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Articles Microcalorimetric Studies on the Thermodynamic Properties of Cationic Gemini Surfactants Guangyue Bai,† Haike Yan,*,† and Robert K. Thomas‡ Center for Molecular Science, Institute of Chemistry, Chinese Academy of Sciences, Beijing 100080, People’s Republic of China, and Physical and Theoretical Chemistry Laboratory, Oxford University, South Parks Road, Oxford OX1 3QZ, U.K. Received October 20, 2000. In Final Form: April 10, 2001 Microcalorimetric measurements have been made as a function of temperature on the series of gemini surfactants [CMH2M+1(CH3)2N(CH2)SN(CH3)2CMH2M+1]Br2, designated CMCSCM, where M and S indicate the numbers of carbon atoms in the side chains and spacer, respectively, for M ) 12 and S ) 3, 4, 6, 8, 10, and 12. The maximum at S ) 4-6 in the values of the critical micelle concentrations (cmc’s) as a function of spacer length was found to be associated with a strong minimum in the magnitude of the exothermic enthalpy of micellization, ∆Hmic. A strong maximum in the corresponding entropies of micellization, ∆Smic, almost cancels out the effect of ∆Hmic, so that ∆Gmic and hence the values of the cmc show a much weaker variation with spacer. The values of ∆CP pass through a marked minimum and are large, corresponding to a large variation of ∆Hmic with temperature. This variation of ∆Hmic has some interesting consequences for the variation of the cmc. Thus, for C12C12C12 the contribution of ∆Hmic to ∆Gmic at room temperature is about 36%, but this increases through 61% to 84% at 308 K. Over this range the changes in ∆Gmic remain within about 10% of each other, almost within the experimental error.
Introduction Gemini surfactants have a number of properties different from those of their single-chain counterparts, for example, lower critical micelle concentrations (cmc’s), better wetting properties, lower limiting surface tensions, unusual aggregation morphologies, etc.1-9 We have shown in a previous paper10 that the enthalpies and entropies of micellization of the cationic gemini and related doublechain surfactants at 303.15 K showed a striking pattern which is completely concealed by the behavior of the Gibbs energy of micellization (from cmc data). This greater detail in the thermodynamic information may prove helpful in understanding the factors behind the differences between gemini and other surfactants. In the previous paper we only reported measurements at a single temperature. Here, we extend them to other temperatures, thus obtaining heat capacity data as well as the variation of enthalpy and entropy of micellization with temperature. * To whom correspondence should be addressed. E-mail: hkyan@ public.east.cn.net. Fax: 86-010-62559373. Phone: 86-010-62562821. † Chinese Academy of Sciences. ‡ Oxford University. (1) Menger, F. M.; Littau, C. A. J. Am. Chem. Soc. 1991, 113, 14511452. (2) Rosen, M. J. CHEMTECH 1993, 23, 30-33. (3) Menger, F. M.; Littau, C. A. J. Am. Chem. Soc. 1993, 115, 1008310090. (4) Zana, R.; Talmon, Y. Nature 1993, 362, 228-230. (5) Karaborni, S.; Esselink, K.; Hilbers, P. A. J.; Smit, B.; Kartha¨user, J.; van OS, N. M.; Zana, R. Science 1994, 266, 254-256. (6) Bhattacharya, S.; De, S. J. Chem. Soc., Chem. Commun. 1995, 651-652. (7) Zana, R. Curr. Opin. Colloid Interface Sci. 1996, 1, 566-571. (8) Kim, S. S.; Zhang, W. Z.; Pinnavaia, T. J. Science 1998, 282, 1302-1305. (9) Regev, O.; Zana, R. J. Colloid Interface Sci. 1999, 210, 8-17. (10) Bai G. Y.; Wang, J. B.; Yan, H. K.; Li, Z. X.; Thomas, R. K. J. Phys. Chem. B 2001, 105, 3105-3108.
