Thermodynamic Study of the Micellization of Zwitterionic Surfactants

May 13, 2014 - Thermodynamic Study of the Micellization of Zwitterionic Surfactants and Their Interaction with Polymers in Water by Isothermal Titrati...
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Thermodynamic Study of the Micellization of Zwitterionic Surfactants and Their Interaction with Polymers in Water by Isothermal Titration Calorimetry César Brinatti, Laura Bissoli Mello, and Watson Loh* Institute of Chemistry, Universidade Estadual de Campinas (UNICAMP), CP 6154, CEP 13083-970 Campinas, SP Brazil S Supporting Information *

ABSTRACT: The micellization of a homologous series of zwitterionic surfactants, a group of sulfobetaines, was studied using isothermal titration calorimetry (ITC) in the temperature range from 15 to 65 °C. The increase in both temperature and the alkyl chain length leads to more negative values of ΔG0mic, favoring the micellization. The entropic term (ΔS0mic) is predominant at lower temperatures, and above ca. 55−65 °C, the enthalpic term (ΔH0mic) becomes prevalent, figuring a jointly driven process as the temperature increases. The interaction of these sulfobetaines with different polymers was also studied by ITC. Among the polymers studied, only two induced the formation of micellar aggregates at lower surfactant concentration: poly(acrylic acid), PAA, probably due to the formation of hydrogen bonds between the carboxylic group of the polymer and the sulfonate group of the surfactant, and poly(sodium 4-styrenesulfonate), PSS, probably due to the incorporation of the hydrophobic styrene group into the micelles. The prevalence of the hydrophobic and not the electrostatic contributions to the interaction between sulfobetaine and PSS was confirmed by an increased interaction enthalpy in the presence of electrolytes (NaCl) and by the observation of a significant temperature dependence, the latter consistent with the proposed removal of hydrophobic groups from water.

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approaches for ionic10 and nonionic surfactants,3 and the enthalpy of micellization, ΔH0mic. Accordingly, the entropy of micellization, ΔS0mic, is then obtained by difference. It is widely known that ionic surfactants interact with ionic11−14 and nonionic15−23 polymers starting at concentrations below the cmc at a given surfactant concentration called the critical aggregation concentration, cac, forming aggregates with the polymer chains in a cooperative way. In calorimetric experiments, this is seen by an abrupt change in the enthalpy values at low surfactant concentration.9,13,14,17,18,20,21,23 This type of micellar aggregate formation increases with increases in the surfactant concentration, where a second critical concentration is achieved, as denoted by C2. This is where the polymer no longer influences the micellization and the concentration of free unimers in solution reaches the cmc and therefore unbound micelles start to form in solution. There are only a few studies in the literature regarding the interaction between zwitterionic surfactants and polymers, and none of them using isothermal titration calorimetry. Velázquez et al.24 investigated the interaction between sulfobetaine SB 312 (dodecyl chain and a spacer with three carbon atoms) and poly(sodium 4-styrenesulfonate), PSS, by tensiometry and ascribed a cac value at low surfactant concentration. In their

INTRODUCTION Zwitterionic surfactants are a class of surfactants that possess both positive and negative charges in their polar headgroups. Usually, the positive charge is an ammonium group, while there are different types of groups responsible for the negative charge, such as carboxylates, sulfates, and sulfonates.1,2 The presence of both charges on the same molecule renders an intermediate hydrophilicity between ionic and nonionic surfactants,3 and this feature can be changed by an increase or decrease in the length of the spacer group,3,4 known as the tether group. This kind of surfactant is preferably used in cosmetics because of its mildness to the eyes and skin.1−3 These surfactants behave similarly to the other classes, and a simplified model of a zwitterionic micelle was proposed by Baptista et al.,5 where the micelles are represented as concentric spheres resembling a capacitor, therefore presenting an ionic dipolar feature. This model explains the interactions between zwitterionic micelles and ions by electrostatic interactions but fails to explain the ionbinding specificity regarding cations and anions following the Hofmeister’s series. Such a feature was studied by the group of Nome6−8 along with the catalytic reaction properties of zwitterionic surfactants due to the chameleon nature of zwitterionic micelles. Amidst the many techniques used to follow self-assembly, isothermal titration calorimetry (ITC) has been showing growing use.9 It allows one, in a single experiment, to obtain thermodynamic parameters such as the cmc and hence the Gibbs free energy of micellization, ΔG0mic, using different © 2014 American Chemical Society

Received: September 13, 2013 Revised: April 25, 2014 Published: May 13, 2014 6002

