2236
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979
solutions in D20 with a Varian XL-100-15 NMR spectrometer operating a t 100 MHz in the continuous wave mode. Linear dichroism measurements were performed with a JASCO 5-40 circular dichroism spectrometer as described previously." The sample was subjected to a linear gradient in a thermostated Couette cell, consisting of two concentric cylinders separated by 0.05 rnm.'O Acknowledgment. Signe Cravsholt's work was supported by the Danish Natural Science Research Council. References and Notes (1) S. Gravsholt, J . Collaidlnterface Sci., 57, 575 (1976). (2) (a) S. Gravshoit, Proceedings of the VIIth International Congress of Surface Active Agents, Moscow, 1976, Vol. 2 (II), 1978, p 906; (b) S. Gravshott in "Polymer Colloids", Vol. 11, R. Fitch, Ed., Plenum Press, New York, 1979. (3) Every iiquid shows some viscoelastic behavior but the text refers to the spectacular viscoelastic behavior discussed in the Introduction. (4) H. Wennerstrorn and J. Ulmius, J . Magn. Reson., 23, 431 (1976). (5) J. Ulmius and H. Wennerstrbm, J . Magn. Reson., 28, 309 (1977). (6) F, Perrin, J. Phys. Radium, 5, 497 (1934); 7, 1 (1936). (7) K. Fontell, private communication. (8) H. Wennerstrom and B. Lindman, fhys. Rep., 52, 1 (1979). (9) L. B.-A. Johansson, G. Lindblorn, and 8. Nordbn, Chem. Phys. Lett., 39, 128 (1976). (10) L. B.-A. Johansson, G. Lindblom, S. Gravsholt, and B. NordGn, J. Colloid Interface Sci., 69, 358 (1979). (11) L. B.-A. Johansson, A. Davidsson, G. Lindblom, and B. Nordbn, J. fhys. Chem., 82, 2604 (1978). (12) N. Pilpel, Trans. Faraday Soc., 50, 1369 (1954); J . ColloidSci., 9, 285 (1954). (13) N. Pilpel, Trans. Faraday Soc., 62, 1015 (1966). (14) A. J. Hyde and Q. W. M. Johnstone, J . Colloid Interface Sci., 53, 349 (1975). (15) D. Saul, G. J. T. Tiddy, B. A. Wheeler, P. A. Wheeler, and E. Willis, J . Chem. Soc., Faraday Trans. 1 , 70, 163 (1974). (16) 6. A. Barker, D. Saul, G. J. T. Tiddy, B. A. Wheeler, and E. Willis, J . Chem. Soc., Faraday Trans. I , 70, 154 (1974). (17) d. W, Larsen, L. J. Magid, and V. Payton, Tetrahedron Lett., 29, 2663 (1973).
M. Abu-Hamdiyyah and L.
Al-Mansour
(18) T. Nash, J. Colloid Sci., 13, 134 (1958). (19) J. N. Israelachvili, D. J. Mitchell, and B. W. Ninham, J. Chem. Soc., Faraday Trans. 2, 72, 1525 (1976). (20) C. Tanford, J . Phys. Chem., 76, 3020 (1972). (21) C. Tanford, J. fhys. Chem., 78, 2469 (1974). (22) C. Tanford, "The Hydrophobic Effect", Wiley, New York, 1973. (23) G. Lindblom, B. Lindman, and L. Mandell, J . Colloid Interface Sci., 42, 400 (1973). (24) H. Hoffmann and W. Ulbrlcht, Z . fhys. Chem., 106, 167 (1977), and personal communication. (25) J. C. Eriksson and G. Gillberg, Acta Chem. Scad., 20, 2019 (1966). (26) U. Henriksson, Thesis, Stockholm, 1975. (27) P. Ekwall, L. Mandell, and K. Fontell, J. Colloid Interface Sci., 29, 639 (1969). (28) J. Uimius, B. Lindman, G. Lindblom, and T. Drakenberg, J. Collold Interface Sci., 65, 88 (1978). (29) P. Mukerjee, Ber. Bunsenges. ,Phys. Chem., 82, 931 (1978). (30) U. Henriksson, T. Claesson, L. Odberg, and J. C. Eriksson, Chem. fhys. Lett., 52, 554 (1977). (31) H. Wennerstrom, N. 0. Persson, and B. Lindman ACS Symp. Ser., No. 9, 253 (1975). (32) Y. Hamnerius, I. Lundstrom, L. E. Paulsson, K. Fontell, and H. Wennerstrom, Chem. fhys. Lipids, 22, 135 (1978). (33) J. G. Kirkwood, Recl. Trav. Chim. fays-Bas, 68, 649 (1949). (34) J. G. Kirkwood and P. L. Aner, J. Chem. Phys., 19, 281 (1951). (35) J. D. Ferry, "Viscoelastlc Properties of Polymers", 2nd ed., Wiley, New York, 1974. (36) J. D. Ferry, Acc. Chem. Res., 6, 60 (1973). (37) R. Ullrnan, Macromolecules, 2, 27 (1969). (38) I. F. Efremov, Surface Colloid Sci., 8, 85 (1976). (39) I. F. Efremov and 0. G. Us'garov, Usp. Khim., 45, 877 (1976). (40) A. Kose, M. Ozaki, K. Takano, Y. Kobayashi, and S.Hachisu, J. Wbki Interface Sci., 44, 330 (1973). (41) P. A. Hiltner and I. M. Krieger, J. Phys. Chem., 73, 2386 (1969). (42) J. C. Brown, P. N. Pusey, J. W. Goodwin, and R. M. Ottewill, J. Phys. A , 8, 664 (1975). (43) J. P. Cotton and M. Moon, J . Phys., 37, L-75 (1976). (44) H. Kiessig and W. Philipoff, Naturwissenschaften, 27, 593 (1939). (45) J. Stauff, Kolloid. Z., 89, 224 (1939). (46) K. Hess and J. Gundermann, Berichte, 70, 1800 (1946). (47) G. S. Hartley, Nature (London), 163, 787 (1949). (48) E. Wigner, fhys. Rev., 46, 1002 (1934). (49) R. Williams and R. S. Crandall, Phys. Lett. A , 48, 225 (1974). (50) D. G. de Gennes, P. Pincus, R. M. Vebsco, and F. Brochard, J. Phys., 37, 1461 (1976).
