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Adsorption of an Aqueous Mixture of Surfactants on Silica A. Thibaut,† AM. Misselyn-Bauduin,† J. Grandjean,‡ G. Broze,§ and R. Je´roˆme† Center for Education and Research on Macromolecules, Chimie Fine aux Interfacesm University of Lie` ge, Sart-Tilman, B6, 4000 Lie` ge, Belgium, Colgate-Palmolive R&D, 4460 Milmort (Herstal), Belgium Received March 1, 2000. In Final Form: August 1, 2000 The behavior of mixtures of a nonionic surfactant, pentaethylene glycol monodecyl ether (C10E5), and an anionic surfactant, sodium dodecyl sulfate (SDS), was studied at the water-silica interface. In contrast to the C10E5 surfactant, which is adsorbed at the silica-water interface at pH = 6, the adsorption of the anionic surfactant is not observed. In the presence of SDS, the adsorption of C10E5 is severely restricted as result of the formation of mixed micelles in solution. This behavior is consistent with the adsorption free energy, ∆Gads, of C10E5 and the free energy of mixed micellization, ∆Gmix.mic, of SDS and C10E5. Flow microcalorimetry (FMC) results showed that the addition of C10E5/SDS mixtures to silica precoated with C10E5 resulted in the release of the preadsorbed nonionic surfactant; the whole phenomenon is exothermic. Combination of calorimetric data (FMC and isothermal titration calorimetry), self-diffusion coefficients, and the regular solution theory led to the conclusion that the main driving force for the C10E5 desorption was the formation of mixed micelles by unimeric SDS and the released nonionic surfactant.
Introduction The adsorption of surfactants at solid-liquid interfaces plays a crucial role in many important industrial processes, such as detergency, flotation, and painting. The actual formulations used in these applications are complex mixtures of several components, including surfactants, polymers, perfumes, thickeners, dispersants, pigments, etc. The partition of the surface active additives between the bulk solution and the solid surface is essential to the success of the product. The investigation of the adsorption equilibria occurring in these multicomponent systems is a challenge. However, the study of binary mixtures of surfactants, that is, comparatively simple systems, is feasible and could improve the understanding of the adsorption phenomena and contribute to optimize formulations. The solution properties of mixtures of surfactants usually indicates deviation from ideal mixing. Several theories were proposed in this respect,1-5 the most widely used being the pseudo-phase separation model, which is based on the approximation of regular solution (RST).1 In this theory, an interaction parameter, β, is defined as the change in enthalpy when the two surfactants are mixed together. It allows the critical micelle concentration (cmc), the micelle composition, and the unimers concentration to be predicted. The adsorption of mixtures of surfactants at solid-liquid interfaces was also investigated. From the adsorption of mixed surfactants on mineral substrates,6,7 two main situations emerge. In the first case, the two surfactants can be adsorbed onto the substrate,8-10 and the issue of †
Center for Education and Research of Macromolecules. Chimie Fine aux Interfacesm. § Colgate-Palmolive R&D. ‡
(1) Rubingh, D. N In Solution Chemistry of Surfactants; Mittal, K. L., Ed.; Plenum press: New York, 1979; Vol. 1, p 337. (2) Blankschtein, D.; Puvada, S.; Sarmoria, C. Langmuir 1992, 8, 2690. (3) Holland, P. M.; Rubingh, D. N. J. Phys. Chem. 1983, 87, 1984. (4) Holland, P. M. In Phenomena in Mixed Surfactant Systems; Scamehorn, J. F., Ed.; ACS Symposium Series; Washington, 1986; Vol. 311, p 102. (5) Rodenas, E.; Valiente, M.; Villafruela, M. J. Phys. Chem. B 1999, 103, 4549.