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Krafft point being near 30 °C. All of the experiments were performed at least in duplicate. All data were obtained and treated with Origin 7.0 software. All of the thermodynamic parameters were obtained with the procedures detailed by Olofsson et al.9 The relationship between the cmc and the Gibbs free energy of micellization (ΔG0mic) is shown in eq 1, as used for zwitterionic surfactants.3

study of fluorescence using Nile red as a probe,25 now with surfactant 3-[(3-cholamidopropyl) dimethylammonium]-1-propanesulfonate, CHAPS, and polyelectrolytes PSS and poly(diallyldimethylammonium chloride), PDADMAC, the same kind of interaction was reported. Bakshi et al.26 used a fluorescence technique, and Mahajan et al.27 used cyclic voltametry technique to study the interaction between a homologous series of sulfobetaines and triblock copolymers PEO-PPO-PEO with different molecular weights. Both of them evaluated the synergism of the mixture as well as the hydrophobic contributions of the surfactants and the polymers. With this background, the aim of the present study is to use isothermal titration calorimetry to follow both the self-assembly of a homologous series of sulfobetaines and their interactions with nonionic and ionic polymers. For the self-assembly study, the temperature ranged from 15 to 65 °C. As a result, thermodynamic parameters such as the Gibbs free energy (ΔG0mic), enthalpy (ΔH0mic), entropy (ΔS0mic), and heat capacity 0 changes (ΔCpmic ) were obtained. Regarding the interaction with polymers, the following polymers were chosen: nonionic poly(ethylene oxide), PEO; poly(propylene oxide), PPO; poly(acrylic acid), PAA; cationic poly(diallyldimethylammonium chloride), PDADMAC; and anionic polymers poly(sodium acrylate), PA; and poly(sodium 4styrenesulfonate), PSS, in mixtures with SB 3-12 and SB 3-14. The literature has enough information from ITC measurements on polymer−surfactant systems, recently reviewed.9,23 We would like to investigate the interaction of the betaine dipolar surfactants with polymers with different features such as electrostatic charges, hydrophilicity, and hydrophobicity in order to compare their interaction process with characteristics determined in previous reports, with special emphasis on the evaluation of their chameleon-like behavior in interacting with both small cations and anions.6−8

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0 ΔGmic = RT ln cmc

(1)

The only difference was SB 3-10, whose titration curves were fitted using the following equation29−31 due to its distorted sigmoidal shape (Supporting Information). 0 ΔHdil =

a1c + a 2 + a4C + a5 ⎡c − a ⎤ 1 + exp⎣⎢ d 3 ⎦⎥ x

(2)

where a1−a5 = fitting parameters; C = surfactant concentration; dx = Δcmc = ± 15% cmc value; a2 = ΔH0obs; and a3 = cmc. Dynamic Light Scattering. The experiments were performed on a CGS-3-based compact goniometer system from ALV-GmbH, Langen, Germany, which is equipped with a detection system in pseudocross geometry with a 22 mW He−Ne laser (λ = 632.8 nm) and an ALV 7004 multitau correlator with a measurement angle of θ = 150°. cis-Decalin was used as the refractive index matching liquid, and the temperature was maintained at 25.00 ± 0.01 °C. The experiments were performed at least in duplicate. All data were treated with ALVcorrelator software version 3.0.

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RESULTS Micellization Thermodynamics of Sulfobetaines. Figure 1 shows a typical set of titration calorimetric curves for the micellization, in this case, for sulfobetaine SB 3-14, in the temperature range from 15 to 65 °C.

