500
Langmuir 1987, 3, 500-511
cellar aggregates leading to incorrect values of NWg.For this reason the system Ru(bpy):+/Fe(CN):has to be avoided for the characterization of inverse-micellar aggregates. If dynamic quenching occurs, the efficiency of a probe/quencher system is related to the quenching rate constant of the system in the micellar aggregates. For small aggregates (NWg< 80) all probe/quencher combinations proposed are efficient. Because of the small micellar volume, the probability of collision of probe and quencher (k,) is sufficiently high for all systems proposed. With PSA as probe at these small R values, more information is available compared to PTSA as probe. The long fluorescence lifetime of PSA allows the measurement of the relatively low k, of these aggregates (k, < lo9 M-I s-l1. This is not possible with the short-living PTSA. For the large aggregates (Nqg> 80), PTSA with negative quencher is the only effective system for AOT in n-hexane. In AOT reverse-micellar aggregates, both PTSA and I- or SCN- are repelled from the interphase. The high k,, for these probe/quencher systems results from the fast diffusion in the "free" water of the micellar water pool. This is in contrast to a slower diffusion of positive quenchers which are attracted to the interphase of AOT. Therefore, an efficient probe/quencher system must be chosen depending on the surfactant system studied and more precisely the charge of the polar headgroup.
The use of fluorescence quenching for the characterization of reverse-micellar aggregates is limited by the dynamics of those aggregates. In the case of k,[M] > k,,, no correct information for N and k, can be obtained from the fluorescence decay of %e probes in these micellar aggregates. For an even higher exchange rate, k,[M] >> kqm,the fluorescence decay of a probe is one exponential. Only values for k, can be obtained from these decay curves. There is good agreement between the Naggmeasured with fluorescence quenching and those reported in the literature, determined with alternate methods (Figure 6). The rate constants for intermicellar exchange obtained with fluorescence quenching are of the same order of magnitude as those measured by Pileni et al.25but are larger than those reported by R o b i n ~ o n . ~ " ~ ~
Acknowledgment. Financial support by IWONL-Agfa, the FKFO, the University Research Fund, and the ERO Research Fund are gratefully acknowledged. A.V. is indebted to IWONL for a fellowship. Registry No. AOT, 577-11-7; PTSA, 108292-62-2;PSA, +, Fe(CN)63-,13408-62-3; I-, 59323-54-5; Ru(b ~ y ) ~ *15158-62-0; 20461-54-5; SCN-, 302-04-5; Cu, 7440-50-8;hexane, 110-54-3. (25) Pileni, M. P.; Furias, J. M. In Reverse Micelles: Biological and TechnologicalReleuance;Luisi, P. L., Straub, B. E., Eds.; Plenum: New York, pp 175-181.
Effects of Counterions on Surfactant Surface Aggregates at the Alumina/Aqueous Solution Interface Daryl Bitting and Jeffrey H. Harwell* Institute for Applied Surfactant Research and School of Chemical Engineering and Materials Science, The University of Oklahoma, Norman, Oklahoma 73019 Received November 18, 1985. I n Final Form: January 15, 1987 The adsorption of Li', Na+, K", and Cs+ salts of dodecyl sulfate on alumina is studied as a function of pH, counterion type, and counterion concentration in order to elucidate the role of counterions in the formation of surfactant aggregates at the aqueous/mineral oxide interface. Numerous similarities to specific counterion effects in micelle formation are observed, but there are unexpected findings. The degree of counterion binding on adsorbed surfactant aggregates is estimated and appears to be significantlyhigher than bindings on micelles. Other phenomena presented and discussed include counterion concentration induced aggregate formation by an ionic surfactant on a like-charged surface and the adsorption of simple, monovalent counterions between the adsorbed surfactant aggregate and the mineral surface. All of these phenomena are explained by a simple picture of the role of counterions in aggregate formation which does not depend on specific chemical interactions but does invoke the surface complexation of simple cations with the oxide surface. It is found that the counterion giving the highest adsorption of surfactant at a given surfactant and added electrolyte concentration depends on both pH and fractional surface coverage. This effect is hypothesized to result from a combination of steric and surface complexation considerations. Introduction Surfactant adsorption from aqueous electrolyte solutions onto mineral oxide surfaces has received considerable attention. Reasons for this include the usefulness of surfactants in froth flotation, the growing importance of surfactant adsorption in new separations processes, the role of surfactant adsorption in detergency, and the continued interest in enhanced oil recovery by both low-tension surfactant flooding and surfactant-assisted water flooding. Surfactant adsorption behavior is complicated by the multicomponent nature of these systems, the importance
* To whom correspondence should
be addressed.
