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Langmuir 1997, 13, 1486-1495
Thermodynamics of Micellization and Adsorption of Zwitterionic Surfactants in Aqueous Media J. Zajac,* C. Chorro, M. Lindheimer, and S. Partyka Laboratoire des Agre´ gats Mole´ culaires et Mate´ riaux Inorganiques, Universite´ des Sciences et Techniques du Languedoc, pl. E. Bataillon, 34095 Montpellier Cedex 05, France Received September 23, 1996. In Final Form: December 19, 1996X Titration microcalorimetry has been applied to study the micellization and adsorption of the following zwitterionic surfactants: (1) the C12 and C14 homologues of (alkyldimethylammonio)ethanoate and (alkyldimethylammonio)-1-propanesulfonate (referred to as C12N1C, C14N1C, C12N3S, and C14N3S, respectively); (2) (dodecyldimethylammonio)butanoate (C12N3C); (3) three surfactants of the [(dodecyldimethylammonio)alkyl]phenylphosphinate type with the intercharge alkanediyl group being C3H6, C6H12, and C10H20 (C12N3PPh, C12N6PPh, and C12N10PPh, respectively). In regard to the basic principles of both the micellization and the adsorption on a hydrophilic silica surface, the zwitterionic surfactants studied fall into line with polyoxyethylenated nonionics. Enthalpy of dilution measurements allowed determining the standard molar enthalpies of micellization (∆micho) and thus, basing on the appropriate critical micelle concentration values, estimates could be made of the standard molar energies of micellization (∆micgo) to furnish the corresponding entropy changes, ∆micso. At room temperatures, the negative free energy term is the result of an unfavorable positive enthalpy change and a favorable increase in entropy, the latter contribution dominating. The positive values of ∆micho decrease with rising temperature, addition of NaCl, decreasing hydrophilic character of the zwitterionic headgroup, and increasing length of the alkyl chain. Adsorption isotherms for the same surfactants on two hydrophilic silica surfaces, Spherosil XOB015 and precipitated silica RP 63-876, were measured using the solution depletion method. The related changes in the differential molar enthalpy of displacement (∆dplh) were detected calorimetrically as a function of surface coverage, Θ. Adsorption of surfactant betaines is the interplay of two mechanisms: (1) ion-dipole bonding between the ionic surface sites and the zwitterionic headgroups oriented with the quaternary nitrogen group close to the surface and the anionic substituent group away from it; (2) formation of surface aggregates, induced by the hydrophobic effect. At higher coverages, ∆dplh is a constant function of Θ, the trends with changing temperature, salt content, and molecular structure of the surfactant essentially paralleling the corresponding ∆micho changes. Strong cooperativity in the adsorption phenomenon results in extended surface aggregates or even bilayers in the adsorption plateau region; otherwise the adsorbed phase is fairly fragmented.
Introduction Surface-active inner salts containing a quaternary ammonium cation, the so-called surfactant betaines, have achieved much commercial importance. The principal applications of these zwitterionic surfactants are based on their unique properties, such as mildness to the skin and eyes, excellent water solubility, broad isoelectric ranges, and resistance to hard water.1 In addition, the effect of pH, temperature, and added electrolyte on their solution and surface-active properties has been found to be minimal. The presence of both positively and negatively charged hydrophilic groups in the same molecule leads to the headgroup hydrophilicity intermediate between the ionic and conventional nonionic classes.2 Moreover, the hydrophilic character can be modified by varying the number of methylene groups separating the charged centers.3-8 An interesting and important property is the synergism between surfactant betaines and other types * To whom correspondence should be addressed. E-mail: zajac@ univ-montp2.fr. X Abstract published in Advance ACS Abstracts, February 1, 1997. (1) Ernst, R.; Miller, E. J., Jr. In Amphoteric Surfactants; Bluestein, B. R., Hilton, C. L., Eds.; Marcel Dekker: New York, 1982; Chapter II. (2) Laughlin, R. G. Langmuir 1991, 7, 842. (3) Chevalier, Y.; Germanaud, L.; Le Perchec, P. Colloid Polym. Sci. 1988, 266, 371. (4) Chevalier, Y.; Le Perchec, P. J. Phys. Chem. 1990, 94, 1768. (5) Chevalier, Y.; Stortet, Y.; Pourchet, S.; Le Perchec, P. Langmuir 1991, 7, 848. (6) Weers, J. G.; Rathman, J. F.; Axe, F. U.; Crichlow, C. A.; Foland, L. D.; Scheuing, D. R.; Wiersema, R. J.; Zielske, A. G. Langmuir 1991, 7, 854. (7) Chevalier, Y.; Me´lis, F.; Dalbiez, J. P. J. Phys. Chem. 1992, 96, 8614.
S0743-7463(96)00926-2 CCC: $14.00
of surfactant.9 They are most commonly employed in combination with anionic and nonionic compounds but may also be found in systems containing cationic surfactants. Surfactant betaines, especially those having a carboxylate and sulfonate ion, have been already studied using a variety of experimental techniques.3-20 Their solubility behavior and surface activity at the air-water interface are very well known. The ability to adsorb at the solidwater interface has received much less attention.12-20 A (8) Kamenka, N.; Chevalier, Y.; Zana, R. Langmuir 1995, 11, 3351. Kamenka, N.; Chorro, M.; Chevalier, Y.; Levy, H.; Zana, R. Langmuir 1995, 11, 4234. Chevalier, Y.; Kamenka, N.; Chorro, M.; Zana, R. Langmuir 1996, 12, 3225. (9) Rosen, M. J. Langmuir 1991, 7, 885. (10) Brochsztain, S.; Filho, P. B.; Toscano, V. G.; Chaimovich, H.; Politi, M. J. J. Phys. Chem. 1990, 94, 6781. (11) Graillat, C.; Dumont, B.; Depraetere, P.; Vintenon, V.; Pichot, C. Langmuir 1991, 7, 872. (12) Clint, J. H. J. Colloid Interface Sci. 1973, 43, 132. (13) Belambri, N. O.; Vanel, P.; Schumann, D. J. Colloid Interface Sci. 1987, 120, 224. (14) Amin-Alami, A.; Kamenka, N.; Partyka, S. Thermochim. Acta 1987, 122, 171. (15) Brode, P. F., III Langmuir 1988, 4, 176. (16) Mannhardt, K.; Schramm, L. L.; Novosad, J. J. Colloids Surf. 1992, 68, 37. (17) Leaver, I. H.; Jurdana, L. E. J. Colloid Interface Sci. 1992, 153, 552. (18) Partyka, S.; Lindheimer, M.; Faucompre´, B. Colloids Surf., A 1993, 76, 267. (19) Zajac, J.; Chorro, M.; Chorro, C.; Partyka, S. J. Thermal Anal. 1995, 45, 781. Zajac, J.; Chorro, M.; Chorro, C. Prog. Colloid Polym. Sci. 1995, 98, 29. (20) Chorro, M.; Kamenka, N.; Faucompre´, B.; Partyka, S.; Lindheimer, M.; Zana, R. Colloids Surf. 1996, 110, 249. (21) Zajac, J.; Partyka, S. In Adsorption on New and Modified Inorganic Sorbents; Dabrowski, A., Tertykh, V. A., Eds.; Elsevier: Amsterdam, 1996; Chapter III.6.