EXPERIMENTAL SECTION

Chemicals. The homologous series of zwitterionic surfactants (nalkyl-N,N-dimethyl-3-ammonio-1-propanesulfonate), with even carbon numbers, from decyl to hexadecyl homologues was purchased from Anatrace with the highest purity available, and they were used as received. In the text, the acronym SB 3-X will be used to refers to the sulfobetaines, 3 is the number of carbon atoms present at the tether group (propanesulfonate group), and X refers to the number of carbon atoms in the alkyl chain. All of the polymers were purchased from Sigma-Aldrich and used without purification. The polymers are the nonionic poly(ethylene oxide), PEO, M = 100 000 g mol−1; poly(propylene oxide), PPO, M = 1000 g mol−1; poly(acrylic acid), PAA, M = 2000 g mol−1; cationic poly(diallyldimethylammonium chloride), PDADMAC, M = 100 000−200 000 g mol−1, 20 wt % aqueous solution; and anionic poly(sodium 4-styrenesulfonate), PSS, M = 70 000 g mol−1. PAA aqueous solutions display pH 3.5. For some experiments, the pH was adjusted by adding 0.1 mol L−1 NaOH until pH 7.5 and 11.7 to produce sodium poly(acrylate), PA. From previous results of the titration of PAA solutions,28 we estimate that the polymer degrees of ionization (α) at these pH values are 0.06, 0.92, and 1.00, respectively. The water used to prepare all solutions was Milli- Q grade (18.2 MΩ cm−1). Isothermal Titration Calorimetry. The calorimeter used was the MicroCal VP-ITC (Northampton, MA, USA). Small aliquots ranging from 3 to 15 μL were added stepwise by an automatic injection syringe containing 270 μL of a concentrated surfactant solution (typically 10 times above their cmc values) into a reaction cell of 1.43 mL that contains either water or polymeric solution, with a 5 min interval between each injection. For the micellization study of the homologous series, the temperature ranged from 15 to 65 °C for the entire series, except for SB 3-16, which was studied from 35 to 65 °C due to its

Figure 1. Calorimetric titration curves obtained for sulfobetaine SB 314 at different temperatures. The syringe has a solution of 6.0 mmol kg−1 (0.2 wt %).

These ITC curves contain information that can be deconvoluted into the most important thermodynamic features associated with the micellization process as recently reviewed.9,23 The temperature dependence of the cmc is 0 controlled by negative ΔCpmic values, which are an important signature of the hydrophobic effect. The thermodynamic parameters obtained for the whole homologous series are presented in Table 1. 6003

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Table 1. Thermodynamic Parameters Obtained for the Homologous Series in the Studied Temperature Range SB 3-10a T/°C 15 25 35 45 55 65 T/°C 15 25 35 45 55 65 T/°C 15 25 35 45 55 65 T/°C 35 45 55 65 a

cmc/mmol kg 40.1 34.6 34.6 34.7 34.7 37.3

± ± ± ± ± ±

−1

0.2 0.3 0.5 0.3 0.5 0.2

cmc/mmol kg−1 3.6 3.4 3.4 3.5 3.9 4.5

± ± ± ± ± ±

0.1 0.1 0.2 0.1 0.1 0.2

cmc/mmol kg−1 0.37 0.34 0.34 0.41 0.44 0.52

± ± ± ± ± ±

0.01 0.01 0.03 0.01 0.01 0.01

cmc/mmol kg−1 0.057 0.057 0.060 0.068

± ± ± ±

0.003 0.003 0.001 0.006

−1

ΔG0mic/kJ

mol

−7.7 −8.3 −8.6 −8.9 −9.2 −9.2

0.1 0.1 0.1 0.1 0.1 0.1

± ± ± ± ± ±

ΔG0mic /kJ mol−1 −13.5 −14.1 −14.5 −15.0 −15.1 −15.2

± ± ± ± ± ±

0.1 0.1 0.1 0.1 0.1 0.1

ΔG0mic /kJ mol−1 −18.9 −19.8 −20.5 −20.6 −21.1 −21.2

± ± ± ± ± ±

0.1 0.1 0.2 0.1 0.1 0.1

ΔG0mic /kJ mol−1 −25.0 −25.9 −26.5 −27.0

± ± ± ±

0.1 0.1 0.1 0.2

ΔH0mic/kJ mol−1 10.2 6.8 3.0 −0.4 −2.2 −3.8 SB 3-12

± ± ± ± ± ±

0.2 0.2 0.4 0.1 0.1 0.1

ΔH0mic /kJ mol−1 7.6 3.3 −1.0 −4.6 −9.0 −12.8 SB 3-14

± ± ± ± ± ±

0.1 0.1 0.1 0.1 0.3 0.3

ΔH0mic /kJ mol−1 5.7 0.5 −4.8 −9.3 −13.6 −17.4 SB 3-16

± ± ± ± ± ±

0.1 0.1 0.2 0.3 0.4 0.1

ΔH0mic /kJ mol−1 −5.7 −8.4 −10.9 −16.7

± ± ± ±

0.1 0.1 0.6 0.5

TΔS0mic/kJ mol−1 17.9 15.1 11.6 8.5 7.0 5.5

± ± ± ± ± ±

0.2 0.3 0.5 0.1 0.2 0.1

TΔS0mic/kJ mol−1 21.1 17.4 13.5 10.4 6.2 2.3

± ± ± ± ± ±

0.1 0.1 0.1 0.1 0.3 0.3

TΔS0mic /kJ mol−1 24.7 20.3 15.6 11.4 7.5 3.9

± ± ± ± ± ±

0.1 0.1 0.4 0.3 0.4 0.1

TΔS0mic /kJ mol−1 19.4 17.4 15.6 10.3

± ± ± ±

0.2 0.1 0.6 0.7

0 ΔCpmic /J mol−1°C−1

−279 ± 3

0 ΔCpmic /J mol−1°C−1

−409 ± 7

0 ΔCpmic /J mol−1°C−1

−462 ± 1

0 ΔCpmic /J mol−1°C−1

−368 ± 18

See Experimental Section.