their competition for adsorption depends on the pH. The structure and composition of the adsorbed layer is the fingerprint of the relative affinity of the two surfactants for the surface. As a rule, the specific interactions of the surfactants with the solid surface dominate the properties of the surfactants mixture in solution. In the second case, only one surfactant is known for adsorption onto the surface. In this case, the adsorption pattern depends strongly on the system.9,11-18 In this study, the adsorption of binary mixtures of a nonionic surfactant, pentaethyleneglycol monodecyl ether (C10E5), and an anionic surfactant, sodium dodecyl sulfate (SDS), will be investigated at the silica-water interface. The adsorption will be studied by flow microcalorimetry and measurement of the adsorption isotherms. The dependence of the adsorption on the monomer concentration will be discussed on the basis of the regular solution theory and data of self-diffusion coefficients. The sequential addition of the surfactants will also be compared with the adsorption of the premixed surfactants. Experimental Section Materials. The silica used in this study was the porous silica X015 M supplied by Biosepra, France. According to the manu(6) Harwell, J. H.; Scamehorn, J. F. In Mixed Surfactant Systems; Ogino, K., Abe, M., Eds.; Surfactant Science Series; Marcel Dekker: New York, 1993; Vol. 46, p 263. (7) Somasundaran, P.; Krishnakumar, S. Colloids Surf. A 1997, 123-124, 491. (8) Penfold, J.; Staples, E. J.; Tucker, I.; Thompson, L. J. Langmuir 1997, 13, 6638. (9) Penfold, J.; Staples, E. J.; Tucker, I.; Thompson, L. J.; Thomas, R. K. Int. J. Thermophys. 1999, 20, 19. (10) Schwuger, M. J.; Smolka, H. G. Colloid Polym. Sci. 1977, 255, 589. (11) Somasundaran, P.; Snell, E. D.; Fu, E.; Q, X. Colloids Surf. 1992, 63, 49. (12) Somasundaran, E.; Fu, E.; Xu, Q. Langmuir 1992, 8, 1065. (13) Huang, Z.; Yan, Z.; Gu, T. Colloids Surf. 1989, 36, 353. (14) Huang, L.; Maltesh, C.; Somasundaran, P. J. Colloid Interface Sci. 1996, 177, 222. (15) Capovilla, L.; P, L.; Reverdy, G. Langmuir 1991, 7, 2000. (16) Meguro, K.; Adachi, T.; Fukunishi, R.; Esumi, K. Langmuir 1988, 4, 1160. (17) Gao, Y.; Yue, C.; Lu, S.; Gu, W.; Gu, T. J. Colloid Interface Sci. 1984, 100, 581. (18) Scamehorn, J. F.; Schechter, R. S.; Wade, W. H. J. Colloid Interface Sci. 1982, 85, 494.
10.1021/la000302e CCC: $19.00 © 2000 American Chemical Society Published on Web 10/31/2000
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facturer, the particle size covered the 39-112 µm range, and the average pore size and the specific area were 1260 Å and 30.8 m2, respectively. SDS from BDH showed no minimum in the surface tension vs concentration curve, thus supporting high purity. The nonionic surfactant, C10E5, was purchased from Flukka and certified to be monodisperse. Methods. Surface Tension Measurements. Surface tension was measured at 298 K with an automatic Kru¨ss K12 tensiometer, equipped with a platinum Wilhelmy plate, to measure the critical mixed micelle concentration (cmc*) for SDS/C10E5 mixtures of different molar ratios, R. Glassware was cleaned carefully by dipping for several hours in 10% H2SO4 solution and then washed thoroughly with deionized water. The cleanness of the glassware was checked by the surface tension measurement of highperformance liquid chromatography (HPLC) grade water, which also was used for the preparation of all the surfactant aqueous solutions. Each surface tension measurement was repeated three times and accepted when the data did not differ from each other by more than 0.25 mN/m. The cmc* was determined at a constant surfactant composition from the surface tension measured for ca. 12-14 solutions of different concentrations. These solutions were prepared by dilution of a 40 mM stock solution of the surfactants mixture of well-defined composition. NMR Experiments. Self-diffusion coefficients, Ds, were measured at 298 K by the pulsed field gradient NMR technique19,20 with a Bruker AM 300WB spectrometer operating at the proton Larmor frequency of 300 MHz. The basic sequence was used with pulsed field duration, δ, of 6 ms and a time interval, ∆, between the two gradient pulses of 22 ms. The echo attenuation, A, was recorded as a function of the gradient amplitude, g, and calibrated with octanol assuming that Ds ) 1.9 10-10 m2 s-1 at 20 °C.21 The signal intensity decreased exponentially, as predicted by the theory:
A ) A0 exp(-γ2δ2g2Ds(∆ - δ/3)
(1)