There have been a few measurements of changes in the heat capacity on micellization.11-15 The geminis investigated were the N,N′-bis(dimethyldodecyl)-R,ω-alkanediyldiammonium dibromide compounds with the general structure
[CMH2M+1(CH3)2N(CH2)SN(CH3)2CMH2M+1]Br2 for which we use the simpler notation CMCSCM, where M and S stand for the numbers of carbon atoms in the free alkyl chains and the spacer, respectively. Many properties of these surfactants have previously been determined, for example, the cmc, degree of ionization of the micelle, morphology of the micellar aggregate, phase behavior, rheology, and surface area per molecule adsorbed at the air/water interface.16-25 For our purposes the most in(11) Bashford, M. T.; Woolley, E. M. J. Phys. Chem. 1985, 89, 31733179. (12) Kresheck, G. C. J. Am. Chem. Soc. 1998, 120, 10964-10969. (13) Andersson B.; Olofsson, G. J. Chem. Soc., Faraday Trans. 1 1988, 84 (11), 4087-4095. (14) Muller, Norbert Langmuir 1993, 9, 96. (15) De Lisi, R.; Fisicaro, E.; Milioto, S. J. Solution Chem. 1988, 17 (11), 1015-1041. (16) Camesano, T. A.; Nagarajan, R. Colloids Surf., A 2000, 167, 165-177. (17) Devı´nsky, F.; Lacko, I.; Imam, T. J. Colloid Interface Sci. 1991, 143, 336-342. (18) Zana, R.; Benrraou, M.; Rueff, R. Langmuir 1991, 7, 10721075. (19) Alami, E.; Levy, H.; Marie, P.; Zana, R. Langmuir 1993, 9, 940944. (20) Alami, E.; Beinert, G.; Marie, P.; Zana, R. Langmuir 1993, 9, 1465-1467. (21) Frindi, M.; Michels, B. Langmuir 1994, 10, 1140-1145. (22) Danino, D.; Talmon, Y.; Zana, R. Langmuir 1995, 11, 14481456. (23) Hirata, H.; Hattori, N.; Ishida, M.; Okabayashi, H.; Frusaka, M.; Zana, R. J. Phys. Chem. 1995, 99, 17778-17784.
10.1021/la001472u CCC: $20.00 © 2001 American Chemical Society Published on Web 06/19/2001
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Figure 1. A typical curve of the output of the microcalorimeter from the experimental measurement of C12C3C12.
teresting observation is that, for a given value of M, the cmc of a CMCSCM series passes through a maximum at S ) 5, but the surface excess at the air/water interface is not a maximum until S ) 10-12, for which a theoretical model has been given by Diamant and Andelman.26 Zana27,28 has also postulated a relationship between the Gibbs energy of micellization and the Gibbs energy of transfer of an alkyl chain from water to the hydrophobic core [∆Gmic(chain)] and used it to explain the pattern of cmc values of the geminis compared with their monomeric counterparts. Experimental Section The C12CSC12 series of surfactants were synthesized and purified according to the method of Menger.3 All the surfactant solutions were prepared using doubly distilled water. An improved LKB-2107 isothermal titration microcalorimeter with a 1 mL sample cell was used for the measurements. It has a precision of electrical calibration better than (1%, and its accuracy was tested by measuring the dilution enthalpy (-0.643 ( 0.015 kJ/mol) of a concentrated sucrose solution, giving results in good agreement with the literature values (-0.653 kJ/ mol)29. For the experiments on the surfactants, the sample cell and the reference cell of the calorimeter were initially loaded with 0.5 and 0.7 mL of pure water, respectively. A concentrated solution of surfactant was injected into the stirred sample cell via a 500 µL Hamilton syringe controlled by a Braun 871182 pump. A series of injections, each of 10-20 µL, was made until the desired range of dilution had been covered. The experiments were performed at temperatures of 298.15, 303.15, and 308.15 ( 0.02 K.
Results A typical plot of the output of the microcalorimeter is shown in Figure 1. All the present measured heats were endothermic because the experimental technique involved dilution of concentrated solutions of micelles. Therefore, the enthalpies of micelle formation are negative. On the plot a clear break corresponding to micelle formation was observed, allowing identification of the cmc. The enthalpies of micellization were obtained from the typical curves as shown in Figure 2 for all three temperatures for all the studied geminis. When the overall concentration is in the premicellar range, the added micelles from the concentrated solution dissociate into monomers and the mono(24) Li, Z. X.; Dong, C. C.; Thomas, R. K. Langmuir 1999, 15, 43924396. (25) Fielden, M. L.; Claesson, P. M.; Verrall, R. E. Langmuir 1999, 15, 3924-3934. (26) Diamant, H.; Andelman, D. Langmuir 1995, 11, 3605-3606. (27) Zana, R.; Levy, H.; Papoutsi, D.; Beinert, G. Langmuir 1995, 11, 3694-3698. (28) Zana, R. Langmuir 1996, 12, 1208-1211. (29) Gucker, F. T., Jr.; P. H. B.; Planck, R. W. J. Am. Chem. Soc. 1939, 61, 459.