0743-7463/87/2403-0500$01.50/0
of the structure of water in surfactant aggregation, and the complex interactions which occur between surfactant, surfactant aggregates, counterions, potential-determining ions, and mineral oxide surface groups. Thus, there remain important aspects of surfactant adsorption on mineral oxides that have received inadequate attention; the interaction of counterions with adsorbed aggregates is one such aspect. Numerous aspects of surfactant behavior in these systems are widely accepted as established. Most workers regions divide typical system isotherms into three or (1) Somasundaran, P.; Fuerstenau, D. W. J.Phys. Chern. 1966, 79,90.
0 1987 American Chemical Society
Counterion Effect on Surfactant Surface Aggregates
IV
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EQUILIBRIUM CONCENTRATION(pM)
Figure 1. Schematic of a typical surfactant adsorption isotherm from aqueous solution onto a mineral oxide.
(Figure 1). Region I corresponds to low-surface coverage by surfactants and an absence of indication of surfactant aggregate formation on the surface. The region I/region I1 transition occurs at the onset of surfactant aggregate formation on the mineral oxide surface,14 as indicated by the sharp increase in the slope of the i ~ o t h e r m .These ~~ aggregates form locally on the surface, i.e., by patchwise a d ~ o r p t i o nthey ; ~ are referred to as hemimicelles' or admicelless to emphasize the micelle-like aspects of their structure and b e h a v i ~ r . ~ The ~ ~concentration ,~ at which the region I/region I1 transition occurs is referred to as the hmc (hemimicelle concentration),l the cac (critical admicelle concentration),8 or the chmc (critical hemimicelle con~entration).~ Region I11 begins where the slope of the adsorption isotherm starts to decrease; this decrease is believed to arise from either the electrostatic repulsion of ions from the interface because of a change in the sign of the surface charge1J2 or from the distribution of patch adsorption e n e r g i e ~ . The ~ , ~ region III/region IV transition occurs either at the critical micelle concentration (cmc) or upon completion of bilayer coverage of the surface;2sssome surfactant/mineral oxide systems exhibit either an adsorption plateau or an adsorption maximum above the cmc,1@12but this phenomenon was not observed in this study. ~
~~
(2) Scamehorn, J. F.; Schechter, R. S.; Wade, W. H. J. Colloid Interface Sci. 1982, 85, 463. (3) Giles, C. H.; MacEwan, T. H.; Nakhwa, S. N.; Smith, D. J . Chem. SOC. 3973. ... 1960. . (4) Bisio, P. D.; Cartledge, J. G.; Keesom, W. H.; Radke, C. J. J. Colloid Interface Sci. 1980, 78, 225. (5) Gaudin, A. M.; Fuerstenau, D. W. Trans. Am. Inst. Min., Metall. Pet. Eng. 1955,202, 958. (6) Gaudin, A. M.; Fuerstenau, D. W. Trans. Am. Inst. Min., Metall. Pet. Eng. 1955, 202, 66. (7) Goujon, G.; Cases, J. M.; Mutaftschiev, B. J. Collid Interface Sci. 1976, 56, 587. (8) Harwell, J. H.; Hoskins, J. C.; Schechter, R. S.; Wade, W. H. Langmuir 1985, 1, 251. (9) Chander, S.;Fuerstenau, D. W.; Stigter, D. In Adsorption from Solution; Ottewill, R. H., Rochester, C. H., Smith, A. L., Eds.; Academic: London, 1983; p 197. (10) Adamson, A. W. Physical Chemistry of Surfaces; Wiley: New
Langmuir, Vol. 3, No. 4, 1987 501
