Thermodynamics of Micellization and Adsorption of Three

Langmuir 2004, 20, 11387-11392

11387

Thermodynamics of Micellization and Adsorption of Three Alkyltrimethylammonium Bromides Using Isothermal Titration Calorimetry Steven P. Stodghill,*,†,‡ Adam E. Smith,§ and John H. O’Haver§ Department of Pharmaceutics, School of Pharmacy, The University of Mississippi, University, Mississippi 38677, National Center for Natural Products Research, Thad Cochran Research Center, University, Mississippi 38677, and Department of Chemical Engineering, School of Engineering, The University of Mississippi, University, Mississippi 38677 Received August 16, 2004. In Final Form: September 30, 2004 Studies of the thermodynamic properties of micellization, as well as the enthalpy change of adsorption (displacement), were conducted using isothermal titration calorimetry (ITC). The cationic surfactants, dodecyltrimethylammonium bromide, tetradecyltrimethylammonium bromide, and hexadecyltrimethylammonium bromide or cetyltrimethylammonium bromide were used. Adsorption studies were performed utilizing HiSil 233 precipitated silica as the substrate. The thermodynamics of micellization were studied at 28, 30, and 35 °C using ITC. ∆Hdil was calculated and graphed versus concentration in order to determine ∆Hmic and critical micelle concentration. From these data, ∆Gmic and ∆Smic were also determined and found to be in agreement with values previously determined using traditional temperature-variation methods. The thermodynamics of adsorption were also determined at the above temperatures using ITC. Using the results, the heat of displacement, Qdisp, and enthalpy of displacement, ∆Hdisp, were calculated. A plot of ∆Hdisp as a function of concentration was generated and used to interpret the mechanism of adsorption in the four regions of a typical cationic surfactant adsorption isotherm.

Introduction The micellization process of surfactants has been extensively studied since the time McBain first described surfactant aggregates in solution.1 Additionally, the thermodynamics of adsorption at solid/liquid interfaces has been studied for many years. Investigating the enthalpy, entropy, and Gibbs free energy of adsorption, researchers can gain insight into the mechanisms by which surfactants adsorb at an interface.2 The utilization of calorimetry for examining the process of micellization and adsorption is not a novel concept, but the application of calorimetry to surfactant systems has encountered difficulties. Until recently, calorimeters did not possess the sensitivity required to study long-chain surfactant systems which have low critical micelle concentrations (CMCs). However, the advent of highly sensitive, highly accurate isothermal titration calorimetry (ITC) has greatly benefited the study of surfactant systems. Before the use of high-sensitivity calorimetry, thermodynamic values were typically obtained by determining the effect of temperature on the CMC. The change in Gibbs free energy, ∆Gmic, can be evaluated by the CMC using one of several models relating the two quantities. The first model, the phase separation (PS) model, assumes 100% binding of counterions. The PS model does not, however, accurately model the ∆Hmic for ionic surfactants and is not a suitable model for this study.3 Another model, * Author to whom correspondence should be addressed. Address: 104 Faser Hall, Department of Pharmaceutics, University, MS 38677. Phone: (662) 915-5164. Fax: (662) 915-1177. E-mail: [email protected]. † School of Pharmacy, The University of Mississippi. ‡ Thad Cochran Research Center. § School of Engineering, The University of Mississippi. (1) McBain, J. W. Trans. Faraday Soc. 1913, 9, 99-101. (2) Rosen, M. J. Surfactants and Interfacial Phenomena, 2nd ed.; Wiley: New York, 1989. (3) Bijma, K.; Engberts, J. Langmuir 1994, 10, 2578-2582.

the mass-action (MA) model, takes into account the fraction of counterions bound to the micelle and is much better suited for modeling ionic surfactants. Equations 1 and 2 can be used to determine ∆Gmic and ∆Hmic from the temperature dependence of the CMC.