magnitude of the enthalpic contribution overweighs the entropic one at 55 °C for SB 3-12 and SB 3-14 and 65 °C for SB 3-16, while for SB 3-10 it does not happen, with the magnitude of the entropic term always being higher than the magnitude of the enthalpic term in the temperature range studied. These results agree well with the well-known phenomenon of enthalpy−entropy compensation.32,33 Such a compensation is only partial because both terms do not vary equally. This compensation is also responsible for the slight variation of the term ΔG0mic within this temperature range. As for the entropic term (TΔS0mic), it is always positive for all surfactants in the temperature range, although increasing temperature causes a decrease in its value due to the decreasing contribution from the hydrophobic effect. Contributions from the Headgroup and the Alkyl Chain to the Thermodynamic Parameters. The thermodynamic functions of micellization for the homologous series were analyzed through the additivity34 scheme represented by

For all surfactants except SB 3-16, increasing the temperature of the experiment causes a decrease in their cmc values until a minimum value is reached, and then the cmc starts to increase again for the other homologous series. Micellization is a cooperative and temperature-dependent process (eq 1), explaining why the values of ΔG0mic become more negative with increases in both the temperature and the alkyl chain length of the surfactant. This temperature effect is also shown in the enthalpy values, which are positive at lower temperatures and become negative at higher temperatures (remembering that the enthalpy changes are related to the derivative of the cmc values with respect to temperature). Moreover, the minimum cmc values correspond to the temperature at which the enthalpy of micellization vanishes (plot in the Supporting Information). For the longer homologous hexadecylbetaine, due to its Krafft point being around 30 °C, only experiments above this temperature were recorded and all processes are exothermic, not allowing the determination of its minimum cmc value. It is possible to note a decrease of approximately 10 °C in the minimum cmc temperature with the increase of two methylene units (−CH2−) on the alkyl chain, which agrees with previous findings.32 At lower temperatures, the micellization process is 0 endothermic (ΔHmic > 0), but with an increase in the temperature of the system, the process becomes exothermic (ΔH0mic < 0) for all surfactants. At higher temperatures, the

0 0 0 ΔX mic = ΔX mic HG + mΔX mic CH 2

(3)

where ΔX0mic = measured thermodynamic property; ΔX0micHG = headgroup contribution for this property; m = number of carbon atoms in the alkyl chain; and ΔX0mic CH2 = contribution of each methylene group for this thermodynamic property. According to this equation, we have chosen to assume the contributions from methyl (−CH3) and methylene (−CH2−) groups as being equal. In general, as seen in Table 2 the 6004

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Table 2. Contributions of the Headgroup (HG) and the Alkyl Chain (CH2) to the Micellization Thermodynamic Properties (kJ mol−1) headgroup (HG) T/°C 15 25 35 45 55 65 a

ΔΔG0mic 20.3 20.3 18.7 19.2 19.7 20.3

± ± ± ± ± ±

0.1 0.1 0.2 0.3 0.1 0.4

ΔΔH0mic 21.2 22.4 17.2 13.1 11.1 15.5

± ± ± ± ± ±

0.8 0.8 0.8 0.1 1.2 0.7

alkyl chain (CH2) ΔTΔS0mic 0.9 2.0 −1.5 −6.1 −8.6 −4.8

± ± ± ± ± ±

0.8 0.7 0.6 0.2 1.1 1.1

ΔΔG0mica −2.8 −2.9 −2.8 −2.8 −2.9 −3.0

ΔΔH0mic −1.1 ± −1.6 ± −1.5a −1.4a −1.5 ± −2.2 ±

0.1 0.1

0.1 0.1

ΔTΔS0mic 1.7 ± 1.3 ± 1.3a 1.4a 1.4 ± 0.8 ±

0.1 0.1

0.1 0.1

The uncertainty is smaller than ±0.1 kJ mol−1.