where γ is the proton gyromagnetic ratio. In SDS, A was measured at different gradient amplitudes, g, for the signal of the highest intensity at 1.3 ppm (side chain protons). The other proton signals, particularly the 4.05 ppm peak (protons in R position of the SO4 group), were less convenient because of the poor signal-to-noise ratio under the experimental conditions of this study. In C10E5, the oxyethylenic protons (δ ) 3.7 ppm) were very well suited to the determination of Ds. Self-diffusion coefficients were calculated by fitting 13 experimental data by eq 1. All the reported values were the average of three independent measurements. Adsorption Isotherms. The adsorption isotherms were measured according to the classical total concentration depletion method at 298 K. Solution (10 mL) of the two surfactants at the desired concentrations was gently shaken overnight with 100 mg of silica. At the equilibrium, the solid was separated by centrifugation (5000 rpm; 20 min). Concentration of the C10E5 nonionic surfactant in the supernatant was analyzed by gas chromatography (GC), whereas the SDS concentration was measured by the Epton method.22 All the reported values were the average of three independent measurements whose standard deviation never exceeded 2%. Flow Microcalorimetry Measurements. The Microscal Flow Microcalorimeter was designed to quantify thermal effects that result from the interaction of a solid surface with a compound in solution. Actually, all the heat effects (heat of dilution, heat of interaction of all the constituents) occurring in the cell are recorded. Moreover, the enthalpy of adsorption should be strictly designated as ‘enthalpy of replacement’, ∆Hrepl, because the adsorption requires the displacement of solvent molecules from the solid surface by molecules of solute.23,24 (19) Stejskal, E.; Tanner, J. E. J. Chem. Phys. 1965, 42, 288. (20) Stilbs, P. Prog. Nucl. Magn. Reson. Spectrosc. 1987, 19, 1. (21) Herden, H.; Karger, J.; Pfeifer, H.; Kube, C.; Schollner J. Colloid Interface Sci. 1992, 152, 281. (22) Epton, S. R. Trans. Faraday Soc. 1948, 44, 226. (23) Rouquerol, J. Pure Appl. Chem. 1985, 57, 69. (24) van Os, N. M.; Haandrikman, G. Langmuir 1987, 3, 1051.
Figure 1. Sketch of the flow microcalorimeter. The Microscal Flow Microcalorimeter was detailed elsewhere.24,25 The apparatus used in this work is schematized in Figure 1. It consists of two main parts: the calorimetric cell and a refractometer downstream to the cell. The microcalorimeter cell (0.17 mL in volume) delimited by an inlet and an outlet Teflon connectors was fitted in a metal block within a draft-proof enclosure. It was filled with a wellknown amount of silica. Water or aqueous solutions were passed through the cell with a syringe pump at a constant flow rate (1.5 mL/h). The effluent flew through the downstream refractive index detector. Temperature changes in the microcalorimeter cell were detected by two thermistors adjacent to the cell and transmitted to a chart recorder and a computer, as a function of time. These changes in temperature were recorded as peaks, whose area was proportional to the total heat released or absorbed. The calibration experiment was performed by replacing the commonly used outlet connector by the calibration outlet connector in which an electrical coil is encapsulated. The calibration experiment consisted of flowing water through a bed of adsorbent (silica) until equilibration (stable baseline). Variable controlled amounts of heat were produced at the thermistors by passing an electrical current through the calibration coil during a given period. Recorded peak areas were related linearly to the heat produced by the coil allowing the calibration factor to be calculated. A calibration factor had to be calculated for each sensitivity (different bridge supply and amplifier setting) and each fluid flow rate through the flow microcalorimetry (FMC) cell. The reported heat of displacement was based on two or three independent experiments reproducible with standard deviations in the heat of displacement of about 2-5%. The amount of compound adsorbed onto silica was calculated as follows. The concentration of the effluent was monitored as a function of time. The same experiment was then repeated with the cell filled with Teflon (polyperfluoroethylene) (blank run) instead of silica. Because of the very low surface area (0.1 m2/g) and surface tension of Teflon, it was assumed that no adsorption occurred on the solid surface. The amount of solute retained on silica was then calculated as the difference between the refractometer signals for the two experiments. Typical adsorption experiments performed in this work consisted of four steps. The first step was the setup of the baseline by wetting the silica powder (= 60 mg) with distilled water (overnight), supplied by one of the syringe pumps at a flow rate of 1.5 mL/h. In the second step, a C10E5 solution (13.2 or 2.6 mM) in distilled water was supplied by the second syringe pump at (25) Groszek, A. J. Thermochim. Acta 1998, 312, 133.