utylurea on the Critical Micelle Concentration of Sodium Lauryl Sulfate in Water Mohammad Abu-Hamdlyyah* and Lulua Al-Mansour Chemistry Department, University of Kuwait, Kuwait (Received January 75, 7979: Revised Manuscript Received April 72, 1979) Publication costs assisted by the University of Kuwait
The critical micelle concentration (cmc) of sodium lauryl sulfate (NaLS) was determined in aqueous n-butylurea (BU) solutions at 15,20,25,30,35, and 45 "C. A t each of these temperatures the cmc initially decreases, reaches a minimum, and then starts to increase with increasing BU concentration. Corresponding changes in the thermodynamic parameters of micellization show similar trends except at 45 "C where both AH",and ASo, tend to increase only with increasing BU concentration. The degree of micelle ionization (a)was calculated from the slopes of the conductance curves and shows the characteristic behavior of increasing, reaching a maximum, and then decreasing as the concentration of additive is increased. This was explained as due to increasing BU solubilization in the micelle, reaching a maximum, and then decreasing as the BU concentration is increased. The results are discussed in terms of the effect of BU and temperature on the structure of the solvent and of micelle penetration. n-Butylurea is strongly adsorbed at water-air and water-dodecane interfaces and increases the dielectric constant of water, the molar increment being about 3 f 1 at 20 "C. Introduction The spontaneous association of a large number of surfactant molecules (ions) (micelle formation) to give a large aggregate (micelle) in aqueous solutions is often taken as a model for hydrophobic bonding1 which is believed to be partly responsible for the stability of the native state of proteins2 and most likely must have been the first step
in the genesis of the living cell. The critical concentration at which aggregation occurs (crnc) reflects the balance of the various factors governing micelli~ation~,~ and thus is an important parameter to study. Urea, which is well known denaturant of proteins, increases the cmc of ionic and nonionic surfactant^.^^^^ This weakening of hydrophobic bonding by urea must be related
0022-3654/79/2083-2236$O?.QC/O0 1979 American Chemical Society
Cmc of
NaLS
in n-Butylurea Solutions
to its hydrogen bonding ability with water, forming mixed clustersS1l such that cavity formation for the accommodation of the monomeric hydrophobic groups or molecules becomes easier as these clusters around the hydrophobic groups have less regular structures than those formed around hydrophobic groups in pure water.12J3 Addition of an alkyl-substituted urea to liquid water introduces an extra factor, the effect of the alkyl (hydrophobic) group, which would oppose the effect of the hydrophilic (urea) part, and compete with the hydrophobic group of the surfactant since nonpolar groups tend to dissolve interstitially in water.8 This effect depends on the relative size of the alkyl group as in the case of alkanols,' except that the hydrophilic effect in alkylureas would be stronger in view of the larger hydrophilic head (urea) and larger hydrogen-bonding capacity which means that an alkanol would be more hydrophobic than an alkylurea with same alkyl group. It is expected that alkylureas with large hydrophobic groups would behave like the corresponding alkanols and would strengthen hydrophobic bonding at low concentrations and weaken it at high concentrations. Alkyl-substituted ureas have been reported both to be ineffe~tive'~ and effective15denaturants. However, studies on the effect of alkyl-substituted ureas on the cmc of surfactants in water reported in the literature1@ showed only an increase in the cmc (weakening of hydrophobic bonding) at all concentrations examined. The substituted ureas in the latter studies were methyl and ethyl ureas which are not very effective denaturants.17 In the present work, a study of the effect of BU on the cmc of NaLS in aqueous solutions at several temperatures was carried out. n-Butylurea has a relatively large alkyl substitutent, is moderately soluble, and has been found to be a very effective denaturant of proteins.15
Experimental Section Chemicals. n-Butylurea, purum quality, purchased from Fluka, was recrystallized from hot benzene. The melting point of the recrystallized product was 97 "C. Sodium lauryl sulfate was a "special pure" quality purchased from British Drug Houses and was used without further purification. Its purity was shown to be satisfactory as the cmc values obtained in water agreed with those in the literature.18 Dodecane, acetone, methanol, nitrobenzene, and ethylene glycol were all analytical reagents and used without further purification. Sodium chloride was of puriss quality. Apparatus and Procedure. Electrical conductance was used to determine the cmc values in the various solutions. The LKB 5300 A conductolyzer was used as a manually operated conductivity bridge, with flow cells which have platinum electrodes, with inlet coils 1 m long and 1-mm i.d. Solutions were injected in the cell which was immersed in a thermostated water bath and allowed to equilibrate before measuring the resistance. The temperature of the water bath was controlled by a proportional controller to f0.02 "C as checked by a Hewlett-Packard digital thermometer. For most of the temperatures investigated the solutions were preequilibriated at the desired temperature prior to injection in the cell. The resistance was measured a t 2000 Hz. All solutions were made up volumetrically. For molar fraction calculations densities of BU solutions were measured with a digital densitometer. Surface tension measurements were carried out with a du Nouy tensiometer. For the partition coefficient of BU between water and dodecane, aqueous BU solutions of different concentrations were shaken with equal volumes of liquid dodecane for 48 h and then the two layers were separated and analyzed by
The Journal of Physical Chemistry, Vol. 83, No. 17, 7979 2237
too90
~
80-
I
I
'
O
b
I 10
0
10
20
40
30
50
60
70
t'c
Figure 1. Slopes of specific conductance-concentration curves for NaLS in water below the cmc (a) and above the cmc (b) as obtained by Goddard and BensonZ0(X), Evans" (O), Mukerjee et aLZ2(e),and this work (0).
refractometry. The dielectric constants of BU aqueous solutions were measured at 20 "C with a Dekkameter DK 03 and Cell MFL3 (Wissenschaftich-Technische Werkstatten). The instrument was calibrated with acetone, methanol, nitrobenzene, ethylene glycol, and distilleddeionized water. For the solubility of BU in water, a saturated solution with excess solid was shaken for more than 50 h at the desired temperature and then left to settle for about 48 h in the thermostat. A certain amount of the filtered supernatant liquid was taken and analyzed by refractometry a t 25 "C.
