Thermodynamic magnitudes of aqueous solutions of the zwitterionic surfactant CHAPS

Thermodynamic Magnitudes of Aqueous Solutions of the Zwitterionic Surfactant CHAPS M. A R A N Z A Z U P A R T E A R R O Y O , * FI~LIX M. GOI~II,*'I AND ISSA A. K A T I M E t Departments of*Biochemistry and t Physical Chemistry, University of the Basque Country, P.O. Box 644, Bilbao 48080, Spain Received February 5, 1988;acceptedNovember 8, 1988

The thermodynamic behavior of 3-((3-cholamidopropyl)dimethylammonio)-l-propanesulfonate (CHAPS) in aqueous solutions has been studied over a wide range of temperatures (298-363 K) by vapor pressure osmometry.Osmoticand activitycoefficients,excessGibbs free energy,and excessentropy functions have been determined. Apparent aggregationnumbers were estimated from vapor pressure osmometry in the temperature range 298-363 K. The experimental results are compatible with the presence of highlyattractive CHAPS-waterinteractions. The latter are decreasedat highertemperatures; thus micelle formationis facilitated, in accordancewith previous studies. © 1989AcademicPress,Inc. INTRODUCTION The study of surfactants in aqueous solutions is important not only from the physicochemical point of view, but also because of the extensive application of soluble amphiphiles in biochemical studies ( 1 - 3 ) . 3-((3Cholamidopropyl)dimethylammonio)- 1 propanesulfonate (CHAPS) constitutes a recent addition to the list ofbiochemically useful surfactants. First described by Hjelmeland (4) it has found application in the purification of a number of intrinsic membrane proteins ( 5 7). We have recently characterized its interaction with phospholipid bilayers, both at sublytic and lyric concentrations (8). Its critical micellar concentration has also been assessed by a variety of methods and under different conditions (9). In the present paper, we intend to describe the thermodynamics of CHAPS solutions using vapor pressure osmometry (VPO). The application of VPO to the thermodynamic study of binary solute-solvent systems is extensive ( 10-14). VPO offers possibilities To whom correspondence should be addressed.

for the thermodynamic study of systems that cannot be explored using other physical methods (15). The technique allows the determination of highly reproducible thermodynamic parameters. We have used VPO to explore the behavior of CHAPS in aqueous solutions in the 298-363 K temperature range. Our results suggest the existence of an equilibrium between the intermolecular association phen o m e n o n and the participation of water molecules in the micellization process. MATERIALS AND METHODS

CHAPS was obtained from Sigma and used without further purification. Double-distilled, deionized water was used throughout this work. Measurements were carried out in a Knauer Model 1974 vapor pressure o s m o m eter, connected to an analog recorder. The osm o m e t e r was equipped with a universal probe that had been previously calibrated with sodium chloride along the temperature range used in this study. The temperature o f the measuring cell was kept constant within ___0.25 K. In VPO a difference in vapor pressure between solution and pure solvent leads to a 22

0021-9797/89 $3.00 Copyright © 1989 by Academic Press, Inc. All fights of reproduction in any form reserved.

Journal of Colloid and Interface Science, Vol. 132, No. 1, October 1, 1989

23

C H A P S / W A T E R SYSTEMS

temperature difference between two thermistors; this is in turn proportional to a difference in electrical resistance zkR. At least five values were measured for each solution and average values were taken. Drop size was kept constant and equal in both thermistors, within experimental error. RESULTS A N D DISCUSSION

Experimental values of &R have been found for different CHAPS concentrations and temperatures. Plots of zXR vs molality for calibration (NaC1/water) and surfactant (CHAPS/ water) systems have been constructed at nine temperatures between 298 and 363 K; an example is shown in Fig. 1. From these, the experimental (rh) and theoretical (m) molalities can be compared, and osmotic coefficients (q5 = rfi/m) calculated for each temperature and concentration (10 different concentrations, between 0.01 and 0.10m, have been used). The reference system used in this work is an asymmetric one; therefore we can distinguish between solvent (subindex 1 ) and solute (subindex 2). The relationship between the

200

molal osmotic coefficient (qS) and the solute activity coefficient (3`2) c a n be derived from the Gibbs-Duhem equation ( 16-18): In

3`2

~

(q~

--

1)

+ f]

m-ln l

For dilute solutions, the solvent mole fraction Xl. ~ 1 and ln3'2=(q~-

1)+ I n q~- ldm. do

m

[21

Activity coefficients for the solute (3"2) may then be obtained by analytical integration of the above equation. The solvent activity coefficient (3"1) is immediately obtained from the Gibbs-Duhem relation (18), as

ln 3"l = fom X2 d l n 3"2.