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the same flow rate and flowed through the silica bed until the adsorption of C10E5 onto silica was at equilibrium (=180 min). In the third step, a solution mixture of C10E5 and SDS in distilled water was flowed at the same flow rate and for =120 min. In the final step, the flow was switched to distilled water, and the desorption process was monitored. The same experiment was repeated with Teflon instead of silica. All the experiments were performed at 298 K. Thermometric Measurements. An OMEGA, ITC, microcalorimeter from Microcal was used to measure the micellization enthalpy of C10E5. A concentrated solution of the surfactant (28.65 mM) was added stepwise (25 steps; 4 µL each) with a 250-µL microsyringe to water in the 1.325-mL calorimeter cell under constant stirring. The calorimeter measured the difference in the heat flux into or out of the titration cell with respect to a reference cell that contained water. The instantaneous data (in microwatts) were recorded by the computer and integrated over time, so giving the total heat evolved for each injection in kilojoules per mole of C10E5. The enthalpy of mixing of SDS with C10E5 was measured by adding stepwise a solution of SDS and C10E5 (25 steps; 4 µL each) in the 1.325-mL calorimeter cell containing a C10E5 solution (13.2 mM) in distilled water under constant stirring. The thermogram was recorded, and the heat evolved for each step was cumulated throughout the 25 steps. The same experiment was repeated for a different SDS/C10E5 ratio, the C10E5 concentration being kept constant at 13.2 mM. The reported values were the average of two independent measurements. Pseudo-phase Separation Model; Regular Solution Theory (RST). In a previous paper,26 the RST was shown to be appropriate to the analysis of binary SDS/C10E5 mixtures. The interaction parameter, β, was interpreted as the excess heat of mixing. It was determined from the mixed cmc and the cmcs of each surfactant by the iterative solution of eq 21:
[
β ) ln
][
] [
][ ]
(1 - R)cmc* 1 R.cmc* 1 ) ln cmc1x1 (1 - x12) cmc2(1 - x1) x12
(2)
where R is the mole fraction of SDS in the mixture, cmc1 and cmc2 are the cmcs for each surfactant, cmc* is the cmc of the binary surfactant mixture, and x1 is the mole fraction of surfactant 1 in the mixed micelles. The unimer concentration was derived by the iterative solution of eqs 3 and 41:
C1M )
-(C - ∆) + x(C - ∆)2 + 4RC∆ f2.cmc2 2 -1 f1.cmc1 M
C2
(
(
)
)
C1M ) 1f cmc2 f1cmc1 2
(3)
(4)
where f1 ) exp β(1 - x)2; f2 ) exp βx2; ∆ ) f2cmc2 - f1cmc1, and C is the total concentration of the surfactants 1 and 2.
Results and Discussion Adsorption of C10E5. Figure 2 shows the adsorption isotherm for C10E5 as measured by the total concentration depletion method. The isotherm has a typical S-shape, which supports the idea that the adsorption is basically cooperative. Just below the critical surface aggregation concentration (csac), no significant adsorption occurs, whereas above it, the extent of the adsorption increases dramatically. Slightly above the cmc, the adsorption levels off and a plateau is reached corresponding to a maximum coverage, Γmax, of 5.3 µmol/m2. (26) Misselyn, A.-M.; Thibaut, A.; Grandjean, J.; Broze, G.; Jerome, R. Langmuir 2000, 16, 4430.
Figure 2. Adsorption isotherm of C10E5 on silica (standard deviation ) 2%).
Figure 3. (a) The two-step adsorption of nonionic surfactant on a negative solid surface. (b) Equilibrium reaction corresponding to K20.
These experimental data fit very well the adsorption isotherm predicted by Zhu and Tiren27 (solid line, Figure 2) from the law of mass action for the two-step mechanism schematized in Figure 3a. This mechanism comprises the adsorption of unimers (anchors) at low concentration followed by aggregates (admicelles or hemimicelles) of an average aggregation number, n, at higher concentration. The adsorption equation proposed for this two-step model is reported in eq 5.