It has long been accepted that the primary forces involved in formation of ionic surfactant surface aggregates are the same as those involved in micelle formation, including the importance of counter ion^.^^^^^ These primary forces are the Coulombic interactions arising from the charges on the surfactant hydrophilic moieties and the mineral surface and the surfactant tail-tail interactions, now explained in terms of the hydrophobic effect.13-19 Various attempts at modeling admicelle formation as a local phase transition have also yielded thermodynamic parameter values approximately the same as analogous quantities found in micelles.l$a These results indicate that counterion interactions with adsorbed aggregates of surfactants should be very important. Systematic studies of the interactions of counterions with surfactant surface aggregates, however, have not been previously reported. In contrast, there are numerous studies of counterion-specific effects in micelle format i ~ n Among . ~ ~ the ~ results reported are that whereas the fraction of "bound" (Le., electrostatically associated) counterions varies little with the type of monovalent counterionZ1in micelle formation, and there is little observed effect of counterion type on the cmc (especially for anionic surfactants),22effects are marked regarding micelle shape and size. Counterion affinity for micelles varies inversely with the radius of anionic counterions or the hydrated radius of cationic counter ion^.^'-^^ Counterion specificity is also more marked for cationic micelles than for anionic micelles. Besides the lack of systematic experimental studies to examine the extent to which these generalizations are applicable to admicelles, there is an absence of theoretical models which include counterion effects. While numerous theoretical treatments at one level or another exist in the l i t e r a t ~ r e , ~ - " none ~ ~ ~ -include ~~ consideration of any counterion/aggregate interaction analogous to counterion binding on micelles. This almost complete neglect of admicelle counterion bindings in theoretical studies of surfactant adsorption is in marked contrast with theoretical studies of counterion (13) Tanford, C. The Hydrophobic Effect: Formation of Micelles and Biological Membranes; Wiley: New York, 1980. (14) Ben Naim, A. Hydrophobic Interactions; Plenum: New York, 1980. (15) Fuerster,au, D. W.; Healy, T. W.; Somasundaran, P. Trans. Am. Inst. Min., Metall. Pet. Eng. 1964, 229, 321. (16) Wakamatau, T.; Fuerstenau, D. W. Adu. Chem. Ser. 1968, 79,161. (17) Dick, S. G.; Fuerstenau, D. W.; Healy, T. W. J. Colloid Interface Sci. 1971, 37, 595. (18) Roy, P.; Fuerstenau, D. W. J. Colloid Interface Sci. 1968,26,102. (19) Bass, B.; Fuerstenau, D. W. Discuss. Faraday SOC.1971,52,371. (20) Lindman, B.; Wennerstrom, H. Top. Curr. Chem. Ser. 1980,87. (21) Kamenka, N.; et al. C.R. Acad. Sci. Paris 1977, 284, 403. (22) Mukerjee, P.; Mysels, K. J.; Kapauan, P. J. Phys. Chem. 1967, 71, 4166. (23) Kale, K. H.; Zana, R. J. Colloid Interface Sci. 1977, 61, 312. (24) Tmamushi, B.; Tamaki, K. Roc. Int. Congr. Surf. Act., 2nd 1957, 3, 449. (25) Trogus, F. J.; Thach, S.; Schechter, R. S.; Wade, W. H. SOC.Pet. Eng. J. 1977, 17, 337. (26) Dobias, B. Colloid Polym. Sci. 1977, 255, 682. (27) Grahame, D. C. Chem. Reu. 1947,41,441. (28) Bisio, P. D.; Cartledge, J. H.; Keesom, W. H.; Radke, C. J. J. Colloid Interface Sci. 1980, 78, 225. (29) Zelenka, R. L.; Radke, C. J. The Role of Surface Charge in Surfactant Adsorption on Oxides; presented at the 1983 Spring National Meeting of the AIChE, Houston, TX, March 27-31, 1983. (30) James, R. 0.;Parks, G. A. In Surface and Colloid Science; Egon Matijevic, E., Ed.; Plenum: New York, 1982; Vol. 12. (31) Anderson, M.; Bauer, C.; Hansmann, D.; Loux, N.; Stanforth, R. In Adsorption of Inorganics at Solid-Liquid Interfaces;Anderson, M., Rubin, A. J., Eds.; Ann Arbor Science: Ann Arbor, MI, 1981. (32) Wilson, D. J. Sep. Sci. 1977, 12, 447. (33) Wilson, D. J.; Kennedy, R. M. Sep. Sci. Technol. 1979,14, 319. (34) Kiefer, J. E.; Wilson, D. J. Sep. Sci. Technol. 1980, 15, 57.
