is explained in terms of the decrease in the electrostatic repulsions in the headgroup region of the SDS/BTD mixed micelles. As more butanediol is added, the aggregation numbers of the SDS/BTD micelles decrease and the degrees of counterion binding decrease as well.31 This would lead to an increase in the interactions between the alkyl chains as the distance between neighboring charged headgroups is decreased, which is reflected in the negative ∆micH° values. The cmc values, the ∆micG° values, and the ∆micS° values for the SDS/BTD mixtures are also presented in Tables 1-4. We note that for the SDS/1,2-BTD and SDS/2,3BTD systems the cmc values tend to decrease with an increase in the amount of added diol at low weight percent diol. We note, however, that for the SDS/1,3-BTD and SDS/1,4-BTD systems the cmc values increase slightly with an increasing amount of diol in the mixed-solvent system. This is due to the differences in the structuremaking ability of the 1,2-BTD and 2,3-BTD in water and the structure-breaking ability of the 1,4-BTD and 1,3BTD in water. The ∆micG° values were obtained from the application of eq 4 using the calorimetrically derived cmc values and the R values from conductivity measurements.31 It can be seen that for the SDS/1,2-BTD and SDS/ 2,3-BTD systems the ∆micG° values tend to decrease as the amount of diol in the mixed solvent is increased. For the SDS/1,4-BTD and SDS/1,3-BTD systems, the ∆micG° values appear to increase slightly. In the case of the micellization entropies, we observe that for the SDS/BTD systems at 298 K, the ∆micS° values remain essentially constant at low diol concentrations and decrease substantially thereafter. The contributions to the ∆micS° values are from the partial molar entropy changes of the water, surfactant, and alcohol upon micelle formation. For the water molecules, the partial molar entropy change is positive, due to the release of structured water around the surfactant and alcohol hydrophobic chains (hydrophobic effects). For the surfactant and the alcohol molecules, the partial molar entropy changes upon micelle formation of both components depend on the number of surfactant and alcohol molecules per micelle (i.e., the micelle size) and the conformational changes incurred by the surfactant and alcohol upon micelle formation. Hence, the relative magnitude of the ∆micS° value depends on the magnitude and sign of the entropy changes for the surfactant, alcohol, and water upon the formation of the aggregates. The dominant contribution to the micelle formation entropy for most of the systems investigated here is likely the entropy increase of water, as almost all of the ∆micS° values are positive in the present work. The partition constants of all of the BTDs are the same within experimental error (see below), and the total aggregation numbers of the micelles change little with increasing diol concentration.31 Therefore, the contribution to the ∆micS° values from the surfactant and alcohol aggregating would be essentially the same for each diol at the same additive concentration. At 298 K, the fact that the ∆micS° values for the SDS/ butanediol systems are fairly constant up to 3 wt % diol indicates the alcohols are making small, additional

contributions to the hydrophobic effects at both 298 and 308 K. In the case of SDS/C4OH mixed micelles, the micelle formation entropies increase significantly upon increasing alcohol concentration in the mixed solvent, which is due to the contribution of the small alcohol chain to the hydrophobic effects. This difference in the trend in the ∆micS° values between the monohydroxy alcohol and diol systems is due to the presence of the additional hydroxyl group on the alcohol chain, which would lead to a decrease in the amount of ordered water molecules around the hydrophobic chains. Hence, the contribution to the hydrophobic effects from the SDS alkyl chains dominates. At higher diol concentrations, the decrease in the entropies of micelle formation is due to the breakdown of the water structure by the diol. 2. NMR-PRE Experiments. The degree of solubilization, p, of the BTD molecules in SDS micelles can be obtained from NMR relaxation experiments using the equation18

p)1-

R1p(obs) - R1(obs) R1p(aq) - R1(aq)

(6)

Here, R1p(obs) and R1(obs) are the spin-lattice relaxation rates for the BTD molecules in micellar solution in the presence and absence of paramagnetic ions, respectively, while R1p(aq) and R1(aq) are the BTD relaxation rates in aqueous solution with and without paramagnetic ions, respectively (R1 ) 1/T1). These values are presented in Table 5 for the isomeric BTDs in a 100 mM SDS solution. From the p value, we have calculated the mole fraction based distribution coefficient, Kx, defined as