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Figure 2. Variation of the observed enthalpies (∆Hobs) with concentration (C) of C12C3C12 (initial concentration 5 mM) in the microcalorimetry dilution experiment at 298.15, 303.15, and 308.15 K. Table 1. Cmc Values (mol/L) for the C12CSC12 Series in Water at 298.15, 303.15, and 308.15 K microcalorimetrya (104cmc)
conductivityb (104cmc)
surface tensionc (104cmc)
S
298.15 K
303.15 K
308.15 K
298.15 K
298.15 K
3 4 6 8 10 12
10.9 ( 0.2 11.4 ( 0.2 8.9 ( 0.2 7.1 ( 0.1 3.9 ( 0.1 2.6 ( 0.1
8.4 ( 0.3 10.1 ( 0.1 8.3 ( 0.2 6.8 ( 0.2 2.8 ( 0.3 2.1 ( 0.3
8.4 ( 0.1 9.5 ( 0.2 8.3 ( 0.3 6.6 ( 0.3 2.2 ( 0.2 1.6 ( 0.2
9.6 ( 0.3 11.7 ( 0.4 10.3 ( 0.4 8.3 ( 0.3 6.3 ( 0.3 3.7 ( 0.1
9.1 10.0 11.2 8.9 3.2 2.8
a These are our experimentally determined values. b These are values reported in ref 18. c These are values reported in ref 20.
Table 2. Enthalpies and Heat Capacities of Micellization for the C12CSC12 Series in Water at 298.15, 303.15, and 308.15 K S
298.15 K
∆Hmic (kJ/mol) 303.15 Ka
308.15 K
3 -5.70 ( 0.11 -15.69 ( 0.47 -24.85 ( 0.75 4 -4.15 ( 0.10 -9.38 ( 0.28 -14.55 ( 0.44 6 -3.67 ( 0.07 -8.85 ( 0.25 -11.99 ( 0.36 8 -4.80 ( 0.09 -9.53 ( 0.33 -13.87 ( 0.44 10 -6.74 ( 0.14 -13.95 ( 0.55 -20.82 ( 0.69 12 -10.10 ( 0.15 -18.03 ( 0.64 -26.04 ( 0.92 a
∆CP (kJ/(K mol)) -1.92 ( 0.15 -1.04 ( 0.08 -0.83 ( 0.07 -0.91 ( 0.08 -1.41 ( 0.13 -1.59 ( 0.14
From ref 10.
mers are diluted. However, when the overall concentration reaches the cmc and above, the added micelles merely become diluted without any dissociation. The enthalpy of micellization, ∆Hmic, is obtained from the difference between the observed enthalpies of the two linear segments of the plot. ∆Gmic is calculated using an expression in the literature,28 and the entropy of micellization, ∆Smic, can then be derived using the values of ∆Hmic and ∆Gmic. The mean value of ∆CP over the whole temperature range was derived from the variation of ∆Hmic with temperature. The complete set of results including the values at 303.15 K from the previous paper is given in Tables 1-3. The values of the cmc and their variation with S for the C12CSC12 series agree well with those obtained by noncalorimetric methods.18,20 The values of ∆Hmic as a function of spacer length S are shown for the three different temperatures in Figure 3. Values of ∆Hmic for two gemini compounds at 308.15 K have been previously reported by Grosmaire et al.30 and were -12 kJ mol-1 for C12C6C12 and -15 kJ mol-1 for C12C10C12. Our corresponding values were -11.99 ( 0.36 and -20.82 ( 0.69 kJ mol-1, (30) Grosmaire, L.; Chorro, M.; Chorro, C.; Partyka, S.; Zana, R. Prog. Colloid Polym Sci. 2000, 115, 31-35.
Thermodynamic Properties of Gemini Surfactants
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Table 3. Gibbs Energies and Entropies of Micellization for the C12CSC12 Series in Water at 298.15, 303.15, and 308.15 K -T∆Smic (kJ/mol)b
∆Gmic (kJ/mol)a
S 298.15 K 303.15 K 308.15 K 298.15 K 303.15 K 308.15 K 3 4 6 8 10 12
-30.1 -29.2 -29.1 -27.9 -28.4 -28.3
-31.8 -30.2 -29.8 -28.5 -30.1 -29.4
-32.3 -31.0 -30.3 -29.1 -31.5 -30.9
-24.4 -25.1 -25.4 -23.1 -21.7 -18.2
-16.7 -20.9 -21.0 -19.0 -16.2 -11.4
-7.5 -16.5 -18.4 -15.2 -10.7 -4.9
a Calculated from ref 28 using the equation ∆G mic ) (1 + β)RT ln(cmc); β can be taken from ref 18. b Calculated from the equation ∆Gmic ) ∆Hmic - T∆Smic.