∆GMA mic ) (1 + β)RT ln(Xcmc)

(

2 ∆HMA mic ) -(1 + β)RT

(1)

)

∂(ln(Xcmc)) ∂T

P

(2)

In the above equations, R is the ideal gas constant, T is the absolute temperature, and Xcmc is the mole fraction of surfactant at the CMC. The parameter β represents the fraction of counterions that bind to the micelle. Once these values are obtained, ∆Smic can be evaluated by the Gibbs equation (eq 3).

∆Smic )

∆Hmic - ∆Gmic T

(3)

As previously stated, this method’s accuracy has been questioned because it relies on the temperature dependence of the CMC to evaluate ∆Hmic. The concern is that the large temperature variation that is required for this method can affect the micellar structure.4 Researchers have realized that the previously discussed method needs to be verified by calorimetry.5,6 Another method used to investigate the thermodynamics of micellization is the use of calorimetric experiments to find ∆Hmic. Using eq 1 to evaluate ∆Gmic with the ∆Hmic found using calorimetry provides a more-accurate depiction of the thermodynamics of micellization. Table 1 shows (4) Paredes, S.; Tribout, M.; Sepulveda, L. J. Phys. Chem. 1984, 88, 1871-1875. (5) Holtzer, A.; Holtzer, M. F. J. Phys. Chem. 1974, 78, 1442-1443. (6) Kresheck, G. C. J. Phys. Chem. B 1998, 102, 6596-6600.

10.1021/la047954d CCC: $27.50 © 2004 American Chemical Society Published on Web 11/19/2004

11388

Langmuir, Vol. 20, No. 26, 2004

Table 1. Previously Published Thermodynamic Values of Micellization Using Calorimetry DTAB ref

temp (°C)

∆Hmic (kcal mol-1)

∆Gmic (kcal mol-1)

∆Smic (cal mol-1 K-1)

9 8 11 10 10 10

25 25 30 25 30 35

-0.525 -0.549 -1.218 -0.389 -0.846 -1.216

-8.885 -8.598 -4.395

28.017 26.918 10.485

ref

temp (°C)

∆Hmic (kcal mol-1)

∆Gmic (kcal mol-1)

∆Smic (cal mol-1 K-1)

8 11 13 12 10 10 10

25 30 25 25 25 30 35

-1.17 -2.033 -1.072 -1.027 -1.087 -1.622 -2.207

-10.509 -5.756

31.408 12.301

ref

temp (°C)

∆Hmic (kcal mol-1)

∆Gmic (kcal mol-1)

∆Smic (cal mol-1 K-1)

8 11 13 7 7 7 7 12 4

25 30 25 25 29.9 40 50.1 25 25

-2.603 -3.32 -2.03 -2.866 -3.248 -4.419 -5.685 -1.935 -2.19

-12.348 -8.121

32.674 15.835

CTAB

14.8

some thermodynamic values of micellization that have been found via calorimetry. As can be seen, the literature values for these fundamental thermodynamic properties vary significantly and need to be researched further to determine the true values. Miller has performed research into the thermodynamic properties of adsorption for homologous series of cationic and nonionic surfactants using two theoretical models: the Frumkin and the reorientation model.14 Other methods to determine the thermodynamic properties experimentally involve using eq 4 with an adsorption isotherm from which the Gibbs free energy of adsorption, ∆Gads, can be found.

∆Gads ) -RgasT ln(K)

T2 ∆Sads )

TTAB

-6.6

∆Hads

(4)