Results The cmc of NaLS in each solution was determined from the intersection of the two straight lines obtained on plotting specific conductance vs. concentration of NaLS. The values of the coefficients a and b of the straight lines k = a + bc below and above the cmc were obtained by the least-squares method are available as supplementary material.lg Below the cmc the slopes of the straight lines, at a given temperature, are independent of butylurea concentration. Specific conductance data for the NaLS-H20 system were also reported by Goddard and BensonZ0and their data and those of others21,22are compared with our data in Figure 1,noting that Goddard and Benson used molality and we are using molarity. Keeping this in mind, our values for the slopes below the cmc are systematically lower than those of Goddard and Benson. The slopes of the lines above the crnc in the absence of BU are, however, in reasonable agreement with those of Goddard and Benson except at 45 "C where our value is lower by about 10%. In the presence of BU the slopes of the lines above the cmc are dependent on the concentration of the additive as shown in Figure 2. The slopes increase almost linearly with increasing butylurea concentration until about 0.2 M, thereafter the rate of this increase decreases and the slopes tend to reach a maximum and then decrease. At each temperature the cmc decreases as the concentration of BU increases, reaches a minimum, and then
2238
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979
0.1
0.2
0.3
0.4
0.5
M. Abu-Hamdiyyah and L. AI-Mansour
0.6
Butyturra ( m o t . I - ' )
Flgure 2. Slopes of specific conductance-concentration curves above the cmc at different temperatures.
starts to increase as shown in Figure 3. The minimum in the cmc-BU concentration curve occurs at 0.24 M at 15 " C and a t 0.27 M at 45 "C, i.e., tends to shift to higher BU concentrations with increasing temperature. When the cmc values are plotted against the temperature we get in the absence of BU the typical characteristic curve with a shallow minimum and these cmc values are in good agreement with those in the literature.18p20The shallow minimum appears a t about 25 " C in agreement with the literature.lVm Addition of BU affects the shape of this curve such that the negative slope gradually disappears and the curve tends to become linear as the additive concentration is increased. The preliminary results on the surface and interfacial tensions show that n-butylurea is strongly adsorbed at the water-air and water-dodecane interfaces. Figure 4 shows the variation of the surface tension with the logarithm of
additive concentration. The interfacial tension is less than the surface tension by about 20 dyn cm-l even in the absence of the additive. After shaking aqueous BU solution with liquid dodecane for 48 h no significant distribution of BU in liquid dodecane was observed. The concentration of BU in the water layer was unaffected by mixing it with dodecane. The solubility of BU in water at 7.3, 15.5, and 24.5 " C was found to be 0.54, 0.71, and 1.1 mol/L, respectively. Ignoring activity coefficients which we do not know, and converting to mole fractions we get from the slope of log (solubility) vs. (TI)the standard enthalpy of BU dissolution in water, 6.7 kcal mol-I. AG," = 2.6, 2.5, and 2.3 and TAS," = 4.1, 4.2, and 4.4 kcal mol-', respectively, at 7.7, 15.5, and 24.5 "C. The dielectric constant of aqueous butylurea solutions was measured at 20 "C. The meter readings were 156 f 0.5, 934 f 0.5, 1023 f 0.7, 1331 f 1, 3435 f 1, 3458 f 0.1, 3458, 3463.8 f 0.2, 3503 f 0.3, and 3547 f 1.5 for acetone ( t = 21.4), methanol ( t = 33.6), nitrobenzene ( E = 35.7), ethylene glycol ( t = 39.8), deionized water ( 6 = 80.4), 0.05 M BU, 0.20 M Bu, 0.28 M BU, 0.40 M BU, and 0.50 M BU, respectively. From this a molar increment of about 3 f 1 for BU in water is obtained. Thermodynamic Quantities of Micellization. Micelle formation has been shown to be an association dissociation equilibrium following the law of mass a ~ t i o n Some. ~ ~ ~ ~ ~ times it is aproximated as a phase ~ e p a r a t i o n . ~ The ~ standard thermodynamical quantities estimated depend on the formulation and model used for micelle formation.25 In the following the subscripts ma indicate mass action, ps phase separation, and CY the degree of micelle ionization. We have25for partly ionized micelles AGO,, = (2 - a)RT. In (crnc),for electroneutral micelles AGO,, = AGO, = 2RTIn (crnc) and for fully ionized micelles AGO,, = AGO,,, = R T In (cmc), where AGO,,, is the standard free energy change accompanying the addition of one monomer to a micelle of the most probable si2e.l All cmc concentrations are in mole fraction units and the standard state of any species in solution is that of mole fraction unity such that the solution properties are those of infinitely dilute solutions. The corresponding standard enthalpy changes are AH",,= -(2 - a)R% d In (cmc)/dT, for partly ionized micelle, AH",, = AHo,, = -2RT2 d In (cmc)/dT, for
I
' 1
20
0.1
0.2
0.3
0.4
0.5
I.