[3]

X1

The variation ofln 3"2with temperature for all CHAPS concentrations tested is shown in Fig. 2. Two temperature regions can be clearly distinguished: above 313 K, In 3"2is constant and less negative than below. Then, excess Gibbs free energies of the CHAPS/water systems under study have been calculated from the activity coefficients using the well-known equation AG E =

n l R T l n 3"1 + n 2 R T l n 3"2.

[4]

160

vn<1

120

8C

4£

0

0.04

0.08 c(mol/Kg)

0.12

FIG. 1. The calculation of experimental and theoretical molalities by vapor pressure osmometry. Instrumental response ( A R ) at 298 K is plotted versus concentration for (O) NaC1 and (B) CHAPS in water. The former is supposed to follow an ideal behavior.

Figure 3 shows the variation of the excess Gibbs free energy (G E) as a function of temperature for the molality range studied. The overall negative values obtained suggest a strong compatibility between surfactant and water, which in turn implies that the latter is involved in the mechanism of micelle formation (19). These values are much more negative than those previously found for sodium dodecyl sulfate or sodium tetradecyl sulfate (20), indicating that, in the present case, surfactant-water interactions are stronger than those with the previously mentioned detergents. Those interactions will tend to hinder the self-association process of CHAPS. Another important thermodynamic magJournal of Colloid and Interface Science, Vol. 132,No. 1, October 1, 1989

24

PARTEARROYO, GO]qll, AND KATIME T(K) 303

o

323

T(K) 343

363

~J

0 293 Oj--~ j ,

=

313

333

353

-2

-1.C m= 0.Ol -1.5 c~

-3 _c -2.5

•

-4

-3X rn = 0 . 1 0

FIG. 4. Temperature variation of the solute-solventinteraction parameter x~2.

- 3 . ~=

FIG. 2. The variation of the natural logarithm of the activity coefficientof CHAPS (In 3'2) with temperature, for all solute concentrationstested. Concentrationsincrease from the upper curve(n = 0.01 ) downward,with molality interactions (21 ). As in the case of activity increments Am = 0.01. coefficients (Fig. 2) two temperature regions are detected: above 313 K, x~2 is constant with nitude to be evaluated is the solute-solvent temperature and less negative than below. One interaction parameter X12, calculated by hypothesis that could explain these results inmeans of the Guggenheim-Stokes equation volves the thermotropic transition of the linear hydrocarbon part of CHAPS molecules. Lin(21), ear hydrocarbons are known to exhibit gel-to~ G E = N R T x ] x2X12, [5 ] fluid thermotropic transitions (22) above a where N is the total mole number ( N = n~ critical temperature. An increase in intrachain motion (in our case above 313 K) would de+ n2). The variation of ×]2 with temperature is crease the probability of CHAPS-water intershown in Fig. 4. Large negative values are oh- actions, thus indirectly favoring micelle formined, indicative of attractive CHAPS-water mation. The small but significant decrease in critical micellar concentrations with temperature (9) would support this hypothesis, which T(K) 303 323 343 363 is also in accordance with the more positive values of excess entropy (calculated as AS z = -- d A G E / d T ) found at higher temperatures 100 ~ ~' • " • (Fig. 5). 200. ~ = " " ~' Vapor pressure osmometry can also be used to obtain an estimate of the apparent number= aoo. ~ . . . . average molar mass )l~,,a, thus the apparent aggregation numbers, of CHAPS micelles. For that purpose a calibration constant K v e is cal600. " ~ ' - - " : ~ culated, by extrapolating to zero concentration a zS,R / c vs c plot for a substance of known 600. molecular weight. In our case, NaC1 was used FIG. 3. The variation of excess Gibbs free energy for for calibration, correction factors for nonideCHAPS-water interaction, as a function of temperature, forall soluteconcentrationstested. Concentrationsincrease ality (between 0 and 0.1 m) being applied as from the upper curve(m = 0.01) downward,with molality in (23). Once the Kve value is known, JQ,, increments Am = 0.01. may be obtained by extrapolating to zero conu

Journal of Colloid and Interface Science, Vol. 132, No. 1, October 1, 1989

25

CHAPS/WATER SYSTEMS 298 K

400

TABLE I Values Obtained for Apparent Number-Average Aggregation Numbers (nn,a)of CHAPS in Water, in the 298363 K Temperature Range