In eq 4, Rgas is the ideal gas constant, T is the absolute temperature, and K is the distribution coefficient of surfactant between the surface and the bulk solution. By using the previous equation with adsorption isotherms at varying temperatures, eqs 5 and 6 can be used to evaluate the enthalpy, ∆Hads, and entropy of adsorption, ∆Sads.15 (7) Bergstrom, S.; Olofsson, G. Thermochim. Acta 1986, 109, 155164. (8) Mosquero, V.; Manuel del Rio, J.; Attwood, D.; Garcia, M.; Jones, M. N.; Prieto, G.; Suarez, M. J.; Sarmiento, F. J. Colloid Interface Sci. 1998, 206, 66-76. (9) Muller, N. Langmuir 1993, 9, 96-100. (10) Bai, G.; Wang, J.; Yan, H.; Li, Z.; Thomas, R. K. J. Phys. Chem. B 2001, 105, 9576-9580. (11) Majhi, P. R.; Moulik, S. P. Langmuir 1998, 14, 3986-3990. (12) Lah, J.; Pohar, C.; Vesnaver, G. J. Phys. Chem. B 2000, 104, 2522-2526. (13) Blandamer, M. J.; Cullis, P. M.; Engberts, J. J. Chem. Soc., Faraday Trans. 1 1998, 94, 2261-2267. (14) Miller, R. J. Surfactants Deterg. 2002, 5, 281-286. (15) Clint, J. H. Surfactant Aggregation; Blackie: Glasgow, 1992.

[ ]

∆Gads T )∂T

Stodghill et al.

∂

θ

∆Hads - ∆Gads T

(5)

(6)

A problem with this method is that it is often difficult to choose points on adsorption isotherms at different temperatures that correspond to the same surface coverage. This problem can be overcome by employing calorimetry to directly measure the heat of adsorption.15 There has been a recent trend to use microcalorimetry to study the thermodynamics of surfactant adsorption.16-30 There are two basic methods used to study the enthalpy of surfactant adsorption, substrate immersion, and titration of surfactant into an aqueous solution containing the substrate.31 The preferred method is the titration of surfactant because this method provides data points along the entire length of the adsorption isotherm in one experiment, whereas the immersion of the substrate provides data at only one point. The calorimetry method is not, however, without problems. The foremost difficulty in studying the adsorption of surfactant onto a substrate by titration calorimetry is that the heat effect found is actually the sum of two separate processes: the simultaneous desorption of water and the adsorption of surfactant. Denoyel states that it is better to reserve the term “enthalpy of adsorption” for processes involving only the adsorption of surfactant and that the term “enthalpy of displacement” better describes the heat effect of the combined processes. In this paper, studies of the thermodynamic properties of micellization, as well as the enthalpy change of adsorption (displacement), are conducted using ITC. The cationic surfactants, dodecyltrimethylammonium bromide (DTAB), tetradecyltrimethylammonium bromide (TTAB), and hexadecyltrimethylammonium bromide or cetyltrimethylammonium bromide (CTAB), are used, and comparisons between data obtained via the temperature variation method and the ITC method will be made. Experimental Methods Materials. The surfactants used in this study, DTAB, TTAB, and CTAB, were obtained from Sigma (St. Louis, MO) and were used without further purification. Sodium hydroxide was also (16) Grosmaire, L.; Chorro, M.; Chorro, C.; Partyka, S.; Boyer, B. Thermochim. Acta 2001, 379, 261-268. (17) Zajac, J.; Trompette, J. L.; Partyka, S. Langmuir 1996, 12, 13571367. (18) Zajac, J. Colloids Surf., A 2000, 167, 3-19. (19) Seidel, J.; Wittrock, C.; Kohler, H. H. Langmuir 1996, 12, 55575562. (20) Benalla, H.; Zajac, J.; Partyka, S.; Roziere, J. Colloids Surf., A 2002, 203, 259-271. (21) Denoyel, R.; Rouquerol, F.; Rouquerol, J. Colloids Surf. 1989, 37, 295-307. (22) Tan, X.; Yan, H.; Hu, R.; Hepler, L. Chin. J. Chem. 1992, 10, 200-205. (23) Van Os, N. M.; Haandrikman, G. Langmuir 1987, 3, 10511056. (24) Seidel, J. Thermochim. Acta 1992, 229, 257-270. (25) Kiraly, Z.; Findenegg, G. H.; Klumpp, E.; Schlimper, H.; Dekany, I. Langmuir 2001, 17, 2420-2425. (26) Seidel, J. Prog. Colloid Polym. Sci. 1992, 89, 176-180. (27) Partyka, S.; Lindheimer, M.; Zaini, S.; Keh, E.; Brun, B. Langmuir 1986, 2, 101-105. (28) Kiraly, Z.; Findenegg, G. H. Langmuir 2000, 16, 8842-8849. (29) Findenegg, G. H.; Pasucha, B.; Strunk, H. Colloids Surf. 1989, 37, 223-233. (30) Zajac, J.; Chorro, M.; Chorro, C.; Partyka, S. J. Therm. Anal. 1995, 45, 781-789. (31) Denoyel, R. Colloids Surf., A 2002, 205, 61-71.