Kx )

Xa,mic Xa,aq

(7)

where Xa,mic and Xa,aq are the mole fractions of solubilizate in the micellar and aqueous phases, respectively, defined by

Xa,mic ) Xa,aq )

pcs,t pcs,t + (csurf,t - csurf,mon) (1 - p)cs,t

[(1 - p)cs,t + csurf,mon + cD2O]

(8)

(9)

where csurf,t and csurf,mon are the total and the monomeric concentrations of surfactant, respectively, cD2O is the solvent concentration, and cs,t is the solubilizate concentration in molar units. Equation 10 can be approximated as

Xa,aq )

(1 - p)cs,t cD2O

(10)

The Gibbs energy of transfer of the alcohol from the aqueous phase to the micellar phase (the ∆trG° values)

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Table 6. 13C Chemical Shift Changes for the Surfactant (∆δ(SDS) ( 0.02 ppm) as a Function of the Alcohol Concentration carbon number 1 2 3 4 8 9 10 11 12 1 2 3 4 8 9 10 11 12

1.0%

2.0%

3.0%

5.0%

1,2-BTD -0.07 -0.11 -0.17 -0.25 -0.02 -0.03 -0.02 -0.01 -0.02 -0.01 0.00 0.01 -0.02 -0.03 -0.05 -0.07 -0.02 -0.02 -0.03 -0.05 -0.01 -0.02 -0.02 -0.03 -0.03 -0.03 -0.04 -0.07 -0.02 -0.03 -0.04 -0.06 -0.03 -0.03 -0.04 -0.05 -0.04 0.00 0.01 -0.01 0.00 0.00 -0.01 -0.01 0.00

-0.09 0.00 0.01 -0.02 -0.02 -0.01 -0.02 -0.02 -0.01

2,3-BTD -0.12 -0.20 0.01 0.01 0.02 0.02 -0.04 -0.07 -0.02 -0.05 -0.01 -0.03 -0.02 -0.05 -0.02 -0.04 -0.07 -0.02

7.0% -0.32 -0.01 0.01 -0.09 -0.08 -0.05 -0.09 -0.07 -0.06 -0.26 -0.03 -0.01 -0.08 -0.06 -0.04 -0.07 -0.06 -0.04

9.0%

10.0%

-0.37 -0.06 0.01 -0.11 -0.11 -0.06 -0.10 -0.08 -0.06 -0.32 -0.03 -0.01 -0.10 -0.09 -0.06 -0.09 -0.08 -0.05

-0.29 0.07 0.10 -0.03 -0.23 -0.01 0.00 0.03 -0.35 -0.01 0.01 -0.11 -0.11 -0.07 -0.09 -0.07 -0.03