© 1997 American Chemical Society
Micellization and Adsorption of Zwitterionic Surfactants
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Table 1. Values of the Critical Micelle Concentration (cmc), the Experimental Standard Enthalpy of Micelle Formation at the cmc (∆micho), and the Evaluated Standard Free Energy (∆micgo) and Entropy (∆micso) of Micellization for Surfactant Betaines under Various Experimental Conditions surfactant
solvent
T (K)
cmc (mmol kg-1)
∆micho (kJ mol-1)
∆micgo (kJ mol-1)
T ∆micso (kJ mol-1)
C14N1C C12N1C C12N1C C12N1C C12N1C C12N3C C12N3C C14N3S C12N3S C12N3S C12N3S C12N3PPh C12N6PPh C12N10PPh
water water water 0.1 M NaCl 1 M NaCl water water water water 0.1 M NaCl 1 M NaCl water water water
298 298 308 298 298 298 308 298 298 298 298 298 298 298
0.22 1.9 1.9 1.6 1.0 4.6 4.6 0.32 3.0 2.6 1.7 1.35 0.9 0.47
0 4.6 0 3.1 0 8.8 2.5 0 3.6 3.4 2.6 10.0 7.1 0
-20.9 -15.5 -16.0 -16.0 -17.1 -13.3 -13.8 -19.9 -14.4 -14.7 -15.8 -16.4 -17.4 -19.0
20.9 20.1 16.0 19.1 17.1 22.1 16.3 19.9 18.0 18.1 18.4 26.4 24.5 19
very interesting hypothesis has been formulated based on adsorption studies of some carboxybetaines on negatively charged surfaces of quartz and amorphous silica.15,18-20 Although these compounds carry no formal net charge, they can adsorb as individual molecules with positively charged substituent group directly opposite a negatively charged site on the surface. This individual adsorption stage, beginning at bulk concentrations much below the critical micelle concentration (cmc), is followed by cooperative adsorption, mainly resulting from the hydrophobic association between the surfactant tails. Ultimately, surface aggregation leads to the formation of compact adsorbate structures having a bilayer character. The intention of the present study was to determine the basic thermodynamic parameters of micellization and adsorption in aqueous media in order to improve current understanding of both phenomena. Emphasis was placed on the influence of zwitterionic headgroup size and hydrophilicity, the two structural factors being dependent on the nature of the anionic substituent group and the intercharge separation distance (tether length). Zwitterionic surfactants herein covered belong to three homologous series: (a) (alkyldimethylammonio)alkanoates; (b) (alkyldimethylammonio)alkanesulfonates; (c) [(alkyldimethylammonio)alkyl]phenylphosphinates. The comparative study includes various homologues obtained by lengthening the ion bridge. Thermodynamic parameters are also presented as a function of temperature, added electrolyte, and alkyl chain length. The effect of the last three factors on the surfactant ability to adsorb on a polar surface has already been quantified for the ultra-longchain homologues of (alkyldimethylammonio)hexanoate15 and for (dodecyldimethylammonio)ethanoate.20 This paper reports data for other zwitterions. The molar enthalpy of micelle formation in aqueous solution is calculated directly from enthalpy of dilution measurements and permits estimates to be made of the molar free energy and entropy of micellization. Thermodynamic description of adsorption from aqueous solution onto surface-hydroxylated amorphous silica is inferred from the analysis of experimental adsorption isotherms and calorimetric enthalpies of displacement. The adsorption isotherm is the most popular tool utilized in measuring surfactant adsorption, whereas calorimetry has already proved to be useful in studying the interactions of ionic and nonionic surfactants with solid substrates. In particular, comparison between enthalpies of micellization in the bulk phase and differential enthalpies of displacement at higher surface coverages should shed light on the nature of surface aggregates.
Materials and Experimental Methods Materials. Two surface-hydroxylated amorphous silicas were used as solid supports. Synthetic silica gel in the form of Spherosil XOB015 was supplied by Rhoˆne-Poulenc (France). It was produced by agglomeration of small elementary spheres with the average diameter of 100-200 nm. This macroporous powdered substrate with spherical particles of diameter between 40 and 100 µm had a BET surface area of 25 m2 g-1 (σN2 ) 16.2 Å2), pores being made of the interstices between packed elementary spheres. Mercury porosimetry revealed an average pore size of 125 nm (the cumulative pore volume of 1 cm3 g-1). It was used as received. Silica RP 63-876 was manufactured in Rhoˆne-Poulenc and Aubervilliers Laboratory (France) by precipitation from sodium silicate solution. It was a nonporous adsorbent with a BET surface area of 40 m2 g-1. The mean particle size, observed in a scanning electron micrograph, was 0.13 µm. This sample was used without further treatment. The following surfactant molecules were investigated: (1) (tetradecyldimethylammonio)ethanoate, referred to as C14N1C; (2) (dodecyldimethylammonio)ethanoate, C12N1C; (3)(dodecyldimethylammonio)butanoate, C12N3C; (4)(tetradecyldimethylammonio)-1-propanesulfonate, C14N3S; (5)(dodecyldimethylammonio)-1-propanesulfonate, C12N3S; (6)[(dodecyldimethylammonio)propyl]phenylphosphinate, C12N3PPh; (7)[(dodecyldimethylammonio)hexyl]phenylphosphinate, C12N6PPh; (8)[(dodecyldimethylammonio)decyl]phenylphosphinate, C12N10PPh. They all were kind gifts from Dr. Y. Chevalier, Laboratoire des Mate´riaux Organiques a` Proprie´te´s Spe´cifiques, CNRS, Vernaison, France. The methods used for their synthesis and purification are described elsewhere.3,5,7,22 The values of the critical micelle concentration (cmc) under various experimental conditions, determined in the surface tension measurement, are reported in Table 1. Surface tension was measured by the Wilhelmy method with a Prolabo tensiometer (Tensiomat 2000). In neutral media, a polar headgroup of the carboxybetaines studied exists as a zwitterion.6 The water used throughout all experiments was deionized and purified with a Millipore “Super Q” system. It had a pH value of about 6 and its conductivity ranged between 0.05 and 0.1 µS cm-1. Sodium chloride, utilized to increase the ionic strength, was of analytical reagent grade. Methods. Adsorption isotherms were measured at free pH using the depletion method. The mixtures of known amounts of silica with a known mass of calibrated surfactant solution were equilibrated at a constant temperature for 24 h by slow rotation in glass joint stoppered tubes of capacity 30 mL. Subsequently the solid samples were separated from the supernatant liquids by centrifugation (20 000 rpm during 30 min). The refractometric technique (Water Associates R-403 Differential Refractometer) was employed for measuring the final concentration of the supernatant liquid. In the case of phenylphosphinatobetaines, the UV spectroscopy (UV Varian Spec(22) Amin-Alami, A. Ph.D. thesis, University of Montpellier II, 1989.