Thermodynamics of Micellization and Adsorption

Langmuir, Vol. 20, No. 26, 2004 11389

Figure 1. Plot of ∆Hdil versus concentration for DTAB at 28, 30, and 35 °C.

Figure 3. Plot of ∆Hdil versus concentration for CTAB at 28, 30, and 35 °C.

Figure 2. Plot of ∆Hdil versus concentration for TTAB at 28, 30, and 35 °C.

Figure 4. Plot of ∆Hdil for DTAB at 28 °C showing the method of calculating ∆Hmic.

obtained from Sigma (St. Louis, MO). The precipitated silica used in this study, HiSil 233 (specific surface area of 150 m2 gm-1) was supplied by PPG (Pittsburgh, PA). Deionized water was obtained from a Barnstead E-pure Water System with the resistance kept constant at 18.3 MΩ cm. Thermodynamics of Micellization. The calorimetry experiments were performed using a titration microcalorimeter (Microcal VP-ITC, North Hampton, MA) at 28, 30, and 35 °C. A computer-controlled syringe is used to inject 5 µL aliquots of surfactant solutions at concentrations ranging from 10 to 15 times the CMC. The surfactant solution is injected into a 1.4 mL sample cell containing purified water that has been adjusted to a pH of 8 by the addition of small volumes of sodium hydroxide. After each injection, the computer records the heating rate required to maintain a constant temperature difference between the sample cell and a reference cell filled with purified water. The integration of the rate of heating yields a plot of heat required to maintain the temperature difference after each injection, qinj, versus the increasing concentration of surfactant in the sample cell, Ccell.32 The enthalpy change of dilution, ∆Hdil, is found by dividing qinj by the number of moles of surfactant injected in each 5 µL aliquot. A graph of ∆Hdil versus concentration exhibits two plateaus. The first plateau, found at low concentrations, represents the dilution of a micellar solution into a monomeric solution. The second plateau is representative of a micellar solution being injected into another micellar solution and is far less endothermic than the first plateau. The enthalpy of micellization, ∆Hmic, is found by subtracting the plateau value at low surfactant concentrations from the plateau at high surfactant concentrations. This plot of ∆Hdil versus Ccell not only provides a value for

∆Hmic, but the CMC of the surfactant can also be evaluated by finding the inflection point between the high-concentration plateau and the low-concentration plateau.11 Once ∆Hmic and the CMC at the three temperatures have been determined, the MA model is used to find ∆Gmic and ∆Smic.33 Thermodynamics of Adsorption. The data for the adsorption isotherms for DTAB, TTAB, and CTAB were taken from Dickson.34 Although the adsorption of surfactants decreases with increasing temperature, this variation for ionic surfactants is relatively small and has been considered to be insignificant for the range studied.2 The calorimetry experiments were performed using a titration microcalorimeter (Microcal VP-ITC, North Hampton, MA) at 28, 30, and 35 °C. Surfactant solutions ranging in concentration from 10 to 15 times the CMC were again used for this study. The surfactant solution is injected in 5 µL aliquots into purified water and a silica solution prepared by placing 0.1 g of silica in 20 mL of purified water. Both the purified water and silica solution have been adjusted to a pH of 8 by adding small volumes of a concentrated sodium hydroxide solution. When the silica solution is titrated with surfactant, both dilution and adsorption effects are detected by the microcalorimeter. Equation 7 can be used to determine the heat effect due to the displacement of water by surfactant (Qdisp) at the solid/liquid interface.