1 2 3 4 8 9 10 11 12

-0.05 -0.01 -0.01 -0.02 -0.02 -0.01 -0.02 -0.01 -0.01

-0.08 0.00 0.01 -0.02 -0.02 -0.01 -0.02 -0.01 0.00

1,3-BTD -0.12 -0.18 0.00 0.00 0.00 0.00 -0.04 -0.07 -0.03 -0.06 -0.02 -0.04 -0.03 -0.05 -0.02 -0.04 0.00 -0.01

-0.23 -0.04 -0.02 -0.09 -0.07 -0.06 -0.08 -0.06 -0.03

-0.27 -0.05 -0.03 -0.11 -0.10 -0.07 -0.09 -0.08 -0.04

-0.31 -0.02 -0.01 -0.13 -0.12 -0.09 -0.10 -0.09 -0.02

1 2 3 4 8 9 10 11 12

-0.05 -0.02 0.00 0.00 -0.01 -0.01 -0.01 0.00 0.00

-0.07 -0.03 0.01 0.01 -0.02 -0.02 -0.02 -0.01 0.00

1,4-BTD -0.10 -0.16 -0.04 -0.07 0.00 0.00 0.00 0.00 -0.04 -0.06 -0.03 -0.04 -0.03 -0.05 -0.03 -0.03 -0.01 -0.01

-0.24 -0.10 -0.04 -0.03 -0.09 -0.07 -0.08 -0.07 -0.04

-0.28 -0.13 -0.05 -0.04 -0.11 -0.09 -0.10 -0.08 -0.05

-0.26 -0.18 -0.01 -0.01 -0.11 -0.08 -0.09 -0.06 -0.03

carbon number

0.50%

1 2 3 4 8 9 10 11 12

-0.10 -0.04 -0.01 -0.01 -0.02 -0.04 -0.03 -0.04 -0.06

1.00%

3.00%

5.00%

10.00%

1-Butanol (C4OH) -0.20 -0.44 -0.05 -0.11 -0.00 -0.08 0.01 0.03 -0.03 0.02 -0.01 0.04 0.01 0.07 -0.06 -0.11 -0.09 -0.16

-0.55 -0.15 -0.04 -0.02 0.02 0.05 0.08 -0.14 -0.20

-0.60 -0.17 -0.22 -0.19 -0.21 0.05 0.09 -0.16 -0.22

can be calculated from the relation

∆trG° ) -RT ln Kx

(11)

The calculated values of the distribution coefficients, along with the Gibbs energy of transfer of the BTD from the aqueous to the micellar phase, are also presented in Table 5. It can be seen from Table 5 that the p values of the isomeric BTDs are approximately the same, at about 0.08. These p values, which reflect the hydrophobicity of the solubilizate molecule, are substantially decreased from the p value of the monohydroxy alcohol (1-butanol) in SDS micelles.14,21 We observe that the location of the hydroxyl groups on the C4 backbone makes little difference in the

Table 7. 13C Chemical Shift Changes for the Alcohol (∆δ(alcohol) ( 0.02 ppm) as a Function of the Alcohol Concentration carbon number