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trometer, λmax ) 263.6 nm) allowed the concentration determinations to be achieved with more precision. ‘Montcal’ microcalorimeter was used to measure enthalpies of dilution and displacement. Detailed description of the apparatus and its functional parts, the operating principles, technical specifications, and procedures for operation were published previously.19,21,23 The micellar stock solution of molality about 10 times greater than the cmc was injected into the calorimetric cell in small steps. In the adsorption experiment the calorimetric cell was initially filled with 0.5 g of silica and 17 g of solvent. The differential molar enthalpy of displacement corresponding to a given adsorption step was calculated from the experimentally measured enthalpy change, effects of dilution, and adsorption isotherm.
Results and Discussion Enthalpies of Dilution and Micellization. The values of the molar enthalpy of dilution, ∆dilh, as a function of the solution molality, m, are measured by successive dilution of a micellar stock solution in the solution from a previous injection (in a pure solvent for the first injection). The total change in enthalpy during the ith injection may be expressed in terms of the apparent molal enthalpy, ΦH ) ΦH(m), as follows i
∆injHi ) (
∑ n2 )[ΦH(mi) - ΦH(mo)] k
k)1 i-1
(
∑ n2k)[ΦH(mi-1) - ΦH(mo)] ) ∆dilH(mofmi) k)1 ∆dilH(mofmi-1) (1)
where mo is the molality of the stock solution; mi and mi-1 are the molalities of the equilibrium solution after the injection i and i - 1, respectively, n2k is the number of moles of the solute injected into the calorimetric cell during the kth injection, and ∆dilH is the enthalpy of dilution from a molality mo to a molality m. Since the quantity of the solute injected during a single injection in any run never exceeds several micromoles, ∆dilh values, calculated as ∆dilhi ) ∆injHi/n2i, correspond to the differential molar enthalpies of dilution. A few enthalpic curves of dilution, representative of the surfactant betaines studied in this paper, are presented in Figure 1. An important feature of these systems is the constancy of ∆dilh in the premicellar region, in which the enthalpy effects result from destruction of micelles and dilution of unmicellized species, and in the postmicellar region, where the dilution of micelles occurs. This observation is consistent with dilution of a micellar solution containing monodisperse micelles. A similar conclusion has been inferred from the very little dependence of the aggregation numbers on the surfactant concentration and salt content of the solution.7,8 The differential molar enthalpy of dilution falls to zero at high molalities above the cmc which also suggests that the intermicellar interactions do not have much effect on this function. Moreover, the transition from the predominantly unassociated surfactant to the micellar state occurs over a fairly narrow range of molality, approaching to a great extent true phase separation. With the present surfactants the values of ∆dilh vary in a very regular fashion with the increasing molality and this can be useful for calculations of the standard enthalpy of micelle formation. The standard enthalpy of micellization is defined as the difference between the standard partial molal enthalpy of a surfactant molecule in a (23) Partyka, S.; Keh, E.; Lindheimer, M.; Groszek, A. J. Colloids Surf. 1989, 37, 309.
Figure 1. Differential molar enthalpies of dilution for some surfactant betaines under various experimental conditions. The enthalpic values (taken with the opposite sign) are plotted against the molality of the equilibrium bulk solution in the calorimetric cell.
micellar aggregate and the corresponding quantity for the free surfactant in aqueous solution.24 However, this definition is based on a nonrealistic assumption, according to which one is dealing with a solution infinitely dilute in micelles to permit neglect of nonideality terms arising from intermicellar interactions. Various operational definitions have been proposed in the literature to overcome this difficulty (e.g., see discussion of the problem in paper25). For the purpose of the present studies, the experimental standard enthalpy of micellization, ∆micho, is calculated as the molar change in partial molal enthalpy of the solute when monomers associate into a micelle at the same molality, i.e., at the cmc. Taking account of the relation between apparent and partial molal quantities,26 eq 1 can be transformed to give, for the sufficiently small n2i values
∆dilh(m) ) h2(m) - h2(mo)
(2)
where h2(m) is the partial molal enthalpy of a surfactant molecule in an aqueous solution of molality m. Therefore, to calculate the micellization enthalpy, the linear regions of ∆dilh ) ∆dilh(m) below and above the cmc are extrapolated to the cmc and the difference taken as ∆micho. Other thermodynamic functions of micelle formation may be evaluated from the cmc and calorimetric data. A crude estimate of the standard molar free energy of micellization is obtained from24
∆micgo ) RT ln cmc
(3)
and this, in turn, gives the corresponding standard molar entropy of micellization, ∆micso. Some values of ∆micho for the surfactants investigated under various conditions of NaCl content and temperature, and the related values of ∆micgo and T ∆micso, are listed in Table 1. Although the error in using the phase separation model for micelle formation (eq 3) for the estimation of ∆micgo is generally not severe, the above calculations are insufficiently precise to provide a quantitative explanation (24) Tanford, Ch. The Hydrophobic Effect: Formation of Micelles and Biological Membranes; Wiley: New York, 1973. (25) De Lisi, R.; Ostiguy, C.; Perron, G.; Desnoyers, J. E. J. Colloid Interface Sci. 1979, 71, 147. (26) Rossini, F. D. Chemical Thermodynamics, 3rd ed.; Wiley: New York, 1964.