0 0.1 Butylurea ( m o l . P
Flgure 3. Effect of n-butylurea on the crnc of NaLS in water at different temperatures.
0.2
0.3
0.4
0.5
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979 2239
Cmc of NaLS in n-Butylurea Solutions
TABLE II: Thermodynamics of Micellization of NaLS in Aqueous n-Butylurea Solutions at Different Temperatures Assuming Electroneutral Micelles _ I
t, "C
(BU)
15
H2O 0.05 0.10 0.20 0.30
20
H2O 0.05 0.10 0.20 0.30 0.40 H2O 0.05 0.10 0.20 0.30 0.40 H2O 0.05 0.10 0.20 0.30 0.40 H2O 0.05 0.10 0.20 0.30 0.40 H2O 0.05 0.10 0.20 0.30 0.40
25
30
35
45
kcal mol-' -10.13 t 0.02 -10.43 i 0.03 -10.74 t 0.04 -11.06 t 0.04 -10.96 i 0.05 -10.32 f 0.01 -10.62 1 0.01 -10.89 f 0.02 -11.26 t 0.04 -11.05 t 0.05 -10.88 f 0.05 -10.51 t 0.01 -10.76 i 0.01 -10.98 f 0.02 -11.29 i 0.02 -11.24 j: 0.05 -11.00 i 0.05 -10.70 i. 0.02 -10.96 t 0.02 -11.18 i 0.02 -11.53 i 0.02 -11.45 t 0.01 -11.18 i. 0.05 -10.86 i 0.02 -11.10 t 0.02 -11.34 t 0.02 -11.72 t 0.02 -11.71 t 0.03 -11.12 j: 0.05 -11.08 i 0.02 -11.33 i 0.03 -11.56 i 0.04 -11.75 f 0.04 -11.83 t 0.04 -11.82 i 0.05
AGO,,,
'Ot
3o
%oz' 0
t -3
-2
-I
In
E
Figure 4. Surface tension of aqueous butylurea solutlons at 20 "C vs. natural logarithm of concentration.
electroneutral micelles, and AH",, = AHo,,, = -RTZ d In (cmc)/dT for fully ionized micelles, assuming that in the temperature range examined the aggregation number is constant. Standard entropy changes are obtained from AGO, and AH",. Furthermore solutions a t the cmc may
kcal mol-' 2.2 i 0.3 1.0 f 0.3 -2.5 t 0.3 -3.2 i 0.9 -1.4 t 1.3 1.5 i 0.1 0.02 i. 0.1 -2.6 i 0.3 -3.3 f 0.9 -1.4 t 1.4 -0.7 t 0.6 +0.21 t 0.03 -0.9 c 0.1 -2.7 i 0.4 -3.4 t 0.9 -1.7 t 1.7 -0.7 f 0.6 -1.0 * 0.1 -1.8 t 0.3 -2.7 i 0.4 -3.5 i 1.0 -1.5 i 1.5 -0.7 t 0.6 -2.2 i 0.3 -2.7 i 0.3 -2.8 f 0.4 -3.6 f 1.0 -1.6 f 1.5 -0.8 ?: 0.6 -4.5 i 0.1 -4.3 i 0.3 -2.3 i. 0.3 -3.9 j: 1.1 -1.7 i 1.6 -0.3 f 0.7
AH',,,
ASOrna,
cal mol K-' 42.8 39.6 28.5 27.4 33.0 40.3 36.2 28.3 27.3 32.8 34.8 35.9 33.2 27.8 26.5 31.9 34.6 32.0 27.1 28.0 26.4 33.0 34.1 28.2 27.3 27.6 26.3 32.8 34.7 19.6 21.7 28.5 24.4 31.5 33.1
Asma,
cal mol K 1 7.6 3.5 -8.7 -11.1 -4.9 5.1 0.0 -8.9 -11.2 -4.8 -2.4 +0.7 -3.0 -9.1 -11.4 -5.7 -2.3 -3.3 -5.9 -8.9 -11.6 -5.0 -2.7 -1.1 -8.8 -9.1 -11.7 -5.2 -2.6 -14.6 -13.4 -1.3 -122 -5.3 -2.5
be taken as infinitely dilute so AH,, = AH",, and AS,, = AH,,/T. Thus adopting the phase separation model or assuming electroneutral or fully ionized micelles does not require knowledge of cy for the estimation of the changes in the standard thermodynamic quantities of micellization. The adoption of electroneutral micelles is j u ~ t i f i a b l e . ~ ~ ~ ~ ' Fully ionized micelles, on the other hand, although unlikely, were used in the l i t e r a t ~ r e . The ~ ~ ~phase ~~~~~~~ separation model leads to results similar to those results assuming electroneutral micelles. The standard changes in these functions have been estimated for electroneutral micelles and are shown in Table 11. Inclusion of the neglected terms, the concentration of micelles at the cmc, and the assumption of unity activity coefficients did not show any significant changes in AGO, values from which all other function were derived. Table 11 shows that at each temperature except 45 "C each of the thermodynamic functions decreases, reaches a minimum, and then increases as BU concentration increases. At 45 "C AH" and ASo both tend to increase monotonously with increasing BU concentration. The enthalpy and entropy of NaLS micellization in water are positive a t low temperatures and become less positive as the temperature increases and eventually both AH and A S become negative, passing through zero at about 25 "C. Addition of butylurea is equivalent to raising the temperature; it makes the enthalpy and entropy of micellization less positive for BU concentrations up to about 0.2 M after which further addition would increase (make less negative) AH, or AS,. However, again at 45 " C AH," and AS, both become less negative and AS," tends to become more positive as the BU concentration increases
2240
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979
M. Abu-Hamdiyyah and L. AI-Mansour
-9 1
Ln ( C o u n t e r I o n s 1
Flgure 5. Log monomeric activity (mole fraction) at the cmc vs. log counterion activity (mole fraction) in the sodium lauryl sulfate-butylurea (0.2 M)-water-NaCI system at 15 "C.