502 Langmuir, Vol. 3, No. 4, 1987
effects and counterion association with micelles, despite strong a priori expectations that these effects should be significant. Numerous theoretical studies of the electrostatics of counterion binding to ionic micelles have ap~eared.~"~~ This general neglect of the role of counterions in admicelle formation, in contrast to the extensive treatments for micelle formation, indicates that the importance of this phenomenon is not yet fully recognized. The studies presented here unambiguously demonstrate the importance of counterion binding to admicelles and provide experimentallydetermined guidelines for future theoretical treatments. These guidelines are discussed at the end of the paper. Experimental Section Lithium dodecyl sulfate (LDS) was purchased from Aldrich Chemical Co. (99.9% pure) and used as received. Sodium dodecyl sulfate (SDS) was recrystallized from ethanol and water before use. Analysis by high-pressure liquid chromatography (HPLC) showed only a single, sharp peak. Potassium dodecyl sulfate (KDS) was prepared by precipitation of SDS with KCl. A saturated KCl solution was combined with a concentrated solution of SDS. The combined solution, now containing a white precipitate, was stirred for 2 days with a magnetic stirrer. The precipitate then was filtered and redissolved in distilled deionized water. In order to dissolve the precipitate, the temperature of the solution was raised above 45 "C. The warm solution containing the precipitate from the previous step was then combined with a fresh solution of potassium chloride. The resulting precipitate was then filtered as before and recrystallized in the same manner as the SDS. The product from recrystallization was estimated to be more than 99% KDS based on the ratio of K+ to Na+ used in this process. Cesium dodecyl sulfate (CDS) was prepared in the same manner as the KDS. LDS was also prepared by this same procedure in order to check unexpected results obtained with the commercial LDS. I t was concluded that the behavior of the commercial LDS was indistinguishable from that prepared as described above. Mukerjee et a1.22previously used CDS and KDS to investigate the role of counterions on the cmc. With the cesium salt, he noted that the solubility was at or near the critical micelle concentration at 25 "C with no added salt. Mukerjee et al. also noted, however, that CDS formed a "quite stable supersaturated solution" which permitted measurement of the cmc. For KDS, the stability of a supersaturated solution was not suitable for cmc measurements a t 25 "C, and although cmc data were reported for KDS a t 32 "C, information that was reported on other systems was not reported for KDS because of "internal inconsistencies". Likewise, in this study, the authors were unable to prepare a KDS solution above the cmc at 30 "C. The Krafft point of KDS in 0.15 M KC1 was determined to be approximately 42 "C. For CDS, the Krafft temperature in 0.15 M CsC1 was approximately 32 "C. All studies of CDS in this work were made at 35 "C. Static adsorption isotherms were prepared by placing 0.5 g of A1203in a capped test tube with 15 mL of solution. The filled test tubes were allowed to equilibrate for a minimum of 4 days in a thermostated water bath and then centrifuged at ambient temperature for about 20 min. The clear supernatant solution above the alumina was then decanted for analysis by HPLC. The adsorption was calculated by taking the difference between the (35)Stigter, D. J. Phys. Chern. 1964,68, 12. (36)Stigter, D. J. Colloid Interface Sci. 1974,47, 2. (37)Stigter, D.J.Phys. Chern. 1974,78, 2480. (38)Stigter, D.J. Phys. Chern. 1975,79, 1008. (39)Stigter, D.J. Phys. Chern. 1975,79, 1015. (40)Wennenstrom, H.K.;Jonsson, B.; Linse, P. J.Phys. Chern. 1982, 76, 4665. (41)Hirasake, G. J.; Lawson, J. B. SPE 10921,presented at 57th Annual Fall Technical Conference and Exhibition of the Society of Petroleum Engineers of AIME, New Orleans, 1982. (42)Linse, D.; Gunnarsson, G.; Jonsson, B. J.Phys. Chern. 1982,86, 413. (43)Rathman, J. F.;Scamehorn, J. F. J.Phys. Chern. 1984,88,5807. (44)Mukerjee, P.J. Phys. Chem. 1969,73, 2054.