Γ)
ΓmaxK1C[1/n + K2Cn-1] 1 + K1C(1 + K2Cn-1)
(5)
where Γmax is the maximum surface excess concentration, C is the bulk surfactant concentration, K1 and K2 are the equilibrium constants for the first and second adsorption steps, respectively. K1, K2, and n, characteristic of the adsorption of C10E5 onto silica, were calculated from eq 5 and found equal to 1800, 1.3 × 1049, and 16, respectively. K1 is expectedly much smaller than K2, which indicates that the first equilibrium shown in Figure 3a is far less favorable than the second one. The average number of surfactants per aggregates (n ) 16) is actually the number of surfactants that occupy the same surface area as one surfactant adsorbed flat on the surface during the first step. The adsorption process is driven thermodynamically by the change in free-energy variation, according to the general eq 6
∆Gads ) ∆Hads - T∆Sads ) ∆Gads ) -RTlnK (6) where K is the equilibrium constant of the adsorption process. (27) Zhu, B. Y.; Tiren, G. Adv. Colloid Interface Sci. 1991, 37, 1.
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Table 1. Thermodynamic Parameters for Micellization of C10E5 in Water and for the Adsorption on Silica cmc (mM)
∆Gmic (kJ/mol)
∆Hmic (kJ/mol)
csac (mM)
Γmax (µmol/m2)
∆Gads (kJ/mol)
∆Hrepl (kJ/mol)
0.69
-18
+0.85
-0.35
5.3
-18.7
+0.6
Figure 4. Enthalpy vs C10E5 concentration measured during the 25-step addition of 100 µL of C10E5 (28.65 mM) into 1.325 mL of water. The ∆Hmic is calculated by subtracting the initial (i) from the final (f) enthalpy indicated by the vertical arrow.
The value of K1 extracted from the first part of the adsorption isotherm (Figure 2) is not accurate enough to be considered in eq 6. Rather, a constant K20, which reflects the ability of one surfactant molecule to adsorb on one or more preadsorbed molecules, was calculated according to eq 7 (Figure 3).
K20 ) (K2)1/(n-1)
(7)
From K2 ) 1.3 × 1049, and n ) 16, K20 ) 1880 (eq 7), and ∆Gads ) -18.7 kJ/mol (eq 6) were calculated. To characterize further the adsorption process, the enthalpy of replacement, ∆Hrep, was measured for C10E5 by means of FMC at such a concentration (2.6 mM) that Γ was maximum (5.3 µmol/m2). This value was close to 3 mJ/m2, thus 0.6 mJ/µmol of C10E5 adsorbed (Table 1). In the plateau region, the adsorption process is thus slightly endothermic, which indicates that the adsorption of C10E5 is driven entropically at high surfactant concentration. This type of behavior was reported previously by Seidel et al.,28 in the adsorption of the nonionic C8E4 surfactant on silica. These authors reported that the adsorption of this surfactant was slightly exothermic at concentrations below the csac, although it was endothermic beyond this critical concentration. The exothermic process resulted from the adsorption of individual surfactant molecules to the surface by hydrogen bonding. At higher surfactant concentration, the adsorbed molecules formed aggregates and the replacement process was endothermic.29,30 The enthalpy of replacement, ∆Hrep, found for C10E5 is very close to the enthalpy of micellization, ∆Hmic, that was measured by titration calorimetry (Figure 4, Table 1). The close agreement between ∆Hrep and ∆Hmic is calorimetric evidence that the formation of aggregates at the silica-water interface is driven by the same hydrophobic interactions that cause the formation of micelles in solution and trigger the release of water molecules. (28) Seidel, J.; Wittrock, C.; Kohler, H.-H. Langmuir 1996, 12, 5557. (29) Partyka, S.; Lindheimer, M.; Faucompre, B. Colloids Surf. A 1993, 76, 267. (30) Narkiewicz-Michalek, J.; Rudzinsky, W.; Keh, E.; Partyka, S. Colloids Surf. 1992, 62, 273.
Figure 5. Adsorption isotherms of C10E5 on silica from C10E5 and SDS mixtures at SDS molar fraction of 0 (9), 0.1 (2), and 0.5 (O), respectively (standard deviation ) 2%).