TABLE 111: Degree of Ionization of NaLS Micelles ( a )at Different Temperatures and Butylurea Concentrations Calculated According to Evans 0.10 0.20 0.30 0.35 0.40 0.50 t, "C H,O M M M M M 15 0.23 0.28 0.34 0.36 0.37 25 0.23 0.30 0.35 0.37 0.36 35 0.24 0.32 0.37 0.40 0.38 0.36 *'0.24 0.34) 0.35 0.37 0.37 0.37
trimethylammonium bromide as a function of alcohol concentration in aqueous solution and found that a increases, reaches a maximum, and then decreases as the concentration of ethyl, isopropyl, or tert-butyl alcohol increases. Methanol only increased a. M i y a g i ~ h i ~ ~ measured the degree of dissociation of dodecylammonium chloride as a function of propanol (and of acetone) concentration and found similar behavior, i.e., a increasing, reaching a maximum, and then decreasing with increasing additive concentration. and the minimum disappears. The standard entropy of The degree of ionization was also determined experimicellization is positive at all temperatures and concenmentally with Corrin's From the law of mass trations investigated. actionz5 Despite the good reasons advocated for the adoption of AG",, 1 electroneutral micelles, many studiesz2show that micelles In (cmc) = -(I - a ) In (Na+) - - In F(M") RT n of ionic surfactants are partly ionized. Degree of Micelle Ionization. In order to estimate the where (crnc), (Na+), and (M") are the activities of the unassociated lauryl sulfate ions, total activity of counchanges in the thermodynamic functions of micellization terions, and the concentration of micelles a t the cmc, all for partly ionized micelles the degree of ionization ( a )is in mole fraction units. F is the activity coefficient of the necessary. This was determined by the Evans methodz1 from the slopes of the conductance curves, assuming micelles. The cmc of NaLS was determined in 0.20 M BU as a function of NaCl concentration at 15 "C. Figure 5 monodisperse micelles, with constant aggregation number shows In (crnc) to be linear with In (Na+),the slope equals of lauryl sulfate ions equal to 62 in water30 and in presence -(1 - a ) , and the intercept equals AG,;/RT + l / n In of butylurea, in the temperature range 15-45 "C. The F(M"), which are -0.54 f 0.04 and -15.15 i0.32, remicelles in aqueous butylurea solutions would contain some spectively. Thus at 15 "C in presence of 0.2 M BU the of the amphiphilic additive but the size of the micelle is degree of NaLS micelle ionization equals 0.46 f 0.04. The dependent on the longer and more hydrophobic component linearity obtained indicates either the last two terms in (LS-), the additive diluting the surface charge only. the equation are constant or we have self-compensating Table I11 shows the values of a calculated accordingly. changes.22 The values of A(Na+) used in the calculations were calSeveral investigators have determined a for NaLS culated a t the ionic strengths corresponding to the difmicelles in water by this method. Mukerjee et aLZ2obferent cmc values by using the Onsager equation. ho(Na+) tained a = 0.31 a t 25 "C, H u i ~ m a nfound ~ ~ a = 0.30 at 21 values at different temperatures were taken from the OC, and Mori et found a = 0.32 aat 30 "C and a = 0.29 l i t e r a t ~ r e .A(Nat) ~~ in aqueous BU was found to be little at 15 and 40 "C. Therefore it is reasonable to take a = different from its value in pure water; for example, it was 0.30 in water for NaLS micelles in the absence of butylurea 1%less in 0.1 M BU than in pure water. in the temperature range investigated. Comparing this The values of a obtained in pure water are in good value of a in water alone with that obtained in 0.20 M BU agreement with the experimental values obtained by by the same procedure, we can see that addition of BU electrokinetic2*and by emP2methods. Table I11 also shows causes an increase of about 50% in a. The same relative that a increases almost linearly with butylurea concenincrease is obtained when a was calculated by the Evans tration up to 0.2 M then tends to reach a maximum and method as seen in Table 111, although the absolute values then decrease at high concentrations. Such behavior has been experimentally found in the l i t e r a t ~ r e Larsen . ~ ~ ~ ~ ~ of a are different. That different techniques give different values of a for the same system under same conditions was and T e ~ l e determined y~~ the degree of ionization of cetyl
+
+
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979 2241
Cmc of NaLS in n-Butylurea Solutions
3r
4i
-31t
-4
-5
-4
-
-5
I
-6
0
10
20
40
30
50
60
I
20
10
70
noted and discussed by Mukerjee22and by A n a ~ k e r . ~ ~ Assuming the concentration of micelles at the cmc to be 2% of the tital surfactant, n = 62, ignoring the term l / n In F , and calculating the mean ionic activity coefficients of the monomers and of sodium chloride a t the appropriate ionic strength, we estimated the standard free energy changes in the NaLS-H,O-NaC1-BU system at 15 "C. These are -8.39, -8.55, -8.65, and 8.54 kcal mol-' in aqueous butylurea (0.20 M) solutions containing 0, 0.001, 0.01, 0.1 M NaC1, respectively. In 0.10 M NaCl using n = 82 as obtained by light scattering30 did not produce a significant difference in the value of the free energy change. Thermodynamic parameters of NaLS micellization at different temperatures and different BU concentrations were estimated for partly ionized micelles with CY values obtained by the Evans method. The same trends noted earlier for electroneutral micelles are also obtained here. Figure 6 compares the enthalpy of micellization of NaLS in water obtained by the van't Hoff method by different workers as a function of temperature assuming (Y = 1. Figure 7 compares the enthalpy of micellization in water obtained in this work for partly ionized and electroneutral micelles and that obtained calorimetrically by Kresheck et al.29 Of the two models, the partly ionized micelle appears to give results nearer to those obtained calorimetrically. On the whole our results tend to be more positive on the low temperature and more negative on the high temperature side. Nevertheless, all the values of the standard enthalpies for partly ionized (or electroneutral) micelles, obtained in the temperature range 15-45 " C at the different BU concentrations, when plotted against the corresponding standard entropy changes, fall on a straight line,37Figure 8, with a slope corresponding to a compensation temperature of 308 f 7 K (308 f 13 K). The plot for electroneutral micelles shows more scatter. The corresponding intercepts AHx are -9.5 f 0.2 and -11.0 f 0.4 kcal mol-I for the partly ionized and electroneutral micelles, respectively. Figure 9 shows the variation of AH", for partly ionized micelles in aqueous BU solutions, with temperature. In the absence of BU, AHo, decreases sharply with increasing
60
t0c
t'c
Flgure 6. Enthalpy of NaLS micellization in water at different temperatures as obtained by the van't Hoff method: this work (A), Moroi et (0),Goddard and Benson*' (X), and Flockhart28(O), assuming no counterion binding in all cases.