1.0%

2.0%

3.0%

5.0%

7.0%

9.0%

10.0%

-0.00 -0.05 0.02 0.05

-0.00 -0.05 0.03 0.06

0.12 0.10 0.33 0.13

1 2 3 4

0.01 -0.01 0.22 0.04

0.00 -0.02 0.22 0.03

1,2-BTD 0.00 -0.01 -0.02 -0.03 0.24 0.24 0.03 0.03

1 2

-0.02 -0.02

-0.02 -0.02

2,3-BTD -0.02 -0.02 -0.02 0.00

-0.04 -0.03

-0.04 0.00

-0.01 0.07

1 2 3 4

0.04 0.03 0.03 0.05

0.04 0.03 0.04 0.06

1,3-BTD 0.04 0.04 0.03 0.03 0.04 0.04 0.06 0.06

0.03 0.02 0.03 0.05

0.03 0.01 0.03 0.06

0.03 0.02 0.03 0.05

1 2

-0.02 -0.05

-0.11 -0.05

1,4-BTD -0.02 -0.03 -0.05 0.00

-0.01 0.00

-0.04 0.00

0.02 0.04

carbon number 1 2 3 4

0.50% 0.05 0.21 0.12 -0.02

3.0%

5.0%

10.0%

1-Butanol (C4OH) 0.03 0.00 0.21 0.20 0.12 0.11 -0.03 -0.05

1.0%

-0.01 0.22 0.11 -0.05

-0.02 0.24 0.12 -0.04

Gibbs energy of transfer, at least within the error of the NMR-PRE. From the difference in the p values of the 1,2-BTD and 1-butanol,21 we can estimate that the addition of another hydroxyl group to the C4 backbone results in an increase of 4-6 kJ/mol to the Gibbs transfer energy. This is due to the fact that in the case of the BTD molecules the addition of the second hydrophilic group makes a large positive contribution to the ∆trG° values. Because the total number of carbon atoms is the same for the BTDs and C4OH, this would decrease the driving force for the transfer of the alcohol molecules from the aqueous phase to the micellar phase. 3. Chemical Shift Changes upon Transfer of SDS Micelles from Water to the Water/BTD Mixtures. The 13C chemical shift changes upon micellization can give information about conformational changes that surfactant molecules undergo when forming micelles. Generally, a high-frequency shift is associated with an increase in the number of trans conformers, whereas a low-frequency shift is interpreted in terms of an increasing probability of gauche configurations.23-26 It should be noted, however, that Van der Weerd et al.33 state that the chemical shift changes upon micelle formation can be interpreted in terms of the difference in the van der Waals interactions between the surfactant chains in the micelle vs the interaction of the surfactant chains in the bulk of the solution. In addition, it has been stated in the literature that contributions to the 13C chemical shift changes upon micelle formation may also be due to headgroup interactions and hydration changes.34 However, the conformational arguments for the chemical shift changes upon micelle formation are supported by independent experimental techniques.35-37 (33) Van der Weerd, R. J. E. M.; De Haan, J. W.; van de Ven, L. J. M.; Achten, M.; Buck, H. M. J. Phys. Chem. 1982, 86, 2523. (34) Ahlna¨s, T.; So¨derman, O. Colloids Surf. 1984, 12, 125. (35) Griffin, R. G.; Powers, L.; Pershan, P. S. Biochemistry 1978, 17, 2718. (36) Faucompre´, B.; Landman, B. J. Phys. Chem. 1987, 91, 383. (37) Griffin, R. G.; Powers, L.; Pershan, P. S. Biochemistry 1978, 17, 2718.


Figure 5. ∆δ(SDS) values for the systems (a) SDS/1,2-BTD and (b) SDS/C4OH as a function of the alcohol concentration. (a) SDS/1,2-BTD: ([) 1.0, (9) 2.0, (2) 3.0, (×) 5.0, (*) 7.0, (b) 9.0, and (+) 10.0%. (b) SDS/C4OH: (b) 0.50, ([) 1.0, (2) 3.0, (×) 5.0, and (+) 10.0%.

In the case of the SDS/alcohol/water mixtures, we define the chemical shift changes as follows.

∆δ(SDS) ) δ(0.10 M SDS/% alcohol) δ(0.10 M SDS) (12) Here, δ(0.10 M SDS/% alcohol) is the 13C chemical shift of SDS in an n % alcohol/water mixture and δ(0.10 M SDS) is the chemical shift of SDS in the presence of water only. In a similar fashion, we define the chemical shift change for the alcohol carbons as follows.

∆δ(alcohol) ) δ(0.10 M SDS/% alcohol) δ(% alcohol) (13) The 13C chemical shift changes upon formation of a 0.10 M micellar-containing solution of SDS and the alcohols given in Tables 6 and 7. The chemical shift changes for the surfactant and alcohol for the SDS/C4OH and SDS/ 1,2-BTD mixtures are plotted in Figures 5 and 6. The assignment of the carbon numbers in the SDS micellar solution was accomplished through the 2D HETCORR experiment, and it is in general agreement with the literature.38,39 These spectral assignments are presented (38) Cabane, B. J. Phys. 1981, 42, 847. (39) Chachaty, C. Prog. NMR Spectrosc. 1987, 19, 183.

Langmuir, Vol. 15, No. 2, 1999 435

Figure 6. ∆δ(alcohol) values for the systems (a) SDS/1,2-BTD and (b) SDS/C4OH as a function of the alcohol concentration. (a) SDS/1,2-BTD: ([) 1.0, (9) 2.0, (2) 3.0, (×) 5.0, (*) 7.0, (b) 9.0, and (+) 10.0%. (b) SDS/C4OH: (b) 0.50, ([) 1.0, (2) 3.0, (×) 5.0, and (+) 10.0%.