Bitting and Harwell initial and final concentrations of the surfactant. I t should be noted that this procedure does not eliminate the possibility that some of the surfactant attributed to the adsorbed phase had actually been hydrolyzed to dodecanol. The alumina used in these experiments was primarily a y structure and was purchased from Degussa Corporation (trade name Aluminum Oxide C). The surface area of the alumina is reported by the manufacturer to be 100 (h15) m2/g. It was used as received. Since hydrogen ions are potential-determining ions for alumina in aqueous solutions,45 control of the p H of the adsorption isotherms was essential. In all solutions the final pH was obtained by addition of commercially available standard normality solutions of HC1, NaOH, or LiOH, purchased from Fisher Scientific. A known quantity of acid and base was added to the surfactant feed, and the final p H was measured after equilibration. The initial quantity of added acid or base was determined from a calibration curve relating the quantity of added acid or base and the final pH of the solution in contact with the alumina in the absence of added surfactant, but a t the same solution to solid ratio as used in the surfactant adsorption isotherms. Since the adsorption of surfactant also affected the final pH, this procedure resulted in slight systematic deviations in the fiial pH values of the isotherm points. These deviations were recorded and are taken into account in the interpretation of the adsorption isotherms. For the plateau adsorption isotherms (Figures 2 and 4-6) it was frequently necessary to add additional acid or base, reequilibrate the solution, and remeasure the pH for several repetitions before the desired equilibrium value was obtained. In such cases the adsorption was corrected for the added volume of solution. Due to the previously mentioned high Krafft point of KDS, experiments involving KDS were performed in a circulated water bath a t 50 "C. One problem associated with this system is the handling of the concentrated surfactant to avoid precipitation. However, by using an autopipetter, concentrated solutions of KDS were pipetted and dispensed without visible evidence of precipitation in either the test tubes or the autopipetter. The final concentration of the static adsorption isotherms was determined by HPLC. Details of the analytical procedure are available elsewhere.46 The electrophoretic mobility of charged alumina particles in aqueous solutions was determined with a Zeta Meter System 3.0. For solutions with a high specific conductance, as was the case in most of these experiments, a low voltage was required. Consequently, the velocity of the particles was reduced. Samples for the electrophoresis cell were prepared by placing 1.5 mg of alumina in 30 mL of solution. After 24 h, the p H of the samples was measured and the samples were then loaded individually into the electrophoresis cell. Before loading a sample, the cell was cleaned with distilled deionized water and dried thoroughly with Kimwipes. Approximately 5 mL of the sample was then rinsed through the cell and shaken out before pouring the remainder of the sample into the cell. Final surfactant concentrations were calculated by assuming an adsorption of the surfactant on the basis of the separately determined adsorption data. This correction factor was generally negligible, however, because of the very high solution/solid ratio used in the electrophoretic mobility measurements. Additional details for these measurements are available elsewhere.46 LiC1, NaCl, KC1, CsC1,l.O N HC1, LiOH.H20, 1.0 N NaOH, KOH, and SDS were all Fisher Certified products. LDS was purchased from Aldrich Chemical. All were used as received, except for the recrystallizations described above.
Results Measured plateau adsorption values for SDS in 0.15 M background NaCl as a function of pH at 30 "C are shown in Figure 2. These plateau values were determined by measuring the adsorption from solution in systems where the equilibrium SDS concentration was at least 4 times the (45)Yopps, J. A.;Fuerstenau, D. W. J. Colloid Sci. 1964,19,61. (46)Bitting, D. M.S. Thesis, University of Oklahoma, 1985. (47)Davis, J. A.;James, R. 0.;Leckie, J. 0. J.Colloid Interface Sci. 1978,63,480.
Langmuir, Vol. 3, No. 4, 1987 503
Counterion Effect on Surfactant Surface Aggregates
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approximate value of the cmc of SDS at the temperature and NaCl concentration of the system. As has been previously reported'+ there is a dramatic effect of pH on the maximum obtainable surfactant adsorption. The adsorption decreases steeply from a roughly constant value of 925-1000 pmol/g at pH 2-4 down to approximately zero adsorption at pH 11.3. At 30 "C the point of zero charge (pzc) of the alumina used in this study was found to occur at pH 9.5, as shown in Figure 3. The plateau adsorption of SDS at this pH is approximately 340 pmol/g, over one-third of the maximum value of the adsorption which could be obtained. The plateau adsorption isotherm in Figure 2 is plotted again in Figure 4 along with plateau adsorptions from systems with decreasing concentrations of added NaC1. While the isotherms at lower added electrolyte concen-