40
Figure 7. Enthalpy of NaLS micellization at different temperatures obtained in this work by the van't Hoff method [electroneutral micelles (O), partly ionized micelles (O)] com ared with that obtained directly, calorimetrically, by Kresheck et al.' (X).
t:
3-
2-
1-
-L
0-
0
-Eu
-1
-
0
-2-
.E I 4
-3
-
-4t / -5
1
I
I
10
20
30
40
AS'm ( c a l . m o i ' . d r p l \
Figure 8. Enthalpy-entropy compensation plot for mlcellization of NaLS in aqueous butylurea solutions in the temperature range 15-45 OC, assuming partly inonized micelles.
temperature. The effect of adding BU is to reduce the rate of change of AH", with temperature. A similar plot is obtained for electroneutral micelles. From these results the changes in heat capacity on micellization ACm, in the various solutions were estimated. Despite the large uncertainties involved in the absolute values of ACmpderived from AHo,which were themselves obtained by the Van't Hoff method, a definite trend appears. ACmpbecomes less negative as the concentration of BU in the solution increases. The effect is large at low BU concentrations and tends to level out at high concentrations. For partly ionized micelles, ACm values are -200, -129, -28, -23, -11, and -6 cal deg-l mol-f , in pure water, and 0.05, 0.10,0.20, 0.30, and 0.40 M BU concentrations, respectively. The
2242
The Journal of Physical Chemistry, Vol. 83, No. 17, 1979
M. Abu-Hamdiyyah and L. AI-Mansour
a t low, the other at high concentrations of the additive. At low concentrations of the additive the hydrophobic part dissolves mainly interstitially and thus on adding NaLS +2 to this solution both additive and surfactant monomers would coaggregate as the limit of interstitial mixing is *I attained, with the resultant formation of micelles at a lower concentration (of NaLS) than in pure water. This might be called a “phobing out” effect. This effect is also seen - ‘2 0 in the absence of micelles as the solubility of KLS was E found to be less, in presence of methanol, ethanol, or : -1 propanol, than in pure water, below a certain temperas t ~ r e . The ~ ~ reduction i ~ ~ in the cmc of ionic surfactants by dissolved hydrocarbons in aqueous s o l ~ t i o n s ~most ~-~~ 3 I -2 a probably follows a similar mechanism except that the hydrocarbon would be dissolved in the interior of the -3 micelles and the surface charge density would not be significantly affected, as indicated e ~ p e r i m e n t a l l yfor ~~ -4 CCl& At high concentration of the additive the structure of 15 20 25 30 35 40 45 SO water is sufficiently disturbed that regular mixing of t’C additive and water molecules becomes dominant so that Figure 9. AHo, vs. temperature for partly ionized micelles. when a surfactant is added to such a mixture it tends to dissolve regularly too, so that the driving force for micelle corresponding ACmpvalues for fully ionized micelles are formation disappears. At intermediate concentrations of -121, -90, -17, -11, -5, and -2 cal deg-’ mol-l. the additive both tendencies occur and when they are Krescheck et and Musbally et a1.38,39determined balanced a minimum appears in the cmc-additive conAH”, for NaLS in water calorimetrically at different centration curve. The interstitial and substitutional temperatures and obtained -134 f 10 and -123 f 2 cal dissolution effects of amphiphilic additives is also nodeg-’ mol-l respectively, for ACm,. However, the increasing ticeable in the solubility behavior of nonpolar gases in (becoming more positive) value of ACm, with increasing aqueous solutions. A t low temperatures the solubility of BU concentration is opposite to that f o ~ n dfor ~ pro~ > ~ ~ the nonpolar gas (argon, oxygen, methane, or ethane) panols where ACmpbecomes more negative. increases with increasing Concentration of the amphiphilic additive (methanol or ethanol), reaches a maximum and Discussion then decreases to a minimum before it starts rising again. Micelle formation is related to the dual nature of the The initial increase in solubility of the gas most probably surfactant molecule (ion) and the ability of one part, the is due to the extra coaccommodation of the nonpolar hydrophilic, to interact strongly with water and the inamolecules in the cavity partly occupied by the nonpolar bility of the other part, the hydrophobic, to mix regularly groups of the amphiphilic additive. However, as the with the water molecules because of strong water-water concentration of the additive increases the structure of interactions. These two conflicting tendencies in the water is disturbed so that interstitial dissolution decreases surfactant ion are satisfied in bulk liquid water prior to and substitutional dissolution starts to become important micelle formation if the hydrophobic part dissolves mainly and, at the minimum, they balance each other; thereafter interstitially and the hydrophilic part substitutionally.8 substitutional dissolution becomes more important. The Liquid water because of three-dimensional hydrogen existence of a minimum in the solubility curves of many bonding has a relatively open structure allowing the examphiphilic additives in water47is also a reflection of the istence of dynamic interstices or cavities of certain sizes interstitial substitution dissolution behavior. which can be used to accommodate nonpolar molecules The increase in CY has been explained in terms of the (moities) and allows the formation of interstices by the reduction of