these curves are essentially the same up to the cmc, with slightly less endothermic values at low surfactant concentration in the presence of polymers. The inflection associated with cmc is also essentially the same, and the curves for PEO and PPO differ from each other only above the cmc region. For PPO, a more hydrophobic polymer, the enthalpy values above the cmc present a difference of 1.5 kJ mol−1 from the dilution curve, while for PEO the curve is superimposed on the dilution curve. The curve with PAA is the one that displays the most distinct behavior. The enthalpy values deviate from the dilution curve at lower surfactant concentration for more negative values of enthalpy, indicating surfactant aggregation at concentrations well below the cmc values, allowing the identification of a cac of 0.10 mmol kg−1. After this deviation, the polymer curve reaches a minimum value and then starts to assume less exothermic values but does not merge with the dilution curve within the studied surfactant concentration range. Previous results from PAA titration28 show that a pH value of 3.5 was chosen to ensure that practically all carboxylic groups are protonated (α = 0.06) and thus are able to form intermolecular hydrogen bonds with other polymer chains and with the sulfobetaine. With this result, a way to analyze whether the contribution of the hydrogen bonds is responsible for the polymer−surfactant complex formed is to increase the pH value from 3.5 to 7.5, where the majority of carboxylic groups are deprotonated (α = 0.92), and up to pH 11.7, where they are all deprotonated (α = 1.0), both referred to as PA. We also examined the interaction of the sulfobetaines with a cationic polymer, poly(diallyldimethylammonium chloride), PDADMAC. Results for both interactions with PA and PDADMAC are shown in Figure 3, indicating no significant difference from the dilution curve. These results for PDADMAC are different from the ones reported by Velázquez,25 who proposed a cac for zwitterionic surfactant CHAPS and PDADMAC when performing a fluorescence study. It is also clear that only an electrostatic interaction is not enough to promote such an interaction, as confirmed by the results for PA with the sulfobetaine. With such results, another experiment was performed, with a negatively charged polymer containing a hydrophobic group on its backbone, poly(sodium 4-styrenesulfonate), PSS. The use of a more diluted surfactant solution (as shown in the inset of Figure 4) allowed the determination of a cac of 0.03 mmol kg−1. This cac value does not seem to vary for different polymer concentrations, agreeing with previous studies.17,18 Our obtained value for the cac with SB 3-12 and PSS (0.17 mmol kg−1) is smaller than that reported by Velázquez with the same type of sulfobetaine, 0.67 mmol L−1 by using surface tension methods24 and 0.4 mmol L−1 by fluorescence.25 As can be seen in Figure 4, initially there is an exothermic peak at low surfactant concentration. The enthalpy then

obtained values for the hydrophobic contributions (Supporting Information) are very similar to those reported for other surfactants, indicating that micelles of these zwitterionic surfactants do not differ from those of other nonionic (ethoxylated32 or glucoside35) or ionic surfactants.29 For the headgroup, ΔΔH0micHG is a sum of dehydration energies, interaction between the dipoles, and the confinement of the headgroups at the micelle surface. It is always positive, and its value decreases with temperature. The ΔTΔS0micHG term is positive only at 15 and 25 °C and then becomes negative, and it should reflect the overall loss of degrees of freedom, most likely from the dehydration of the headgroups. Both contribute to the positive value of the ΔΔG0micHG term. As for the alkyl chain, the ΔΔH0micCH2 term is negative over the whole temperature range and thus contributes favorably to micellization. The ΔTΔS0micCH2 term is positive and is related to the gain in the number of degrees of freedom of water molecules that were once hydrating the alkyl chains and are now released into the bulk. Sulfobetaines’ Interaction with Polymers. Because the profiles of the calorimetric curves for both sulfobetaines are the same, we present only the titration curves of SB 3-14 and its interaction with polymers (with the results for polymer interaction with SB 3-12 presented as Supporting Information). The titration calorimetric curves presented in Figure 2 for polymers PEO and PPO are similar and present the same profile of the dilution curve of the sulfobetaine in water. All of

Figure 2. Calorimetric titration curves for the addition of SB 3-14 (6.0 mmol kg−1) at 15 °C to water (□) and to a 0.1 wt % polymeric solution of PEO (●), PPO (△), and PAA (★) at pH 3.5. 6005

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Figure 3. Calorimetric titration curves for the addition SB 3-14 (6.0 mmol kg−1) at 15 °C to water (□) and to a 0.1 wt % polymeric solution of PDADMAC (●), PA at pH 7.5 (△), and pH 11.7 (★).

Figure 5. Comparison of calorimetric titration curves for the addition of SB 3-14 (6.0 mmol kg−1) at 15 °C in the absence (open symbols) and in the presence of 0.25 mol L−1 NaCl solution (filled symbols).