This release is responsible for the increase of entropy which is the driving force for both the surfactant micellization and the surfactant adsorption on silica.28-30 Adsorption of C10E5 in the Presence of SDS. Adsorption Isotherms. Because of repulsive electrostatic interaction, the adsorption of SDS on silica is negligible compared with the adsorption of the nonionic surfactant. Moreover, as reported previously, C10E5 interacts with SDS in aqueous solution, and the average interaction parameter, β, calculated by RST lies between ca. -2 and -3.3.26 It is worth investigating to which extent the anionic surfactant can perturb the adsorption profile of the nonionic surfactant at the silica-water interface as a consequence of the mutual association of these two compounds. The influence of SDS on the adsorption of C10E5 was studied by measuring the adsorption isotherm of the nonionic surfactant in the presence of SDS, the SDS molar ratio, β, in the surfactants mixture being 0.1 and 0.5, respectively. Figure 5 shows that the anionic surfactant strongly affects the adsorption isotherm of pure C10E5, because the addition of 10 mol % of SDS to C10E5 results in a 3-fold decrease of Γmax for C10E5, from 5.3 to 1.7 µmol/m2. However, C10E5 seems to aggregate on silica at lower C10E5 concentrations than for pure C10E5. When SDS and C10E5 are used in equimolar amounts, no clear signature for the surfactant-adsorbed aggregates is observed, and Γmax does not exceed one tenth of the value for pure C10E5. Furthermore, using the total concentration depletion method, the amount of SDS adsorbed on silica under the same experimental conditions (R ) 0.1 and 0.5) was very low even in the presence of the nonionic surfactant. These observations show that, when SDS and C10E5 are mixed together in the presence of silica, the nonionic surfactant adsorption is severely restricted, because formation of mixed micelles in the bulk solution is more favorable than the adsorption of the surfactants. As suggested by the two-step adsorption model (Figure 3a), the adsorption of surfactant proceeds through the adsorption of the unimer, the micelles are not directly adsorbed. The observed decrease of the C10E5 adsorption therefore results from the reduction of the C10E5 unimer concentration as a consequence of mixed micelle formation.
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Figure 6. C10E5 unimer concentration (9) and C10E5 adsorption (0) versus, R, the mole fraction of SDS; [C10E5] ) 2.6 mM.
Thibaut et al.
observed that the adsorbed layer from sodium p-octylbenzenesulfonate/octaethylene glycol mono-n-dodecyl ether also contained the two surfactants. On the basis of the aforementioned observations and comments, the adsorption of premixed surfactants is controlled by two opposite tendencies, that is, the enhanced tendency of mixed micelles to be formed and the formation of a mixed adsorbed layer at the solid-liquid interface. Thus, the favorable formation of mixed micelles in solution or mixed adsorbed layer at the solid surface results from the interplay of both the change of free energy of mixed micellization, ∆Gmix.mic, and of surfactant adsorption, ∆Gads. Similar to eq 8a for the change of enthalpy upon mixed micellization, ∆Hmix.mic:31 mic
∆Hmix.mic ) x1∆Hsurf1
mic
+ (1 - x1)∆Hsurf2
+ HE (8a)
∆Gmix.mic can be expressed by eq 8b: mic
∆Gmix.mic ) x1∆Gsurf1
+ (1 - x1)∆Gsurf2mic + GE (8b)
where x1 is the molar fraction of surfactant 1 in the mixed mic micelles, ∆Gmic surf1 and ∆Gsurf2 are the free energy of micellization of surfactants 1 and 2, respectively, and GE is the excess free energy of micellization related to the mixing of the two surfactants. As already mentioned, SDS/C10E5 mixtures are in agreement with the RST. According to this theory, the excess entropy, SE, is negligible.1 Accordingly, GE may be expressed by eq 9,
GE ) HE ) x(1 - x)βRT
Figure 7. Critical micelle concentrations (cmc*) for C10E5 and SDS mixtures of various compositions.