micellar charge density resulting from the nonpolar molecules (moities) themselves on dissolution. separation of the ionic heads by the hydrophilic (nonionic) The latter type requires energy40 which is partly comgroup of the additive.33 However, our results and those pensated by the interaction of the nonpolar group with the of 0 t h e 1 - s show ~ ~ ~ ~that ~ a increases, reaches a maximum, surrounding water molecules forming the cavity. This and then decreases as the concentration of the amphiphilic interaction is in addition to the hydrogen bonding between additive increases. No explanation was offered for the those water molecules surrounding the hydrophobic group, decrease in a at high additive (alcohol) concentrations. hence the so-called stabilization of water structure by This decrease in a could be explained by the micelles nonpolar groups. Thus stabilization of water structure by becoming progressively richer in the surfactant ion at high nonpolar groups is related to interstitial dissolution which additive concentration so the surface charge density inis central to the hydrophobic effect. Destabilization of creases. In a solution where a micelle is formed from the structure by amphiphilic (polar-nonpolar) additives, on aggregation of two amphiphiles of differing hydrophobic the other hand, is connected with the decrease of incharacters, the weaker one would prefer the bulk solution terstitial and increase of substitutional (regular) dissomore, so less and less of it (the additive) would coaggregate lution. with the surfactant ions until eventually with increasing At a given temperature for a given surfactant only a additive concentration the solution becomes so destruccertain amount of monomers can be accommodated in (a tured that even the surfactant ions loose the tendency to given volume of) solution and any further dissolution of form micelles. the surfactant will be essentially transformed to micelles. The changes made by the additive and by the temThe cmc in pure water is related to this limit. When an perature on the structure of the solvent are reflected in amphiphile such as butylurea is dissolved in water it will and ASo, the enthalpy and entropy changes. Both Mom influence the hydrophobic effect in two ways, one occurring
The Journal of Physlcal Chemistry, Vol. 83, No. 17, 1979 2243
Cmc of NaLS in n-Butylurea Solutions
(as well as AH, and A§,) become less positive with increasing additive concentration corresponding to the progressive decrease in the amount of structured water around the hydrophobic groups as the interstitial solution of nonpolar moities becomes less and substitutional solution increases. This trend continues until a minimum is reached where both types of solution occur to the same extent, after which regular mixing becomes dominant and AH”, and Aso, tend to rise again, however, micelles start to become unstable and eventually would not form. Such trends in AHo, and AS”, were also noted in presence of alcohols.48 We are confining our discussion of the thermodynamic effects of BU addition on the micellization of NaLS to the monomeric ions only. The solvent structural changes are revealed also in ACm, values. The high negative value of ACm, in the absence of the additive is indicative of the loss of structured water around the monomeric surfactant ion in bulk solution on micellization. This loss becomes progressively less as the additive affects the structure of the solvent so less structured water form around the surfactant ion in bulk solution and hence A P , becomes less negative. This trend we obtain is different from that reported in the literature28 for propanols where ACm became more negative in presence of the additives, however, ethanol made A P , less negative29 and it was suggested that the difference between the effect of propanols and ethanol on A P , might be related to penetration of the micelles by propanols but not by ethanol. This does not seem likely since both ethanol and propanol are amphiphiles and are expected to be incorporated in the micelle but to different extents. The extent of penetration may be deduced from the variation of the slope of the conductance curve above the cmc with the (nonionic) additive concentration. The slope above the cmc is proportional to the degree of ionization to a first approximationz2 since the additive does not influence the equivalent conductance of the counterions significantly. Likewise we expect the micelle mobility to be influenced mainly by the degree of dissociation of the micelle. The increase in CY and the corresponding increase in the slope above the cmc result from the decrease in the micelle surface charge density due to the separation of the ionic heads by the (nonionic) hydrophilic head of the amphiphile additive. Thus additives which are unlikely to penetrate the micelle would not increase the slope above the cmc significantly. Such an additive is urea which is polar but not amphiphilic. Urea did not show any variation in the slope,above the cmc in the system NaLS-H20 up to 2 M, the highest concentration used. However, in the presence of ethanol,49n-butylurea, l - p r ~ p a n o land ,~~ l-butano150 the initial rates of the variation of the slope above the cmc with the additive concentration are approximately in the ratio 1:50:100:400, showing butylurea between ethanol and 1-propanol. This appears reasonable as the hydrophilic group in BU is much larger than the hydroxyl group. Another measure of penetration which might also be satisfactory is the initial slope of the cmc-additive concentration curve. For the same system we have for ethanol, butylurea, propanol, and butanol, the variations are approximately in the ratio 1:11:27:114, showing again BU with an intermediate hydrophobic character between ethanol and 1-propanol. The strong adsorption of BU at the water-dodecane interface as well as the result of the partition coefficient measurements of BU between water and dodecane strongly favor the idea of amphiphilic additives penetrating the micelle. Supplementary Material Available: Table I contains the coefficients a and b for the lines h = a b below and