Figure 4. Calorimetric titration curves for the addition of SB 3-14 (6.0 mmol kg−1) at 15 °C to water (□) and to PSS 0.002 wt % (●) and 0.01 wt % (☆).

Figure 6. Comparison of calorimetric titration curves for the addition of SB 3-14 (6.0 mmol kg−1) to water and PSS 0.01 wt % at 15 °C (open symbols) and at 25 °C (filled symbols).

reaches a minimum, indicating the point of maximum cooperativity, and increases upon addition of more surfactant, reaching a plateau. These values are considered to be associated with the saturation concentration, C2, because constant enthalpy values are achieved. From this concentration onward, both polymeric titration curves resemble the dilution curve, only they are now shifted to higher surfactant concentrations and smaller values of ΔH0obs. The dilution and polymer curves eventually merge at a higher surfactant concentration of 1.0 mmol kg−1 for the 0.002 wt % PSS, while for the 0.01 wt % PSS this is close to happening at the highest studied concentration. For further analyses of the nature of this interaction, the same experiment was performed in the presence of NaCl 0.25 mol L−1 (Figure 5) and at a higher temperature (25 °C) (Figure 6). The presence of electrolytes is expected to partially screen the positive and negative charges on the sulfobetaine as well as the polyelectrolyte, reducing possible electrostatic interactions. Also, the solubilization process of the styrene group into the surface of the micelle is favored because its

solubility in water is reduced due to the presence of electrolytes.36 More importantly, temperature changes should also affect hydrophobic interactions due to the important ΔCp contributions. Moreover, the shift to 25 °C would allow the disregarding of enthalpic contributions arising from surfactant association because this is the temperature of athermal micellization. The results shown in Figures 5 and 6 reveal that the interaction profiles are not significantly altered by the temperature change, while a more exothermic interaction appears in the presence of NaCl.

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DISCUSSION Micellization Thermodynamics of Sulfobetaines. Overall, the thermodynamics of micellization of this family of sulfobetaines resembles the features reported for other ionic and nonionic surfactants. The most important differences, as expected, arise from the contributions of the dipolar headgroup.

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Comparing the values obtained at 25 °C for the betaines with values reported for other families of cationic,29,37 anionic,30,38 nonionic31,35 and zwitterionic4 surfactants, one observes that all values are positive and the value obtained by Kresheck31 for a homologous series of alkyldimethylphosphines, ΔΔH0micHG = 23.5 kJ mol−1, is the closest compared to 22.3 kJ mol−1 found in the present study. This agreement may be attributed to the dipolar characteristics of both headgroups. The larger values were obtained for cationic surfactants (16−38 kJ mol−1), while the lower values were reported for anionic surfactants (ca. 10 kJ mol−1), the latter being estimated from van’t Hoff analyses. The values calculated from the reports with nonionic (16−23.5 kJ mol−1) and other zwitterionic (18−22.3 kJ mol−1) surfactants are intermediate. The ΔTΔS0micHG term is positive only below 25 °C and then becomes negative at higher temperatures, corresponding to an unfavorable contribution to micellization. Both enthalpic and entropic contributions of the headgroups may be related to their dehydration associated with micelle formation. The hydrophobic contributions, on the other hand, are all favorable to micellization and agree with the general picture of the hydrophobic effect, including features of enthalpy−entropy compensation (especially with respect to temperature effects). Overall, the ΔΔG0micCH2 values observed are −2.9 ± 0.2 kJ mol−1. This is in agreement with several other studies3,29,34,35,39−41 related to the transfer of organic molecules from a polar solvent to a nonpolar solvent, including micelles. Surfactant Interaction with Polymers. The interaction of betaines with three nonionic polymers was studied. No sign of interaction was observed for the more hydrophilic PEO. With the more hydrophobic PPO, only after micelle formation was there evidence of an endothermic interaction because the titration curve in the presence of PPO was ca. 1.5 kJ mol−1 above the dilution curve. This endothermic contribution is attributed to the dehydration of the hydrophobic PO groups and its incorporation into the micelles, as seen in several studies with ionic surfactants and nonionic polymers.9,15,17,21 Among the nonionic polymers studied, PAA is the one which presents the most distinct result, as can be seen in Figure 2. At pH 3.5, in which the polymer is mostly undissociated, there is clear evidence of induced surfactant aggregation due to interaction with poly(acrylic acid). It is possible to analyze the intensity of interaction between surfactant and polymer18 using the Gibbs energy difference described by