This explanation is confirmed by Figure 6, which compares the adsorbed amount of C10E5 and the C10E5 unimer concentration in solution as a function of the mole fraction of SDS, R. The unimer concentration of C10E5 was calculated from the RST (β ) -2.7) and cmc* at different compositions (Figure 7). The amount of C10E5 adsorbed was measured by the total concentration depletion method. The SDS concentration was changed at each mixture composition, with the C10E5 concentration kept constant at 2.6 mM, a concentration at which the surface coverage is maximum in the absence of SDS. Figure 6 shows a parallel decrease in the unimer concentration and the amount of C10E5 adsorbed when the surfactant mixture contains relatively more SDS, which indicates again the mixed micelle formation in aqueous solution. These observations, however, do not preclude that the layer of adsorbed surfactants is a mixture of the anionic and the nonionic surfactants, although the amount of SDS should be low because the surfactant was not detected by the total concentration depletion experiments. Data published about closely related systems support this suggestion. Indeed, Penfold et al.9 recently studied the adsorption of SDS/C12E6 mixtures on silica by specular neutron reflection, which allows the composition and the structure of the adsorbed surfactants to be analyzed. The adsorbed layer contained both the anionic and nonionic surfactants. Similarly, Somasundaran et al.11
(9)
where x is the micellar molar composition and HE is the excess enthalpy. The interaction parameter, β, being negative for the system under consideration (β ) -2.7), GE will contribute to make ∆Gmix.mic more negative. ∆Gmic values for C10E5 and SDS are -18 and --23.9 kJ/mol,32 respectively. According to eq 8b, ∆Gmix.mic for surfactant mixtures of different compositions is expected to range between -18 + GE and -23.9 + GE kJ/mol. Moreover, ∆Gads for C10E5 is -18.7 kJ/mol (Table 1) and SDS can only contribute to make it more positive because ∆Gads for SDS is close to zero. Therefore, the mixed micellization of SDS/C10E5 mixtures in solution is expected to be more favorable than their adsorption on silica. Order of Addition. In addition to the adsorption of premixed C10E5 and SDS onto silica, it is worth considering the effect that the sequential addition of the surfactants might have on the silica bed. For this purpose, the FMC was used, which allows one surfactant to be preadsorbed on silica (C10E5 in this case), followed by the addition of nonadsorbing SDS. Figure 8 compares the enthalpy and the refractometry profiles observed when a mixture of SDS/C10E5 is flowing through a bed of silica precoated with the nonionic surfactant. A blank run was also performed with Teflon instead of silica. In this experiment, the amount of C10E5 preadsorbed on silica was 4.25 µmol/m2, the maximum coverage obtained by flowing a 13.2 mM surfactant solution through the silica bed. This value is ca. 80% of the value measured by static experiments. This difference can be explained by transport limitations in the flow system as suggested (31) Hey, M. J.; MacTaggart, J. W. J. Chem. Soc., Faraday Trans. 1 1985, 81, 207. (32) Mukerjee, P. Adv. Colloid Interface Sci. 1967, 1, 241.
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Figure 10. Heat of mixing of C10E5 and SDS/C10E5 mixtures measured by FMC and by ITC in relation to the SDS concentration of the mixture solution; ([C10E5] ) 13.2 mM). Figure 8. Refractive index (upper) and calorimetric (lower) profiles when a SDS/C10E5 solution is flown through a bed of C10E5-precoated silica (solid line) and Teflon (dashed line), respectively.
Figure 9. (a) Enthalpy of replacement when C10E5 precoated silica is eluted by a SDS/C10E5 solution ([C10E5] ) 13.2 mM) and variable [SDS]; (b) amount of C10E5 released under the same conditions; flow rate ) 1.5 mL/h.
in the scientific literature.33 A solution of SDS/C10E5 mixture rather than SDS was flowed through precoated silica to prevent C10E5 from being partly desorbed because of partitioning between the solid and the liquid phases. The run through precoated silica (Figure 8) shows a large exotherm, together with a maximum in the refractive index of the effluent before reaching a plateau. So, for a certain time, the surfactant concentration is higher in the effluent than in the injected solution. This observation indicates that C10E5 is desorbed in the presence of SDS. Subsequent addition of water does not lead to further desorption. The same experiment was repeated at different SDS concentrations, while keeping the nonionic surfactant concentration constant at 13.2 mM to check whether C10E5 is desorbed at any molar fraction and to understand the origin of the exotherm. Figure 9 shows how the heat of replacement and the amount of the nonionic surfactant released depends on the SDS concentration. The amount of C10E5 desorbed is independent of the composition of the SDS/C10E5 mixtures. More than 80% (33) Kwok, W.; Nasr-El-Din, H. A.; Hayes, R. E.; Sethi, D. Colloids Surf. A 1993, 78, 193.