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for the process. However, they may be helpful in arriving at some qualitative conclusions. The experimental enthalpies of micellization are positive in most cases indicating that the process is essentially endothermic. The negative free energy of micellization results from a favorable increase in entropy which overrides the unfavorable enthalpy term. Commonly the large entropy increase is ascribed either to structural changes in the solvent, associated with loss of hydration of the hydrophobic tail when the surfactant enters the micelle,24 or to increased freedom of the hydrophobic chain in the interior of the micelle compared to the aqueous medium.27 A detailed analysis of the values reported in Table 1 indicates that these effects occur during micellization of the surfactant betaines studied. Lengthening the hydrophobic surfactant tail greatly reduces ∆micho: the addition of two methylene units to a straight-chain dodecyl group is sufficient to render the process athermal. At the same time, the value of T ∆micso increases a little, rendering ∆micgo more negative by about 2.8 kJ mol-1 per one methylene group. The latter value is similar to those that have been obtained for nonionic surfactants28 (the extreme case of hydrophobic bonding corresponds to about 3.3 kJ mol-1 24). The role of the zwitterionic headgroup in the process cannot be neglected at all. The effective size of the dipolar functional group, its orientation with respect to the micelle surface, and dipole-dipole interactions in the micellar aggregate determine the micelle size, determine the hydration and intermicellar interactions,7,8 and make an important impact on the thermodynamic functions. The dependence of both entropy and enthalpy terms on ion bridge length (n) is complex and seems to follow changes in the relative hydrophilicity of the zwitterionic headgroup.2,3,5,6 The positive ∆micho and T ∆micso increase with n from n ) 1 to n ) 3 for n - 12 carboxybetaines and decrease from n ) 3 to n ) 10 for n - 12 phenylphosphinatobetaines; the molar free energy of micellization becomes respectively less negative and more negative. These changes in ∆micgo, as well as in ∆micho and T ∆micso, are a result of the balance between two opposing effects, which do not vary in a similar fashion as the tether length increases. On the one hand, methylene groups within the ion bridge make the zwitterionic molecule more hydrophobic. The average contribution to ∆micgo per one -CH2- group is much smaller for the polymethylene intercharge chain than it is for a single straight hydrophobic tail (1.1 kJ mol-1 for n - 12 carboxybetaines in the n range 1-3, and 0.4 kJ mol-1 for n - 12 phenylphosphinatobetaines in the n range 3-10). On the other hand, the addition of methylene groups to the chain separating charged centers increases the polarity of the dipolar functional group.2,3,5,6 Additionally, the effective dipole moment depends on the conformational structure of zwitterionic groups. For small n values, it rises linearly with n, but next this behavior is modified by the increasing flexibility of the intercharge chain.3,5,6 Long intercharge groups may even penetrate into hydrophobic micellar core and this reduces both hydration number and repulsive dipole-dipole interactions (steric hindrance effect due to the bulkiness of the headgroup becomes important).8 As a consequence of all the above tendencies, the surfactant betaine systems show a separation distance at which the cmc is a maximum (usually about n ) 3 or 43,5,6). Changing the structure of the anionic substituent group also appears to have a great influence on changes in
thermodynamic functions during micellization. The values of ∆micho and T ∆micso fall in the sequence C12N3PPh > C12N3C > C12N3S. This dependence is somewhat different from the order of effectiveness in decreasing the cmc, which is (for this particular value of n ) 3) -PO2C6H5> -SO3- > -CO2-. The total effect of a given physical factor, like temperature and salt content, is the resultant of its effects on the hydrophobic tail and hydrophilic surfactant headgroup. Unfortunately, the lack of precise thermodynamic data on a sufficiently wide range of systems does not allow exact calculations of these separate contributions. An increase in temperature reduces ∆micho to a great extent. It seems likely that if calorimetric data were accessible for higher temperatures, the related systems would be characterized even by negative values of ∆micho (exothermic micellization). The entropy term decreases with increasing temperature and differences in T ∆micso values essentially parallel ∆micho changes; this accounts for the roughly constant free energy at different temperatures. This behavior seems to be consistent with the hydrophobic effect, because the amount of water structured by the hydrocarbon chain has to diminish at higher temperatures (water loses some of its peculiar structure properties29), disfavoring micellization. Simultaneously, temperature increase induces a partial dehydration of the zwitterionic headgroup, thus creating conditions favorable for micellization. It should be noted that the micelle aggregation number decreases as temperature rises,8 so the degree of nonpolarity of the micelle core and the micelle hydration are modified. The presence of NaCl in aqueous solution makes ∆micho less positive, the effect being dependent on the salt content and the nature of the zwitterionic headgroup. For example, it is much less pronounced for C12N3S irrespective of the ionic strength. For a given surfactant molecule, a marked change is observed only at high ionic strength (1 M NaCl). Variations of T ∆micso are highly irregular. In salt solutions of C12N1C, the entropy term diminishes but to a lesser degree than does ∆micho. For another surfactant molecule (C12N3S), changes in T ∆micso are very small and positive. In both cases, the resultant free energy of micelle formation becomes more negative as ionic strength increases, the magnitude of the changes being only a little dependent on the nature of the zwitterionic group. The salting out and salting in of the hydrophobic chains by the electrolyte can serve as a basis for understanding the effects of NaCl on micellization of surfactant betaines.30 Electrolytes may moderate the repulsive dipole-dipole interactions between the headgroups or screen the electrostatic attractions between the opposite charges within the zwitterionic moiety, giving rise to a larger protrusion of the anionic substituent group toward water. In spite of this, the micelle aggregation number has been found to depend very little on the ionic strength (although the greatest NaCl concentration, used in ref 8, was only about 0.5 M). Moreover, the selective binding of electrolyte cations and anions to zwitterionic micelles modifies intermicellar interactions and affects the micellization process.8 Enthalpies of Displacement. In order to obtain the values of the differential molar enthalpy of displacement, ∆dplh, as a function of the amount adsorbed, Γ, the same experimental procedure of successive injections of a micellar stock solution into a calorimetric cell is applied. Compared to the dilution experiment the only difference
(27) Aranow, R. H.; Witten, L. J. Phys. Chem. 1960, 64, 1643; J. Chem. Phys. 1965, 43, 1436. (28) Benjamin, L. J. Phys. Chem. 1964, 12, 3575.
(29) Kavanau, J. L. Water and Solute-Water Interactions; HoldenDay: San Francisco, CA, 1964. (30) Ray, A.; Nemethy, G. J. Am. Chem. Soc. 1971, 93, 6787.