+
above the cmc, which were obtained by the least-squares method, for each solution used in this study (2 pages). Ordering information is available on any current masthead page-
References and Notes M. F. Emerson and A. Holtzer, J. Phys. Chem., 71, 3320 (1967). W. Kauzmann, Adv. Protein Chem., 14, 1 (1959). C. Tanford, J . Phys. Chem., 78, 2469 (1974). G. Stainsby and A. E. Alexander, Trans. Faraday Soc., 48,587 (1950). (5) W. Brunning and A. Holtzer, J. Am. Chem. Soc., 83, 4865 (1961). (6) P. Mukerjee and A. Ray, J. Phys. Chem., 67, 190 (1963). (7) M. J. Schick, J . Phys. Chem., 68, 3585 (1964). (8) M. Abu-Hamdiyyah, J . Phys. Chem., 6% 2720 (1965). (9) D. W. James, R. F. Armishaw, and R. L. Frost, J . Phys. Chem., 80, 1346 (1976). 110) . , A. I. Sodorova, I.N. Kochnev. L. V. Moiseeva, and A. I. Khaloimov in “Water in Biological Systems”, Consultants Bureau, New York, 1968. or 54. (11) J. C. MacDonald, J. Serphillips, and J. J. Guerrera, J. Phys. Chem., 77, 370 (1973). (12) G. C. Kresheck and L. Benlamine, J. Phvs. Chem., 68 2467 (1964). (13) M. Abu-hamdiyyah and A. .Shehabuddin: manuscrlpt In preparation. (14) J. A. Gordon and W. P. Jencks, Biochemistry, 2, 47 (1963). (15) T. T. Herskovits, H. Jaillet, and B. Gadegbeku, J. Biol. Chem., 245, 4544 (1970). (16) G. Barone, V. Crescenzi, A. M. Liquori, and F. Quadrifoglio, J . Phys. Chem., 71, 984 (1967). (17) L. Levine, J. A. Gordon, and W. P. Jercks, Bk.xhemistty, 2, 168 (1963). (18) P. Mukerjee and K. J. Mysels, Natl. Stand. Ref. Data Ser., Natl. Bur. Stand., No. 36 (1971). (19) See paragraph at the end of text regarding supplementary material. (20) E. D. Goddard and G. C. Benson, Can. J . Chem., 35, 966 (1957). (21) H. C. Evans, J . Chem. Soc., 585 (1956). (22) P. Mukerjee, K. J. Mysel, and P. Kapuar, J . Phys. Chem., 71, 4166 (1967). (23) M. Abu-Hamdiyyah and K. J. Mysels, J. Phys. Chem., 71, 418 (1968). (24) P. H. Elworthy and K. J. Mysels, J . Colloid Interface Sci., 21 331 (1966). (25) E. W. Anacker, “Micelle Formation of Cationic Surfactants in Aqueous Media” in “Cationic Surfactants”, M. Schick, Ed., Marcel Dekker, New York, 1970. (26) K. J. Mysels, ”Introduction of Colloid Chemistry”, Intersclence, New York, 1959. (27) P. F. Mijnlief, J . Colloid Interface Sci., 33, 255 (1970). (28) B. D. Flockhart and A. R. Ubbelohde, J. ColloidSci., 8, 428 (1953). (29) G. C. Krescheck and W. A. Hargraves, J. ColloM Interface Sci., 48, 481 (1974). (30) K. Shinoda, J. Nakagawa, B. Tamamushi, and I.Isemura, “Colloidal Surfactants”, Academic Press, New York, 1963. (31) Handbodc of Chemistry and Physlcs, Chemical Rubber Publlshing Co., Cleveland, Ohio. (32) G. C. Kresceck, “Surfactants” in “Water. A ComprehensiveTreatise”, Vol. 4, F. Franks, Ed., Plenum Press, New York, 1975. (33) J. W. Larsen and L. B. Tepky, J. CobMInterface Sci., 49, 113 (1974). (34) S. Mlyagishi, Bull. Chem. Soc. Jpn., 48, 2349 (1975). (35) H. F. Huisman, Kronikl. Ned. Akad., Weternschap., Proc., Ser. 6 , 64, 367 119641. (36) Y. Moroi, ‘N. Nikikdo, H. Vehara, and R. Matuura, J. Colloid Interface Sci., 50, 254 (1975). (37) c. Jollicoeur and P. R. Philip, Can. J . Chem., 52, 1834 11974). (38) G. M. Musbally, G. Perron, and J. E. Desnoyers, J. Colloa I&terface Sci.. 48. 494 (19741. (39) G.M. Musbaily; G. Perron, and J. E. Desnoyers, J . ColloM Interface Sci., 54, 86 (1976). (40) Formation of cavities by nonpolar molecules is favored by an increase in temperature (the work of creating a cavity becomes less) which, however, would tend to destroy some of the existing cavities because of the increased motion. This is most likely the reason for the initial decreasing solubility of nonpolar gases in water, reaching a minimum and then increaslng with increasing temperature.5’ (41) At a higher temperature but below the Krafft point the solubility sometimes increases arid this mlght be due to the co-accommodation of the surfactant hydrophobic chains in the cavities partly occupied with the nonpolar groups of the additl~e.~‘ (42) K. Shirahama, M. Hayashi, and R. Matuura, Bull. Chem. Soc. Jpn., 42, 1206 (1969). (43) A. Metzer and I. J. Lin, J. Phys. Chem., 75, 3000 (1971). (44) N. 2. Kostova and Z. N. Markina, Kolloid Zh., 33, 551 (1971). (45) N. 2 . Kostova, 2. N. Markina, P. A. Rebinder, and A. E. Kuzimina KolloidZh., 33, 79 (1971). (46) M. Yaacobi and A. Ben-Naim, J . Solution Chem., 2, 426 (1973). (47) N. Nichino and M. Nakomura, Bull. Chem. Soc. Jpn., 51, 1617 (1978). (48) L. Benjamine, J . Colloid Interface Sci., 22, 386 (1966). (49) A. F. Ward, Proc. R . Soc. London, Ser. A , 176, 412 (1940). (50) K. Shirahama and T. kashiwabara, J. Colloid Interface Sci., 36, 65 (1971). (51) R. Battino and H. L. Clever, Chem. Rev., 66, 395 (1966). (1) (2) (3) (4)
- - - I