of the preadsorbed nonionic surfactant is actually desorbed (ca. 3.45 µmol/m2) forming mixed micelles with SDS of different compositions. The enthalpy change at a SDS concentration smaller than 10 g/L shows a linear dependence on this concentration, although the slope decreases at higher SDS concentrations. The enthalpy changes observed when SDS is added can be explained by: (a) desorption of the nonionic surfactant from silica, (b) dilution of the SDS solution, and (c) the interaction between the desorbed C10E5 and the species in the flowing SDS/C10E5 solution. The first two reasons may be disregarded. Indeed, if the amount of the nonionic surfactant released is constant, the enthalpy of desorption should also be constant, which is not the case. Moreover, this contribution must be negligible, because the enthalpy of adsorption was very small. Because the dilution enthalpy of SDS, that is, the demicellization enthalpy, is nearly athermal at room temperature,34 dilution of SDS cannot explain the large enthalpy change observed in Figure 9. The most probable origin of the exotherm is the heat of mixing between the released nonionic surfactant with the flowing SDS/C10E5 solution. Isothermal titration calorimetry was performed to confirm this hypothesis. The heat of mixing of a 13.2 mM C10E5 solution with SDS/ C10E5 mixtures of the same compositions as used in the FMC experiments was measured. Figure 10 shows that the enthalpy changes recorded by FMC and by isothermal titration calorimetry (ITC) follow an evolution similar to the SDS concentration, that is, a linear dependence at low SDS concentration followed by a slower increase at higher concentration. It may thus be stated that the heat of exchange observed by FMC is intimately related to the heat of mixing of the nonionic surfactant released from the surface with the SDS/C10E5 elute. Nevertheless, the comparisons are only qualitative because the concentration of the nonionic surfactant released from the silica surface cannot be measured within the FMC cell. In a further step, the unimer concentration of SDS and C10E5 was estimated in the SDS/C10E5 elute. A previous work reported on the measurement by NMR of the selfdiffusion coefficients of the premixed SDS/C10E5 surfactants. From these experimental data, it is possible to extract the concentration of the unimers formed by each surfactant.26 The concentration of the C10E5 unimers is too low to be measured accurately, in agreement with the values calculated by RST, which decrease rapidly when the SDS (34) Birdi, K. S. Colloid Polym. Sci. 1983, 261, 45.
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Figure 11. (a) ∆H measured by FMC (9) when the SDS/C10E5 eluate is flowed through C10E5-precoated silica ([C10E5] ) 13.2mM and variable [SDS]); flow rate ) 1.5 mL/h; (b) SDS unimer concentration as a function of the SDS concentration of the SDS/C10E5 solution ([C10E5] )13.2 mM) measured by NMR (4) and calculated by RST (solid line).
molar fraction of the binary mixture is increased (Figure 6). Therefore, C10E5 unimers cannot account for the exotherm observed in Figure 8. In contrast, the SDS unimer concentration was calculated from the NMR selfdiffusion coefficients as shown in Figure 11. Figure 11 compares the SDS unimer concentration measured by NMR and calculated by RST with the enthalpy change measured by FMC. The good agreement between these data indicates that the enthalpy change measured by FMC is directly related to the concentration of SDS unimers. Therefore, the observed enthalpy change results from the binding of the SDS unimers in the elute with C10E5 molecules desorbed from silica. So, the driving force for the desorption of the nonionic surfactant is the formation of SDS/C10E5 mixed micelles. 4. Conclusions This article reports on the adsorption of a mixed surfactant system, that is, SDS/C10E5, at the silica-water
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interface. These surfactants form mixed micelles in aqueous solution; their mixing is nonideal as a result of the moderately strong interaction of the headgroups within these micelles. The regular solution theory accounts very well for these interactions and can predict the concentration of the unimers and the mixed micellar composition at various concentrations of the individual surfactants.26 In the presence of silica onto which only C10E5 is spontaneously adsorbed, FMC experiments and adsorption isotherms have shown that mixed micellization in solution is more favorable than adsorption of C10E5 or SDS/C10E5. This behavior is consistent with the free energy of mixed micellization, ∆Gmix.mic, of SDS and C10E5, which is expected to be more negative than the free energy of C10E5 adsorption, ∆Gads. When C10E5-precoated silica is eluted by C10E5/SDS mixtures, the preadsorbed nonionic surfactant is released and an exothermic reaction occurs. Combination of calorimetric data (FMC and isothermal titration calorimetry) with measurements of self-diffusion coefficients and the regular solution theory leads to the conclusion that the main driving force for the C10E5 desorption is the propensity of unimeric SDS to form mixed micelles with the released nonionic surfactant. Nevertheless, even if the nonionic surfactant adsorbed amount decreases with SDS concentration it cannot be excluded that the remaining layer is a mixture of both anionic and nonionic surfactants.
Acknowledgment. A.T., A.-M.M.-B., and R.J. are much indebted to the “Services Fe´de´raux des Affaires Scientifiques, Techniques et Culturelles” for support to CERM in the frame of the “Poles d’Attraction Interuniversitaires: PAI 4/11”. They are also grateful to ColgatePalmolive for a fellowship to A.T. LA000302E