trations retain the same basic shape as the isotherm from 0.15 M NaC1, the maximum amount of adsorption decreases significantly with decreasing concentration of NaCl from -960 pmollg from the 0.15 M solution, to =760 pmol/g from the 0.05 M solution, and finally to ~ 6 4 0 pmollg from the solutions with 0,Ol M or no added NaC1, a decrease of over 30%. The behavior of the plateau isotherms with lesser amounts of added NaCl in the vicinity of the pzc is obscured by the need to use excessive amounts of 1N NaOH to raise the pH of the equilibrium solutions above pH ~ 8 . 5 This . problem arose because of the relatively high solid/liquid ratio in the test tubes and could be avoided by using a lower solid/solution ratio. The alumina adsorbs significant amounts of H+ions. By the time the pH had been adjusted to above -8.5, the amount of Na+ added as base was sufficient to significantly perturb the isotherms. For lesser amounts of A1,0,, less NaOH would need to be added to raise the equilibrium pH. Such a lower solid/solution ratio would decrease the accuracy with which the surfactant adsorptions could be determined, however. The effect on the plateau adsorption isotherm of changing the so-called "indifferent electrolytenmfrom NaCl to LiCl at a constant added electrolyte concentration of 0.15 M and at 30 "C as shown in Figure 5. The effect of pH on the adsorption from the system with added LiCl is dramatically less than the effect on the system with added NaC1. While reaching a maximum of roughly 700 wmol/g at low pH, the isotherm goes through what appears to be a global minimum near pH 8, then actually shows a slight increase with increasing pH. It should be noted that LiOH was used to adjust the pH in the solutions with added LiC1. The LiCl and NaCl surfactant isotherms cross at an adsorption of about 500 pmollg between pH 8 and 9. Plateau adsorption isotherms with KC1 and CsCl as the added electrolyte were not obtained. As already noted, the Krafft point of the KDS was found to be near 45 "C. Clear solutions of CDS (Krafft point 32 "C) at concen-
504 Langmuir, Vol. 3, No. 4, 1987
Bitting and Harwell
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Figure 5. Effect of varying the counterion on the plateau adsorption of surfactant. trations above the cmc were obtained at a temperature of 35 "C. Adsorption data below the cmc were obtained for such CDS solutions, but attempts to measure adsorptions with equilibrium CDS concentrations above the cmc led to anomalously high values, suggesting either that CDS precipitation was occurring within the sample tubes or that CDS was somehow being lost due to precipitation during handling after removal from the temperature bath. It is also possible that 35 "C is not sufficiently above the Krafft point to prevent precipitation of the CDS onto the alumina surface. The effect on LDS plateau adsorption of decreasing the amount of added LiCl is shown in Figure 6. Plateau adsorption isotherms with 0.15 M added LiCl and 0.01 M added LiCl are compared. As with the isotherms with added NaC1, there is a significant decrease in the surfactant adsorption all along the isotherm, although the general shape of the isotherm is unchanged. The value of the minimum in the isotherm decreases from 4 0 0 to =300 pmol/g with decreasing LiC1. It should be noted that NaOH was used to adjust the pH above 8 in the 0.01 M solutions. Comparing the isotherms from the added NaCl solutions in Figure 4, however, indicates that the Na+ has considerably less effect on the adsorption at higher values of the pH than does Li+, so that the effect of increasing pH is actually more clearly seen in the added NaOH isotherm than in the added LiOH isotherms. The almost vertical rise in the low added LiCl isotherm below pH 5 may indicate that LDS precipitate was forming in this region. Apparent LDS adsorptions below pH 4 were anomalously high at 0.01 M added LiCl and are not reported, also indicating the possibility of precipitation. The effect on the adsorption isotherm of changing the counterion at constant pH, temperature, and added electrolyte is shown in Figure 7. Comparing the data below 70 p M , it is seen that the adsorption increases at constant surfactant concentration in the order CDS < SDS < LDS. The isotherms cross near a surfactant concentration of 70 pM, the adsorption increasing above that concentration in the order CDS > SDS > LDS. A comparable adsorption
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EFFECT OF COUNTERION TYPE ON DODECYLSULFATE ADSORPTION
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Figure 7. Effect of counterion type on surfactant adsorption for 0.15 M added chloride salt at pH 3.9 and 30 "C. isotherm for KDS is shown in Figure 8. Though the temperature difference makes a direct comparison difficult, the same trend appears to be followed. One data point for SDS at 50 "C is plotted to confiim that the SDS adsorption does decrease with increasing temperature but that this decrease is not dramatic; hence, it is not without value to compare the KDS at 50 "Cto the SDS at 30 "C. Projection to low-pH surfactant adsorption suggests the order CDS < KDS < SDS < LDS at low surface coverages switching to CDS > KDS > SDS > LDS nearer the plateau adsorption region. Adsorption isotherms at constant pH, added NaC1, and temperature are shown as functions of increasing SDS concentration at two different values of the pH in Figures 9 and 10. Both isotherms exhibit the typical shape of a surfactant isotherm for adsorption on alumina at constant added electrolyte. The region I/region I1 transition was not obtained for the lower pH isotherm because of difficulties in determining small surfactant concentration changes for surfactant equilibrium concentrations below
Langmuir, Vol. 3, No. 4, 1987 505
Counterion Effect on Surfactant Surface Aggregates
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