200

0

-200 3

o E

T (K)

ffp,a

298 303 308 313 318 323 333 363

3.5 3.6 3.5 3.2 3.0 3.0 3.0 2.6

ol/kg)

-400

% -600

-800

-1000

FIG. 5. The variation of excess entropy for CHAPSwater interaction (plotted as -TAS E) as a function of surfactant concentration (between 0.01 and 0.10 m), for all temperatures tested. Temperatures increase from the upper curve downward as follows: 298, 303, 308, 313, 318, 323, 328, 333, and 363 K.

centration a z~kR/Kvp vs c plot (Fig. 6). Apparent aggregation numbers are summarized in Table I; values around 3 are found. Previous measurements of the aggregation number of CHAPS by 13C N M R gave values of 3.8 _+ 0.8 (24), in good agreement with our data. However, Hjelmeland et al. (25) found an aggregation number of l0 for this surfactant, apparently through surface tension measurements. More recently, Neugebauer and Kratohvil (26) have carried out a detailed study of CHAPS micellization by light scat-

Note. Data are calculated from apparent number-average molar masses. See text for details.

tering, including corrections for the variability of m o n o m e r concentration above the CMC and for the effects of the second virial coefficient. These authors obtain a mass-average aggregation number (tiw) of 9.7 for a 0.0163 M CHAPS micellar concentration. From our own (uncorrected) light scattering data that have been published elsewhere (9), we have estimated an apparent aggregation number (ffw,a) of 7.9 for the concentration range 0.010.03 M , in reasonable agreement with the results in (26). Thus, considering the differences between mass- and number-average values, there is an overall agreement on the data of CHAPS aggregation numbers that is not always found in the study of bile salts (27). ACKNOWLEDGMENTS

1.0

%

This work was supported in part by grants from CAICYT (0992/84) and the Basque Government (X-86.047).

0.8

REFERENCES

ee O.4 0.2

20

40

60

c(g/kg)

FIG. 6. The calculation of number-average molar masses of CHAPS micelles from VPO measurements. The data in the figure were taken at 298 K. Kvr is the calibration constant referred to in (14). See text for additional details.

1. Helenius, A., and Simons, K., Biochim. Biophys. Acta 415, 29 (1975). 2. Helenius, A., McCaslin, D. R., Fries, E., and Tanford, C., in "Methods in Enzymology" (S. Fleischer and L. Packer, Eds.), Vol. 56, p. 734, Academic Press, New York, 1979. 3. Lichtenberg, D., Robson, R. J., and Dennis, E. A., Biochim. Biophys. Acta 737, 285 (1983). 4. Hjelmeland, L. M., Proc. Natl. Acad. Sci. USA 77, 6368 (1980). Journal of Colloid and Interface Science, Vol. 132, No. 1, October 1, 1989

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PARTEARROYO, GOlqI, AND KATIME

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Journalof CollaldandInterfaceScience,Vol.132,No. 1,October1, 1989

17. TS'O, P. O. P., and Chan, S. I., J. Amer. Chem. Soc. 86, 4176 (1964). 18. Lewis, G. N., Randall, M., Pitzer, K. S., and Brewer, L., "Thermodynamics," 2nd ed. McGraw-Hill, New York, 1961. 19. Katime, I., and Allende, J. L., "Surfactants in Solution" (E. Mittal, Ed.), Vol. 4, p. 77. Plenum, New York, 1986. 20. Katime, I., and Allende, J. L., Thermochim. Acta 74, 215 (1984). 21. Guggenheim, E. A., and Stokes, R. H., "Equilibrium Properties." Pergamon, New York, 1969. 22. Chapman, D., Williams, R. M., and Ladbrooke, B. D., Chem. Phys. Lipids 1,445 (1967). 23. Weast, R. C. (Ed.), "Handbook of Chemistry and Physics," !o. D-224. CRC Press, Cleveland, 1974. 24. Stark, R. E., Left, P. D., Milheim, S. G., and Kropf, A., J. Phys. Chem. 88, 6063 (1984). 25. Hjelmeland, L. M., Nebert, D. W., and Osborne, J. C., Anal. Biochem. 130, 72 (1983). 26. Neugebauer, J. M., and Kratohvil, J. P., submitted for publication. 27. Kratohvil, J. P., Adv. Colloid Interface Sci. 26, 131 (1986).