in Figure 1. It can be seen from Table 6 that for almost all of the systems studied here the SDS carbons undergo low-frequency shifts upon micelle formation in the isomeric BTD/water mixtures. In all cases, however, we see that the 13C chemical shift changes are the most significant in the headgroup region of the mixed micelles. This is due to the fact that all of these molecules would be primarily partitioned in the headgroup region of the SDS micelles. We would also expect that the magnitude of the R carbon chemical shifts of SDS would be directly related to the concentration of the isomeric BTD in the micellar phase. For the SDS/1,2-BTD and SDS/2,3-BTD systems, the magnitude of ∆δ(SDS) for the R carbon of SDS is slightly larger than that for the 1,3-BTD and the 1,4-BTD systems. This indicates that in order to accommodate the methyl groups of 2,3-BTD or the small hydrophobic chain of 1,2BTD there is a slightly larger tendency for the R-CH2 of the surfactant to adopt a gauche conformation. Note that for certain concentrations of 1,2-BTD and 2,3-BTD the C-2 and C-3 of SDS undergo high-frequency shifts, an indication of an increase in the population of trans conformers in the SDS micelles. For the 1,4- and 1,3-BTDs, the magnitudes of the low-frequency shifts are slightly smaller. This is likely due to the manner in which the BTD molecules interact with the SDS micelles, given the fact that the p values for the diols are equivalent from the NMR-PRE.

436 Langmuir, Vol. 15, No. 2, 1999

The chemical shift changes of the BTD carbons can be interpreted as follows. For 1,2-BTD, we see that carbons 3 and 4 in the diol appear to undergo high-frequency shifts, whereas the two headgroup carbons undergo lowfrequency shifts. This is consistent with 1,2-BTD sitting in the micelle with the two hydroxyl groups in the palisade layer of the SDS micelles, while the hydrophobic tail of 1,2-BTD mixes to a small extent with the first couple of carbons of the SDS hydrophobic chain. In the case of 1,3BTD, we observe small, high-frequency shifts for all of the carbons on the BTD backbone. This is consistent with the 1,3-BTD spreading out about the palisade layer of the SDS micelles in a “zigzag” shape, i.e., an all-trans configuration for 1,3-BTD which may be expected from the positions of the hydroxyl groups on the four-carbon backbone. For the 1,4-BTD and 2,3-BTD micelles, we observe very slight low-frequency shifts upon the transfer of the BTD from water to the interior of the SDS micelles, which may indicate that these molecules prefer to adopt a gauche configuration in the micellar headgroup region. It should be noted that for most of these alcohols, the magnitude of the 13C shift changes is small, which is a reflection of the small degrees of solubilization of each of the isomeric BTDs in the SDS micelles. In terms of the micellar thermodynamic properties, it is fairly clear from the 13C chemical shift changes for the SDS carbons that the effects of the added BTDs are indeed localized in the headgroup region of the micelles. This is consistent with our explanation of the large, negative magnitude of the ∆micH° values. The location of the alcohols in the headgroup region is very effective in screening the electrostatic repulsions between neighboring headgroups. This allows for strong interactions between the alkyl chains of the surfactants, which, of course, make a large negative contribution to the micelle formation enthalpies.

McMahon et al.

In the case of the SDS/C4OH mixed micelle, the magnitude of the R-carbon chemical shift change indicates that the alcohol chain penetrates a little more deeply into the micellar interior. Hence, the screening of the electrostatic interactions is not as effective for the monohydroxy alcohol as it is for the diol, and the positive contribution to the micelle formation enthalpy from the hydrophobic effects would dominate. Conclusions The thermodynamic properties of micelle formation, ∆micH°, ∆micS°, and ∆micG°, have been determined for the micellization of SDS in a series of isomeric butanediols. It appears that the position of the second hydroxyl group on the butanediol backbone has a significant impact on the electrostatic and hydrophobic properties of the BTD/ SDS mixed micelles, when compared to SDS/1-butanol mixed micelles. This is reflected in the variation of the thermodynamic properties of micelle formation with increasing amounts of butanediols in the mixed solvent. The results from the NMR experiments indicate that, even though the distribution of the substituted BTDs does not depend on the position of the second hydroxyl group on the four-carbon chain, the manner in which they interact with the micellar interior is different for each of the BTDs. Acknowledgment. The financial support of NSERC (research grants to R.P. and D.G.M.) and St. Francis Xavier University in the grant of a James Chair Research Fellowship to R.P. is greatly appreciated. B.H. and C.A.M. acknowledge the grants of NSERC Undergraduate Summer Research Awards. We are grateful to Don Hooper, Mike Lumsden, and Rod Wasylishen for stimulating discussions. LA971093+

Calorimetric and NMR Investigations of the Micellar Properties of

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