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is that the calorimetric cell contains a solid sample of a given mass or surface area MS. A small amount of the stock solution injected during the ith injection is diluted in the supernatant liquid inside the cell and some surfactant monomers subsequently adsorb onto solid particles. The partition of solute molecules between the adsorbed and bulk phases is strictly determined by the adsorption equilibrium at a given temperature, and therefore the amount adsorbed of the solute, Γ2i, and the molality of the equilibrium bulk solution, mbi, after the ith injection (at equilibrium) are evaluated from the adsorption isotherm.21 The total change in enthalpy during the ith injection may be written as
∆injHis ) MSΓ2i[h2s(mbi) - h2b(∞)] + MSΓ1i[h1s(mbi) h1o] + MSΓ2i[ΦH(∞) - ΦH(mbi)] MSΓ2i-1[h2s(mbi-1) - h2b(∞)] - MSΓ1i-1[h1s(mbi-1) h1o] - MSΓ2i-1[ΦH(∞) - ΦH(mbi-1)] + ∆dilH(mofmbi) - ∆dilH(mofmbi-1) (4) where Γ1i is the specific amount adsorbed of the solvent after the ith injection, hjs(mb) is the partial molal enthalpy of component j (j ) 1, 2) in the adsorbed phase, being at equilibrium with a bulk solution of molality mb, and h1o represents the molar enthalpy of pure solvent and h2b(∞) the partial molal enthalpy of the solute in the infinitely dilute solution, taken as a reference state for the equilibrium bulk solution; in such a case h2b(∞) ) ΦH(∞). The first two terms on the right hand side of eq 4 represent together the integral enthalpy of displacement, ∆dplH(mbi).21,31 The third term is a correction expression due to dilution of MS Γ2i moles of the solute from a molality mbi to infinite dilution. It is fairly small, because surfactant adsorption is studied for equilibrium bulk molalities mbi lower than the cmc (it may be included in ∆dplH). The overall enthalpy effect of such a displacement process will depend on the strength of both normal and lateral adsorptive bonds as well as the affinity of solvent to the solid support. All solute components other than the surfactant monomers (e.g., electrolytes) are lumped together as constituting a mean solvent. The apparent differential molar enthalpy of displacement may be calculated from eq 4 in the following way, for small n2i values accompanied by small increments in the amount adsorbed
∆dplhi ≈
∆dplH(mbi) - ∆dplH(mbi-1) MS(Γ2i - Γ2i-1)
) ∆injHis - n2i ∆dilhi MS(Γ2i - Γ2i-1)
(5)
The enthalpic curves of adsorption are plotted in terms of differential molar enthalpy of displacement as a function of the corresponding amount adsorbed at a given temperature, e.g., ∆dplhi ) ∆dplhi(Θi,mbi) ) ∆dplhi(Θi(mbi)), where the coverage ratio, Θi, is defined as percentage of the total amount adsorbed in the isotherm plateau. Variations of ∆dplh corresponding to the adsorption of surfactant betaines on two silica surfaces, Spherosil XOB015 and precipitated silica RP 63-876, are illustrated in Figures 2-4. Despite marked differences between the adsorption systems, all enthalpy curves have the same general shape. Since enthalpy of displacement is the (31) Denoyel, R.; Rouquerol, F.; Rouquerol, J. J. Colloid Interface Sci. 1990, 136, 375.
Figure 2. Differential molar enthalpies of displacement for various surfactant betaines adsorbed from aqueous solutions onto Spherosil XOB015 at two temperatures. The enthalpic values (taken with the opposite sign) are plotted against the coverage ratio.
Figure 3. Differential molar enthalpies of displacement for sulfobetaine surfactants adsorbed onto Spherosil XOB015 from various aqueous media at two temperatures. The enthalpic values (taken with the opposite sign) are plotted against the coverage ratio.
resultant macroscopic effect of the intermolecular interactions involved, such a striking similarity between the calorimetric results means that there is the same qualitative mechanism for the phenomenon. At very low coverage ratios, the values of ∆dplh are negative so that the process is exothermic. Large negative enthalpies have been attributed to individual adsorption of surfactant molecules.20 At this stage the surfactant is adsorbing on an empty surface, where there may be only a few adsorbed molecules, and therefore lateral adsorbateadsorbate interactions can be neglected. In aqueous solutions, a mineral surface of silica acquires a positive or negative electric charge depending on the pH of the aqueous phase and the concentration of a background electrolyte.32 Since all experiments were carried out at free pH (pH 6 for Spherosil XOB015 and pH 8.3 for precipitated silica RP 63-876) and in the absence of extra salt (with the exception of systems containing NaCl solutions), the density of negative charge on both silica surfaces was low or moderate at most. Individual adsorption of zwitterionic molecules may therefore occur because of ion-dipole interactions between hydrophilic (32) Bolt, G. H. J. Phys. Chem. 1957, 61, 1166.
Micellization and Adsorption of Zwitterionic Surfactants
Figure 4. Differential molar enthalpies of displacement for phenylphosphinatobetaine surfactants with varying tether length adsorbed from aqueous solutions onto precipitated silica RP 63-876 at 298 K. The enthalpic values (taken with the opposite sign) are plotted against the coverage ratio.
headgroups and negatively charged surface sites. The potential energy of such interaction depends on the orientation of dipoles with respect to the surface. Dipolar headgroups themselves tend to adopt a special orientation with the quaternary nitrogen group toward the surface and the anionic substituent group away from it, especially when the surface charge density becomes important (greater values of pH and ionic strength). The flexibility of ion bridge has also influence on the number of possible conformations. Dispersion forces acting between the alkyl chain and the solid support (possibly also between the tether chain and the surface) constitute a supplementary mechanism of individual adsorption providing there are no large hydrated electrolyte ions in the vicinity of the surface (they prevent attachment of the hydrophobic tails).33 With the molecules lying flat on the surface, the adsorption energy will be reduced by the energy required to remove the corresponding number of water molecules from the adsorbed layer. Dewetting of a mineral surface, especially that accomplished by the hydrophobic tails, renders it more hydrophobic, accounting for a marked increase in the captive bubble contact angle, as observed during adsorption of (tetradecyldimethylammonio)hexanoate on glass at very low surface coverages.15 When the surfactant is obliged to compete with salt ions for ionic surface sites, the enthalpy of displacement will additionally diminish. The behaviour of enthalpy curves, corresponding to various molecular structures of the surfactant, temperatures and ionic strengths, would allow shedding light on the relative orientation of the hydrophobic and hydrophilic parts of zwitterionic surfactants at the silica surface. Unfortunately, a detailed comparison of ∆dplh values is very difficult due to lack of experimental precision at very low coverage ratios. It is even impossible to locate exactly the end of this adsorption stage. In the next section, an attempt at determining the domain of individual adsorption will be made, using adsorption isotherms. The subsequent adsorption stages are increasingly dominated by adsorbate-adsorbate interactions. At higher adsorption densities, the driving force of adsorption will derive mainly from the hydrophobic effect, i.e., lateral chain-chain attractions and the tendency of hydrophobic tails to escape from an aqueous environment. This mode of adsorption is characterized by a constant enthalpy of (33) Zajac, J.; Trompette, J. L.; Partyka, S. Langmuir 1996, 12, 1357.