⎛ cac ⎞ 0 ⎟ ΔG PS = RT ln⎜ ⎝ cmc ⎠

present in the headgroup of the sulfobetaines, favoring the formation of micellar aggregates around the polymer chains. As the surfactant concentration increases, more aggregates are formed around the polymer chains, resulting in more negative values of the observed enthalpy. This hypothesis is confirmed when the pH of the polymer solution increases, leading to greater polymer ionization. No interaction was observed between surfactant and poly(acrylate); even the cmc values remain the same within experimental error. The hydrophobicity of the polymer is already reduced, and the few remaining protonated carboxylic groups are not capable of forming hydrogen bonds, thus confirming the previous proposal that hydrogen bonds are responsible for the interaction between the sulfobetaine and the poly(acrylic acid). Regarding only electrostatic interactions between betaine surfactant and ionic polymers, no interaction at all resembling a polymer-induced aggregation was observed whether the polymer contained negatively charged groups (−COO−) or positively charged groups (RN+Me2), as indicated by the results obtained with PDADMAC (Figure 3). This finding does not agree with the results reported by Velázquez et al.25 based on fluorescence measurements using Nile red as a probe to study the interaction with another sulfobetaine, 3-[(3-cholamidopropyl)-dimethylammonium]-1-propanesulfonate], CHAPS, and polymer PDADMAC. The zwitterionic surfactants used in the two studies display the same headgroup, but the previous one has a hydrophobic group related to bile salts. However, in this previous investigation an interaction of CHAPS with PSS is reported, as seen in our present results (Figures 4−6), suggesting a significant role of PSS in inducing betaine aggregation. The small value ascribed for the cac of these betaines with PSS implies a highly favorable and cooperative interaction in the process of formation of the surfactant−polymer complex and is associated with ΔG0PS values of −7.5 and −6.5 kJ mol−1 for SB 3-12 and SB 3-14 respectively, for a polymer concentration of 0.01 wt %. Considering the fact that the hydrophobicity of SB 3-14 is larger than that of SB 3-12, we would expect it to reflect on the ΔG0PS values as it is for PAA, which is not the case. This may indicate that for both surfactants the most important contribution comes from the removal of the aromatic moiety of PSS, which should be located in similar environments in both micelles. Increasing the polymer concentration leads to more negative values of ΔH0obs, denoting a greater extent of interaction because there are more available sites for surfactant binding. We believe the interaction happens because the styrene groups, which are very hydrophobic, are incorporated into the micelles. Skerjanc et al.42 proposed that the hydrophobic interactions between cationic surfactants and PSS are more important that the electrostatic ones, also agreeing with a report by Hansson et al.43 Different studies involving oppositely charged surfactants and polymers12,13,44 show calorimetric titration curves similar to the ones reported here in the presence of an exothermic peak at low surfactant concentration. After going through a minimum, saturation of the polymer chain starts and the electrostatic repulsion among the micellar aggregates formed around the polymer chains and the surfactant monomers in solution start to increase, causing the enthalpy values to be less negative. Considering that our system does not present electrical repulsion between the polymer-induced aggregates, we tried to investigate whether micellar growth could be occurring due

(4)

The more negative the term, the more favorable the interaction between surfactant and polymer. For comparison, the values for sulfobetaines SB 3-12 and SB 3-14 and PAA are −2.3 and −3.1 kJ mol−1, respectively, confirming a contribution from the surfactant hydrophobicity as reported earlier.18 Moreover, when these values are compared with others reported for the interaction of ionic surfactants with nonionic polymers,17,18 they are smaller than the ones for anionic surfactants such as SDS and slightly above those observed for cationic surfactants, positioning the betaine interaction intensity between that of these two ionic surfactants. There must be a driving force for this interaction to occur, one that is exothermic in its nature. An explanation of this resides in the aforementioned ability of the polymer chains to make hydrogen bonds with the sulfonate groups (−SO3−) 6007

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contributions at higher temperatures, with thermodynamic parameters that are within the range of literature values. Nevertheless, their closest resemblance with dipolar nonionic surfactants such as alkylphosphines is seen by the enthalpic contribution of their headgroups toward micellization. Regarding their interaction with polymers, no sign of interaction was observed with hydrophilic nonionic polymers such as PEO or polyelectrolytes regardless of the sign of their electrical charge. Signs for the incorporation of the more hydrophobic PPO appeared above the cmc, which remained unchanged. Drastic changes in titration curves, however, were observed with hydrophilic polymer PAA, in which hydrogen bonding between the PAA carboxylic group and the sulfonate group of the surfactant is identified as the key interaction. In addition, the interaction of betaines with the more hydrophobic polyanion PSS was also observed, with rather low cac values indicating a favorable interaction that was ascribed to the incorporation of the hydrophobic pendant aromatic rings of the polymer into surfactant aggregates. Overall, this comprehensive investigation revealed two important specific interactions that are extremely relevant for this kind of surfactant and that should be taken into account in other studies that involve mixed systems containing betaines. Considering the wide use of sulfobetaines in formulation and their resemblance to polar groups of important membraneforming phospholipids, we believe a clear indication of how these interactions occur is of great importance to those working with these surfactants.