(32) Bijma, K.; Engberts, J.; Blandamer, M. J.; Cullis, P. M.; Last, P. M.; Irlam, K. D.; Soldi, L. G. J. Chem. Soc., Faraday Trans. 1 1997, 93, 1579-1584.

Qdisp ) Qtotal - Qdil

(7)

In eq 7, Qtotal and Qdil are the experimental heat determined by titrating the silica solution with surfactant and the heat effect of dilution of the surfactant solution into purified water, (33) Stenius, P.; Backlund, S.; Ekwall, P. IUPAC Chem. Data Ser. 1980, 28, 295. (34) Dickson, J. Master’s Thesis, The University of Mississippi, Oxford, Mississippi, 2001.

11390

Langmuir, Vol. 20, No. 26, 2004

Stodghill et al.

Figure 5. Adsorption isotherms for DTAB, TTAB, and CTAB.34

Figure 6. Typical adsorption isotherm of an ionic surfactant adsorbing onto an oppositely charged hydrophilic surface.

Figure 7. Adsorption and the enthalpy of displacement of DTAB plotted against equilibrium concentration.

Figure 8. Adsorption and the enthalpy of displacement of TTAB plotted against equilibrium concentration.

Table 2. Thermodynamic Properties of Micellization for DTAB, TTAB, and CTAB at 28, 30, and 35 °C as Determined by the ITC Method temp (K)

CMC (µM)

301.15 303.15 308.15

15 175 15 510 15 830

301.15 303.15 308.15 301.15 303.15 308.15

∆Hmic (kcal mol-1)

∆Gmic (kcal mol-1)

∆Smic (cal mol-1)

DTAB -0.706 -0.828 -1.273

-8.690 -8.724 -8.846

26.511 26.048 24.578

3825 3925 3970

TTAB -1.413 -1.669 -2.208

-10.150 -10.190 -10.345

29.012 28.105 26.408

925 950 975

CTAB -2.246 -2.656 -3.443

-11.653 -11.702 -11.867

31.237 29.841 27.336

respectively. Once Qdisp is determined, the enthalpy of displacement, ∆Hdisp, is found by dividing Qdisp by the number of moles adsorbed onto the surface in each step as determined from the adsorption isotherm. The enthalpy of displacement is then graphed versus equilibrium cell concentration.

Results and Discussions Micellization Thermodynamics by ITC. As can be seen from Figures 1-3, all ∆Hdil were endothermic. Figure 4 demonstrates how CMC and ∆Hmic are determined from the plot of ∆Hdil versus concentration. All ∆Hmic were exothermic for the studied cationic surfactants. It is interesting to note that the plateau in these enthalpograms does not stay flat. The enthalpograms for DTAB all exhibit

Figure 9. Adsorption and the enthalpy of displacement of CTAB plotted against equilibrium concentration.

a positive slope, whereas the enthalpograms for TTAB and CTAB are all relatively flat. The enthalpogram for CTAB at 28 °C, in fact, has a slightly negative slope. This can be explained by the tendency of the CMC to increase with decreasing tail length. This means that for DTAB, with a CMC of 16 000 µM, there will be a higher concentration of surfactant in the injected solution than for CTAB, which has a CMC of only 920 µM.2 This increase in concentration means that the behavior of the solution being injected and the solution in the sample cell cannot be assumed to be ideal.32 It can also been seen from Figures 1-3 that another consequence of decreasing tail length

Thermodynamics of Micellization and Adsorption

Langmuir, Vol. 20, No. 26, 2004 11391

Figure 10. Comparison of two surface-aggregate theories presented by Somasundaran et al. and Harwell et al.36,40