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displacement. For certain zwitterionic molecules, like C12N3C, C12N3PPh, and C12N6PPh, the constant value of ∆dplh is positive (endothermic adsorption). In other cases, adsorption is almost athermal (e.g., C12N1C, C12N3S in pure water and 0.1 M NaCl solution) or moderately exothermic (e.g., C12N3C and C12N3S at 308 K, C12N10PPh, and C12N3S in 1 M NaCl solution). The average enthalpy of displacement in this adsorption range, ∆dplhc, shows similar dependence upon temperature, ionic strength, and hydrophilic character of the headgroup as the enthalpy of micellization does. The former is decreased (adsorption becomes more exothermic) by an increase in temperature (see Figures 2 and 3) and by a large increase in the ionic strength of the solution (see Figure 3). Raising the temperature from 298 to 308 K makes ∆dplhc for C12N3C less positive by about 4.9 kJ mol-1, whereas the corresponding reduction in ∆micho is 6.3 kJ mol-1. The values of ∆dplhc for C12N3S are more sensitive to the presence of electrolyte: ∆dplhc is decreased by 4.5 kJ mol-1 when NaCl is added to obtain a concentration of 1 M (the analogous decrease in ∆micho is only 1 kJ mol-1). The same qualitative trends have been observed during adsorption of C12N1C onto Spherosil XOB015 at three temperatures (298, 308, and 318 K) and in the presence of salt (1 M NaCl and CaCl2 solutions).20 The values of ∆dplhc become more positive with increasing tether length (n) from n ) 1 to n ) 3 for n - 12 carboxybetaines and with decreasing n from n ) 10 to n ) 3 for n - 12 phenylphosphinatobetaines, differences in ∆dplhc paralleling ∆micho changes. Changing the structure of anionic group has the effect of reducing ∆dplhc in the order C12N3PPh > C12N3C > C12N3S. The same dependence has been obtained for the enthalpy of micellization. The difference between enthalpy values for the last two groups is 3.3 kJ mol-1 for adsorption and 5.2 kJ mol-1 for micellization. An analogous comparison of C12N3PPh with C12N3C or C12N3S is erroneous because both molecules are adsorbed on two different solid supports. One thing is clear from the behavior of ∆dplhc considered above: the structure of the adsorbed layer in the region of constant enthalpy of displacement is the interplay of the same intermolecular forces responsible for micelle formation in the bulk phase. The parallel between bulk solution micellization and the aggregation process occurring at the silica surface for (dodecyldimethylammonio)ethanoate has been emphasized in the previous paper.20 Here the hypothesis is extended on other surfactant betaines. For all surfactant systems, the values of ∆dplhc are less positive than the corresponding values of ∆micho; the formation of surface aggregates is favored energetically to a greater extent. Apparently some constraint placed on the adsorbing molecules by the interfacial environment prevents the adsorbed phase from attaining the same balance of the contributory effects as is attained in the bulk phase. The origins of these energetic and structural differences are as yet unclear, but they surely involve the fact that the zwitterionic headgroups of certain molecules forming a surface aggregate may be bound to a polar surface. This would disturb a subtle balance between various geometrical and packing factors, thus modifying the shape and size of interfacial aggregates (they are likely to be greater than the bulk globular micelles). The formation of surfactant aggregates at the surface is thus controlled by a pseudonucleation step, i.e., individual adsorption. The adsorbed molecules with hydrophobic tails oriented toward the solution act as nucleation centers for future aggregates formed through chain-chain association. The alkyl chains, initially lying flat on the surface, interact weakly with the support and
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may be easily displaced from it by the hydrophilic moieties of adjacent molecules. The driving force of adsorption at this stage results from the combination of adsorbentadsorbate and adsorbate-adsorbate interactions. As the adsorption density increases, the latter contribution becomes prominent and changes in the alkyl chain conformation at the surface are induced. The first small aggregates are subsequently formed. Their size and shape have to be strongly influenced by local surface properties such as charge density, relative affinity for the adsorbate and solvent, and energetic and geometric heterogeneity. It seems likely that the primary surface aggregates are characterized by a marked polydispersity. When they grow into more extended structures, the enthalpy of aggregation becomes more and more positive. As a consequence, the negative enthalpy of displacement in the intermediate region decreases with increasing amount adsorbed, approaching a limiting steady value ∆dplhc (see the enthalpy curves at low values of Θ in Figures 2-4). The values of ∆dplh fall to zero in the vicinity of Θ ) 1. This observation is consistent with the saturation of the adsorbed phase, which is attained at the plateau of adsorption. Adsorption Isotherms. The adsorption isotherms are plotted in terms of amount adsorbed per unit surface area of solid, Γ2, as a function of the molality of the equilibrium bulk solution, m2b, at a given temperature. With the very dilute surfactant solutions used in this work, the quantity of adsorption Γ2 is practically equal to the number of moles of the surfactant adsorbed per square meter of solid surface.21 The experimental adsorption isotherms, representative of surfactant betaines under various conditions of temperature and salt content, are shown in Figures 5-8 on two different scales: double linear and double logarithmic. The latter presentation of adsorption data have been chosen in order to better discriminate between the subsequent adsorption stages (the transition between regions of individual and cooperative adsorption appears much sharper on this scale than it is in reality). In adsorption of ionic surfactants on the oppositely charged mineral surfaces, the domain of individual adsorption is usually identified with the first region where the log-log plot is linear and has a slope of unity.34-36 The adsorption isotherms given in Figures 5a-8a are S-shaped with a well-marked plateau region which is generally reached somewhat above the cmc (note, however, that the presence of mineral impurities in a solid sample, which are released to the aqueous phase,33 is capable of modifying the ionic strength and thereby the actual cmc value). The corresponding values of Γ2 at saturation (Γmax) are given in Table 2. It is worth mentioning that isotherms having the same qualitative shape were obtained for adsorption of polyoxyethylenated alkylphenols (with short and moderate polar chains) on a hydrophilic surface of Spherosil.37 Calorimetric studies of the same nonionic surfactant systems38,39 yielded the enthalpy curves resembling in shape those presented in Figures 2-4. Fluorescence decay spectroscopy, applied by Levitz et al.37 to the micellization and adsorption of nonionics, provided (34) Somasundaran, P.; Fuerstenau, D. W. J. Phys. Chem. 1966, 70, 90. (35) Bo¨hmer, M. R.; Koopal, L. K. Langmuir 1992, 8, 2649 and 2660. (36) Harwell, J. H.; Schechter, R.; Wade, W. H. In Solid-Liquid Interactions in Porous Media; Cases, J. M., Ed.; Technip Publisher: Paris, 1985; p 371. (37) Levitz, P.; El Miri, A.; Keravis, D.; Van Damme, H. J. Colloid Interface Sci. 1984, 99, 484. Levitz, P.; Van Damme, H. J. Phys. Chem. 1986, 90, 1302. (38) Lindheimer, M.; Keh, E.; Zaini, S.; Partyka, S. J. Colloid Interface Sci. 1990, 138, 83. (39) Giordano-Palmino, F.; Denoyel, R.; Rouquerol, J. J. Colloid Interface Sci. 1994, 165, 82.
Zajac et al.