to polymer effects by means of dynamic light scattering measurements. These samples are dilute, but experiments with a polymer concentration 10 times more concentrated than the one used for calorimetric experiments show that there is no significant change in the size of the micelles (Supporting Information), with a hydrodynamic radius of around 3.0 nm for pure micelles, agreeing with the value reported by Wang et al.,4 and 2.0 nm in the presence of polymer. The nature of the interaction between betaines and PSS was also investigated by adding electrolytes to solution, which is expected to shield the electric charges, and if electrostatics is the prevailing interaction, it should lead to a decreased interaction. Regarding the presence of electrolytes in solution (Figure 5), there is a slight decrease in the values of both the cmc and ΔH0mic in the presence of NaCl. Its cmc values vary from 0.37 to 0.31 mmol kg−1 due to a decrease in the repulsion of both the positively and negatively charged groups on the headgroup of the surfactant, and the ΔH0mic value decreases from 5.75 to 5.47 kJ mol−1 probably because of the interaction of ions with the micelle surface, causing them to display a negative charge as reported earlier.6 Moreover, the beginning of aggregation (cac) in the presence of PSS is affected by NaCl, while the enthalpy values for the interaction are more negative. We tried to investigate more precisely the cac region in the presence of salt by diluting the surfactant sample approximately 10-fold. Unfortunately, the signal-to-noise ratio in the calorimeter was too small for us to assess this value. Besides this, the titration curve in polymer solution containing salt resembles the curve in polymer solution. Overall, these are not the changes one would expect if the main reason for the interaction of betaines with PSS was an electrostatic interaction. The same kind of behavior was observed when the temperature of the system increased (Figure 6): a large exothermic peak with the very first injections, resembling the curve at 15 °C in its shape and magnitude, with a ΔH0obs of −9.0 kJ mol−1 at a surfactant concentration of 0.1 mmol kg−1. Enthalpy values for this interaction vary more significantly with temperature, as observed for the enthalpy values of micellization, thus confirming our previous statement that this kind of interaction between PSS and SB is hydrophobic in nature. The comparison of the curves shown in Figures 4−6 allow an estimate of the ratio (n) of bound surfactant molecules per polymer repeat unit, which can be determined from (a) n = (C2 − cac)/Cmon, where Cmon stands for the concentration of monomers, or (b) between the inflection points of the polymer and dilution curves divided by the concentration of monomers.17 For a polymer concentration of 0.01 wt %, these values are 4.0 and 2.7 for SB 3-12 and 1.4 and 1.9 for SB 3-14 using methods a and b, respectively. For a polymer concentration of 0.002 wt %, the values are 3.1 and 2.4, respectively, and do not vary significantly with temperature. These numbers are larger than the ones reported for the interaction of ionic surfactants with nonionic polymers,18 which are typically around 1.

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ASSOCIATED CONTENT

S Supporting Information *

Titration calorimetric curves obtained for sulfobetaine SB 3-14 at different temperatures. ΔH0mic as a function of temperature for the homologous series of sulfobetaines. ΔG0mic, ΔH0mic, and TΔS0mic as functions of the number of carbon atoms on the alkyl chain. Calorimetric titration curves for the addition of SB 3-12. Intensity−time autocorrelation functions. Distribution functions of decay time as a result of the intensity-weighted inverse Laplace transformation of the correlation function. This material is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We acknowledge Brazilian agencies FAPESP and CNPq for financial support. W.L. thanks CNPq for a productivity research grant, and C.B. and L.B.M. thank CNPq for their scholarships. We also thank Gerd Olofsson for valuable discussions of the calorimetric results and Ana Maria Percebom for her help in analyzing the light-scattering data. This is a contribution from INCT-Catálise.

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CONCLUSIONS From the results described in this report, the thermodynamics of micellization of a homologous series of sulfobetaines is found to be very similar to that of other surfactant families, following the same trends: the process is entropically driven at lower temperatures and jointly driven by the enthalpy and entropy

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