is that the endotherms for the dilution process are not as large and ∆Hmic tends to decrease. The enthalpograms for TTAB at all three temperatures and CTAB at 30 and 35 °C would be classified as type A and the enthalpograms of DTAB and CTAB at 28 °C would be classified as type B.32 Type A enthalpograms are considered to be “textbook” examples in which the surfactants behave ideally and do not demonstrate solutesolute interactions.32 In the second type of enthalpogram, type B, the solute-solute interactions begin to impact the enthalpy change, causing the enthalpy of dilution to be more endothermic for each injection until the CMC is reached.32 As noted before, the CMC increases within a series of surfactants as the tail length decreases. From Figures 1-3, it can be noticed that the CMC for a surfactant also tends to increase with increasing temperature. Table 2 shows the CMCs determined from the ITC experiments, which correspond fairly well with values previously published and summarized in Table 1. Using the ITCdetermined CMC values, ∆Gmic and ∆Smic were found using eqs 1 and 3 with β values taken from the literature.11 Table 2 shows the values of ∆Hmic, ∆Gmic, and ∆Smic found in this study. Comparing these values to those in Table 1, it can be seen that the thermodynamic quantities found in this study are very similar to previously published data. The increase in temperature causes a decrease in both the enthalpy and entropy of micellization. This is due to the decrease in the amount of water that is ordered due to its nearness to the hydrophobic alkane chains and the amount of water bound to the trimethylammonium headgroups.2 While the decrease in ∆Hmic and ∆Smic have opposite effects on the Gibbs free energy, it is seen that ∆Gmic also decreases, meaning that the change in ∆Hmic has more of an effect on ∆Gmic than the decrease in ∆Smic. From the data presented in Table 1, it is also noted that, as the chain length increases, ∆Gmic becomes more negative. Previous studies have shown that ∆Gmic becomes more negative by about 3 kJ (0.7 cal) for each additional -CH2- group in the hydrophobic tail.2 This is true of the values represented in Table 2. Another trend evident from the data presented in Table 2 is that at any given temperature the change in each of the thermodynamic parameters varies in a consistent manner with the addition of -CH2- groups. However, due to the presence of an entropy/enthalpy compensation effect, the contribution of each of these terms to the free energy changes. From Table 2 it can be seen that at a given temperature the ∆Gmic decreases by approximately 1.5 kcal mol-1 for each addition of two -CH2- groups. Likewise, the ∆Hmic decreases by ∼0.92 kcal mol-1 and

the ∆Smic increases by ∼1.9 cal mol-1 with the addition of each pair of -CH2- groups. The data suggest that the large negative values of the free energy are primarily due to entropic contributions; however, the percentage of the free energy change that is due to the enthalpic term increases with increasing tail length. These findings appear to support interpretations published previously regarding the entropy/enthalpy compensation effect in micelle formation.35 Thermodynamics of Adsorption by ITC. The values for the adsorption isotherm for the three surfactants are presented in Figure 5. It can be seen that the data for all three adsorption isotherms demonstrate the characteristic shape of an ionic surfactant adsorbing onto an oppositely charged surface, as presented in Figure 6. These data were used to find the amount of the surfactant adsorbed onto the silica surface after each injection. Using eq 7, the heat of displacement is found for each injection by subtracting the heat of dilution from the heat observed in the adsorption experiments. This subtraction of the heat of dilution is necessary because the surfactant injections must be diluted from a micellar syringe concentration to the premicellar cell concentration before any surfactant monomer can adsorb. The enthalpy of displacement, ∆Hdisp, is found by dividing the heat of displacement by the number of moles that adsorb in each step. Figures 7-9 show plots of the adsorption isotherm and ∆Hdisp for each of the three surfactants studied. To understand the complex shape of these plots, it is necessary to discuss the mechanism of adsorption in the three regions of the adsorption isotherm. From inspecting Figures 7-9, it is noticed that two sharp peaks occur at the transition for region I to region II and at the transition between regions II and III. This can be explained by how the surfactant adsorbs onto the surface in each of these regions. Figure 10 illustrates two of the proposed surface aggregate theories for each region of the isotherm. Region I is characterized by the adsorption of surfactant monomers and has little if any interaction with one another. This region is referred to as the “Henry’s law region”.2 In this region, ∆Hdisp is due only to the interactions of monomer with the surface, and there is little if any heat effect due to the interaction of adjacent surfactant monomers. The sharp increase in the slope of region II has been hypothesized to be due to the beginning of surface aggregation. The peak at this transition can be explained by change in enthalpy due to the initial interactions of surfactants with one another on the surface. The other (35) Fisicaro, E.; Compari, C.; Braibanti, A. Phys. Chem. Chem. Phys. 2004, 6, 4156-4166.