Figure 5. Effect of the alkyl chain length on the zwitterionic surfactant adsorption from aqueous solutions onto Spherosil XOB015: adsorption isotherms on a double linear (part a) and a double logarithmic scale (part b). The arrows indicate the critical micelle concentration.
a convincing argument for the existence of finite surface aggregates in the whole adsorption range (isolated adsorbed molecules are present only at very low surface coverages). Depending on the length of a polar headgroup, these surface structures can be seen either as “a selfrepeating micellar-like adsorbed layer” (the average size of surface aggregates is comparable with that of the bulk micelles), existing beyond some very low critical surface concentration, or as “more extended and gradually growing condensed regions” (surface aggregates are too large to be detected). When the hydrophobic and polar chains become similar in length, an intermediate behavior may even lead to the formation of a continuous phase in the plateau of the isotherm.37,40,41 These examples point to essential analogies between the adsorption mechanism for surfactant betaines and that for alkylphenol ethoxylates on hydrophilic surfaces. A detailed analysis of the log-log plots shown in Figures 5b-8b leads to the conclusion that individual adsorption of surfactant betaines onto silica occurs only for very small quantities of adsorption (usually below 0.1 µmol m-2). The initial isotherm parts on the double logarithmic scale are linear (although the number of experimental points is not always sufficient to clearly justify that conclusion); the resulting slopes and domains (up to a coverage ratio of Θo) are listed in Table 2. With the exception of adsorption systems containing NaCl, the slopes are less than unity. In the literature dealing with the adsorption of ionic surfactants, the same result is ascribed to individual (40) Cases, J. M.; Villieras, F. Langmuir 1992, 8, 1251. (41) Bo¨hmer, M. R.; Koopal, L. K. Langmuir 1990, 6, 1478.
Micellization and Adsorption of Zwitterionic Surfactants
Figure 6. Effect of the electrolyte addition on the adsorption of zwitterionic sulfobetaine surfactants from aqueous solutions onto Spherosil XOB015: adsorption isotherms on a double linear (part a) and a double logarithmic scale (part b). The arrows indicate the critical micelle concentration.
adsorption on heterogeneous surfaces.42 In the present case, a nonuniform distribution of negative charge may render the silica surface heterogeneous with respect to ion-dipole interaction. It is, however, more probable that these small slopes reflect a wide variety of adsorbate conformations at the interface, i.e., the combined effect of Coulombic and specific interactions between the surface and the various parts of the surfactant molecules, modified by the local affinity of solvent for the surface and the overall surface heterogeneity. The quantity of individually adsorbed molecules is reduced as the number of carbon atoms in both alkyl and tether chains increases. This observation is in agreement with an increase in the average interfacial area per one adsorbed molecule. There is thus evidence for the quasi-parallel orientation of the alkyl chains with respect to the silica surface. Full hydrophobic association of the tether chain with the surface is less probable, because the anionic substituent group is forced to extend away from the negatively charged surface. Nevertheless, repulsive electrostatic forces acting between the adjacent bulky headgroups are able to diminish the adsorption space accessible to individual adsorption. When the surfactant is adsorbed from a salt solution, the initial slope of the log-log plot becomes equal to or greater than unity. On the one hand, the addition of electrolyte to the bulk phase causes an increase in the surface density of negative charge, followed by adsorption of electrolyte ions at the silica-water interface. On the other hand, the latter should moderate all electrostatic interactions in the adsorbed phase (both attractive and (42) Lajtar, L.; Narkiewicz-Michalek, J.; Rudzinski, W.; Partyka, S. Langmuir 1993, 9, 3174.
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Figure 7. Effect of the intercharge separation distance and temperature on the adsorption of zwitterionic carboxybetaine surfactants from aqueous solutions onto Spherosil XOB015: adsorption isotherms on a double linear (part a) and a double logarithmic scale (part b). The arrows indicate the critical micelle concentration.
repulsive) and may preclude the possibility of the hydrophobic chains being extended on the surface. Low salt content only renders the interface more “homogeneous” and the slope is equal to unity (e.g., C12N3S in 0.1 M NaCl solution); the number of individually adsorbed molecules decreases. For high ionic strength (1 M NaCl), the principal ion-dipole bonding with the surface is probably so weak that surface aggregates have to form at this adsorption stage in order to compete successfully with electrolyte ions for surface sites. Adsorption of surfactant betaines becomes cooperative almost from the beginning, accounting for larger initial slopes. If adsorption of surfactant molecules with the hydrophilic headgroup in contact with the silica surface continues to occur beyond the region of individual adsorption, the surface aggregates will be different from globular bulk micelles. By analogy with nonionic surfactants, the ultimate structure of the adsorbed phase for zwitterionic betaines on a hydrophilic surface will be the total outcome of a cooperation or competition between two phenomena: a direct adsorption of surfactant molecules on the surface (especially that provoked by strong ion-dipole interactions) and association of the hydrophobic tails between themselves (induced by the hydrophobic effect). If both phenomena are strongly cooperative, extended aggregates will be expected to form on the surface. When direct adsorption is essentially reduced, association into threedimensional-like micelles is favored at the interface and the adsorbed layer should be very fragmented. These hypotheses may be easily verified by studying the effect of certain factors on the adsorption process at higher surface coverages.
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Zajac et al.
Figure 8. Effect of the intercharge separation distance and temperature on the adsorption of zwitterionic phenylphosphinatobetaine surfactants from aqueous solutions onto precipitated silica RP 63-876: adsorption isotherms on a double linear (part a) and a double logarithmic scale (part b). The arrows indicate the critical micelle concentration. Table 2. Effectiveness of Surfactant Betaine Adsorption on a Hydrophilic Silica Surface under Various Experimental Conditionsa surfactant
solvent
solid
C14N1C C12N1C C12N3C C12N3C C14N3S C14N3S C12N3S C12N3S C12N3S C12N3PPh C12N6PPh C12N6PPh C12N10PPh
water water water water water 1 M NaCl water 0.1 M NaCl 1 M NaCl water water water water
XOB015 XOB015 XOB015 XOB015 XOB015 XOB015 XOB015 XOB015 XOB015 RP 63-876 RP 63-876 RP 63-876 RP 63-876
T Γmax (K) (µmol m-2) 298 298 298 308 298 298 298 298 298 298 298 308 298
7.6 5.2 5.6 5.6 6.2 4.1 5.7 4.4 3.5 3.2 2.9 2.9 2.7
initial stage Θo
slp
0.009 0.022 0.007
0.8 0.7 0.4
0.013 0.032 0.019 0.012 0.047 0.015 0.004
0.8 2.1 0.7 1.1 3.9 0.5 0.3
a Γ max is the amount adsorbed at saturation; Θo is the upper limit of the initial adsorption stage; and slp is the average slope of the initial isotherm part when plotted on a log-log scale.