Figure 3. Variation of the enthalpy of micellization (∆Hmic) for C12CSC12 with the spacer length (S) at 298.15, 303.15, and 308.15 K.
respectively. The data in Table 2 show that the temperature sensitivity of the enthalpies is quite large, changing about 1 kJ mol-1 K-1. Grosmaire et al. did not give any indications of the precision of either their temperature of measurement or the values of ∆Hmic, so we cannot make any further comments on the discrepancies except to note that the pattern of behavior they observed for the variation of ∆Hmic for the three geminis they studied is entirely in agreement with ours. Discussion The pattern of variation of ∆Hmic for a single-chain surfactant series is exemplified by the results of Woolley et al.,11 and Mosquera et al.31 for the alkyltrimethylammonium bromides (CNTAB). The enthalpies of micelle formation change linearly from a small positive value for N e 10, becoming increasingly negative with the number of methylene groups (N). The corresponding entropies of micelle formation are all positive and also increase with N. Thus, entropy must be the driving force behind micellization for N e 10, but at higher values of N micellization is also assisted by an increasingly negative enthalpy. Woolley et al. also determined values of ∆CP but obtained a rather confusing pattern of behavior. At high temperatures they found that ∆CP was negative and on the order of 1 kJ mol-1 K-1. However, at low temperatures ∆CP was sometimes positive and seemed to follow a more unusual pattern at lower temperatures (10 °C). For the gemini surfactants there is a minimum in the magnitude of the exothermic enthalpy of micellization at S ) 4-6 at all three temperatures. This minimum becomes more pronounced as the temperature increases above room (31) Mosquera, V.; Manuel del Rı´o, J.; Attwood, D.; Garcı´a, M.; Jones, M. N.; Prieto, G.; Suarez, M. J.; Sarmiento, F. J. Colloid Interface Sci. 1998, 206, 66-76.
Figure 4. Variation of the heat capacity of micellization for C12CSC12 with the spacer length (S). The solid line is a polynomial fitting line.
temperature. Following the discussion of Diamant and Andelman,26 there are four contributions to the thermodynamic functions that need to be considered: the van der Waals interaction between the chains, the headgroup repulsion, the hydrophobic interaction, and the energetics of the configuration of the spacer chain. The van der Waals and hydrophobic interactions will always tend to make ∆Hmic negative. These would be expected to increase with the total number of carbon atoms in the surfactant. Thomas and his collaborator reported the results from the neutron scattering, which showed the known actual areas per molecule of this series of gemini surfactants at the cmc increase steadily with S,24 so the headgroup repulsion must also decrease with S. This would also make ∆Hmic more negative. The only contribution that will have the opposite effect on ∆Hmic is any energy that is required to distort the spacer chain. However, this would not seem to be an effect that could contribute adequately at low values of S, and hence it is difficult to see how it could explain the marked exothermicity of ∆Hmic for S ) 3. Comparison with the double-chain singly charged surfactants in the previous paper suggested that short side chains (or spacers) could not take advantage of either the hydrophobic or the van der Waals interactions.10 This would mean that the addition of the first few methylene groups to the spacer would not contribute at all to the exothermicity. Under these circumstances the energy required to configure the spacer chain, which is endothermic, can make a relatively large contribution to the overall ∆Hmic. Once the spacer is longer than six methylene groups, it can assume configurations where it is gaining hydrophobic energy as well as some extra van der Waals interaction energy with the other chains. Hence, ∆Hmic starts to become more negative again. This explanation, based mainly on the parallel behavior of the double-chain singly charged surfactants described in the previous paper, is similar to that given by Grosmaire et al.30 and is also consistent with the ideas presented by Diamant and Andelman.26 There is a weak maximum in the value of the cmc at about S ) 4-6 at all three temperatures, which for two of the temperatures corresponds to a deep minimum in the magnitude of ∆Hmic. Meanwhile there are corresponding large maxima in the values of ∆Smic as shown in Table 3. In combining to give ∆Gmic, and hence the cmc, the large variations in ∆Hmic and ∆Smic compensate each other, and the net variation in cmc is comparatively small. This shows how important it is to study the thermodynamic components of ∆Gmic rather than rely only on the values of the cmc. The values of ∆CP also pass through a marked minimum (plotted as the negative values in Figure 4). The magnitude
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of ∆CP is always large because the variation of ∆Hmic is large. This variation of ∆Hmic has some interesting consequences for the variation of the cmc. Taking C12C12C12 as an example, at room temperature the contribution of ∆Hmic to ∆Gmic is about 36%, but this increases through 61% to 84% at 308 K. However, over this range the changes in ∆Gmic remain within about 10% of each other, almost within the experimental error. Once again, although the values of the cmc do not indicate any interesting physicochemical changes, the more detailed thermodynamic information reveals a much more complex behavior. We do not attempt to give any interpretation of the actual
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values of ∆CP at this stage. There have been few such measurements and little or no interpretation to date. We will extend this type of measurement to other series of surfactants to determine whether there are any welldefined patterns in the ∆CP behavior. Acknowledgment. We are grateful for financial support from the Royal Society, the Chinese Academy of Sciences, and the National Natural Science Foundation of China (Grant 20073055). LA001472U