Lengthening the alkyl chain causes an increase in the effectiveness of adsorption, Γmax (Figure 5a). The intermediate part of the log-log plot becomes more vertical (Figure 5b). Both observations are consistent with a closer packing of longer chains (greater van der Waals attraction between them), resulting in a consequent increase in the average aggregation number at the solid-liquid interface. Comparison between the average cross-sectional area of one surfactant molecule adsorbed at the solid-liquid interface, σS (the reciprocal of Γmax), and its equivalent for the air-solution interface, σo, provides further information on the orientation of the adsorbate at the adsorbentsolution interface. The following values of σo have been
obtained with the C14 and C12 homologues of (alkyldimethylammonio)ethanoate and (alkyldimehtylammonio)1-propanesulfonate: 0.58 nm2 (for CmN1C) and 0.56 nm2 (for CmN3S).20,22 Compared to the flat air-solution interface, the area per molecule is roughly halved when the surfactant is adsorbed onto silica. In the case of C14N1C, the unexpectedly small value of σS/σo (less than 0.5) may be ascribed to the different orientations of the zwitterionic headgroup at both interfaces:15 the headgroup is lying flat at the air-solution interface,20 whereas its arrangement is tilted away from the parallel near the silica surface (note as well that in C12N1C micelles the dipole is oriented perpendicular to the micelle surface8). The chain-length sensitivity of Γmax is an argument against a close-packed bilayer structure of the adsorbed phase for surfactant betaines containing moderately long alkyl chains. The effectiveness of adsorption of sulfobetaine surfactants onto silica is reduced by an increase in the ionic strength of the aqueous phase (Figure 6). Although the presence of electrolyte in the adsorption system induces several new effects, the outcome of decreasing solid-solute interaction seems to dominate over the others, preventing the growth of surface aggregates into more extended structures. The calculated values of σS/σo range between 0.7 and 0.8, being greater for higher ionic strengths. Temperature increase appears to have no effect on the quantity of adsorption in the whole adsorption range (Figures 7 and 8), in spite of the increased exothermicity of the process. This can be explained by the decreased solute-solvent and solid-solvent interactions as the temperature is raised. The former causes the average aggregation number to decrease (as in the bulk phase8). The latter makes the affinity between the surfactant headgroups and the support stronger, thus promoting formation of more extended aggregates. It can reasonably be argued that both opposing effects are small and cancel each other. Simultaneously, less endothermic dehydration of the hydrophobic groups and less endothermic dewetting of the mineral surface result in a more exothermic process of displacement. An increase in the length of the tether chain from n ) 1 to n ) 3 increases the effectiveness of adsorption of carboxybetaine surfactants onto Spherosil XOB015 (Figure 7), the change being relatively small compared to that caused by a longer hydrophobic tail. The same upward tendency in Γmax can be observed in Figure 8 for phenylphosphinatobetaines with decreasing length of the headgroup from n ) 10 to n ) 3. This means that the more hydrophilic the headgroup, the greater is the effectiveness of adsorption. When a few methylene groups are added to the intercharge arm, the cationic and anionic groups become independent of each other and may be randomly located in the interfacial region. Since the electrostatic attraction within the zwitterion diminishes, the ion-dipole bonding with the surface becomes stronger, making the adsorption process more cooperative. With a further increase in the tether length, the steric hindrance between the headgroups increases. A folding of long zwitterionic headgroups and their location in the micelle hydrophobic core (anionic group with a part of the tether chain) do not permit closer packing of the alkyl chains, with a consequent decrease in the surface aggregation number. It should also be noted that the values of Γmax obtained for the three phenylphosphinatobetaines are quite small in comparison with the other surfactants. These molecules are adsorbed on a silica surface which contains some traces of sodium as impurities. The presence of exchangeable sodium cations at the silica-
Micellization and Adsorption of Zwitterionic Surfactants
solution interface33 will disfavor the direct adsorption of zwitterions on the surface. Conclusions The results of calorimetric measurements, applied to micellization of zwitterionic carboxybetaine, sulfobetaine, and phenylphosphinatobetaine surfactants in aqueous solutions and subsequently to adsorption of these molecules on hydrophilic silica surfaces, as well as the observed trends with varying alkyl chain length, tether length (i.e., the distance between charged centers), temperature, and salt concentration, are in good agreement with other experiments. The titration calorimetric technique, which consists of injecting small amounts of a micellar stock solution into calorimetric cell, allows one to determine values of the molar enthalpy of micelle formation. In the adsorption experiment, it is possible to follow evolution of the differential molar enthalpy of displacement with increasing density of surfactant adsorption. The formation of micelles at room temperatures is an endothermic process, displaying behavior typical of what has been termed hydrophobic effect. The positive standard enthalpy of micellization markedly decreases (the process becomes less endothermic) with increase in the length of the hydrophobic tail, temperature, and salt concentration (but only at high ionic strength). The effect of tether length (n) is complex, the value of ∆micho appearing first to increase with n to some maximum (in the neighborhood of n ) 3) and then to decrease with further increase in n. In the literature, the analogous trend in the cmc with increasing n has been ascribed to differences in the hydrophilicity of the headgroup, being a subtle balance between the hydrophobic effect induced by the increasing number of carbon atoms and the hydrophilic effect of the increasing dipole moment.
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Adsorption of surface-active betaines on a hydrophilic silica surface occurs mainly by a combination of ion-dipole bonding (between the zwitterionic headgroups and the ionic surface sites) and hydrophobic bonding mechanisms. At very low surface coverages (usually below 0.1 µmol m-2), the surfactant molecule is individually adsorbed on the surface (lateral adsorbate-adsorbate interactions are negligible), with the quaternary nitrogen group close to the surface and the anionic substituent group away from it. The orientation of the hydrophobic group initially may be parallel to the support; as adsorption progresses, the surfactant tail is displaced from the surface by the hydrophilic headgroup and lateral interactions between adjacent alkyl chains. Individual adsorption has a pronounced competitive character and is generally exothermic. For higher surface coverages, the adsorption of surfactant betaines is characterized by a constant enthalpy of displacement, ∆dplhc. The trends in ∆dplhc with changing temperature, salt content, and molecular structure of the surfactant parallel the corresponding ∆micho changes. The adsorbate structure on a solid surface arises from the competition between the molecular forces that drive the formation of bulk micelles and the constraints of the solid surface (further adsorption of certain molecules with the headgroup in contact with the surface). Strong cooperativity in the adsorption, induced by a decrease in the ionic strength or an increase in the hydrophilic character of the surfactant headgroup, reinforced with a closer packing of longer alkyl chains, leads to the formation of extended surface aggregates (large micelles “flattenedout” on the surface) or even bilayers in the adsorption plateau. LA960926D