11392

Langmuir, Vol. 20, No. 26, 2004

peak occurs at the transition from region II to region III. In region III, the mechanism of adsorption is not well agreed upon. Somasundaran et al. hypothesize that adsorption in region II occurs as monolayer aggregates with the headgroup oriented toward the surface and the change in slope coinciding with the cancellation of surface charge by the charged headgroups.36 Other researchers believe that the second layer of adsorbed aggregates is not present until region III, which accounts for the change in slope.37-38 Harwell et al. have attributed the difference between regions II and III to the presence of adsorption sites with different energies.39-40 In region II, the highenergy sites will be taken, and the slope in region III is due to the beginning of adsorption to the low-energy sites. Whatever the mechanism, the peak can again be attributed to the changing of how the surfactant molecules interact with one another on the surface. From the figures, it can be seen that, as the length of the hydrophobic tail increases, the variation of ∆Hdisp decreases. For DTAB, there is a difference from peak to peak of about 9.1 kcal mol-1. For TTAB and CTAB, this difference is 3.9 and 2.3 kcal mol-1, respectively. It can be noticed from the graphs that there is no clear dependence of the variation from the exothermic to the endothermic peak on the temperature. The three surfactants exhibited the largest variation from exothermic peak to endothermic peak at different temperatures, so no correlation between temperature and ∆Hdisp can be drawn at this time. Conclusions The purpose of this study was to investigate the thermodynamics of the micellization process, as well as (36) Somasundaran, P.; Healy, T. W.; Fuerstenau, D. W. J Phys. Chem. 1964, 68, 3562-3566. (37) Chandar, P.; Somasundaran, P.; Waterman, K. C.; Turro, N. J. J. Phys. Chem. 1987, 91, 148-150. (38) Chandar, P.; Somasundaran, P.; Turro, N. J. J. Colloid Interface Sci. 1987, 117, 31-46. (39) Scamehorn, J. F.; Schechter, R. S.; Wade, W. H. J. Colloid Interface Sci. 1982, 85, 463-477. (40) Harwell, J. H.; Hoskins, J.; Schechter, R. S.; Wade, W. H. Langmuir 1985, 1, 251-262.

Stodghill et al.

the enthalpy change due to surfactant adsorption onto a hydrophilic silica surface, and how changes in temperature affect these thermodynamic quantities. This goal was achieved by using ITC to determine the enthalpy of micellization, CMC, and the change in enthalpy of displacement at three temperatures, 28, 30, and 35 °C, for three cationic surfactants, DTAB, TTAB, and CTAB. The change in enthalpy due to micellization was determined from the plots of the heat of dilution versus concentration. As has been widely reported in the literature, ∆Hmic was found to increase with increasing temperature and with increasing length of the hydrophobic alkyl tail group for the three surfactants studied. The change in Gibbs free energy was found using eq 1 with the CMC determined experimentally from ITC. The Gibbs free energy change was found to be negative, implying, as expected, that micellization occurs spontaneously once the CMC has been reached. The values of ∆Gmic were found to become more negative with increasing temperature due to the increased hydrophobicity of the methyl groups of the surfactant tail. The change in entropy was found to decrease with increasing temperature and with decreasing length of the tail groups. The values of the enthalpy, Gibbs free energy, and entropy of micellization matched literature values previously published. The change in enthalpy due to surfactant adsorption onto a hydrophilic silica surface was determined from the plots of the heat of displacement versus concentration. In the graphs of ∆Hdisp, the transitions between regions I and II and regions II and III were marked by noticeable exothermic and endothermic peaks, respectively. These peaks have been attributed to a change in how the surfactant adsorbs in each of the three regions. No discernible correspondence of the enthalpy of displacement curve with changing temperature was observed. The use of ITC provides an accurate method to measure the thermodynamic quantities to provide some insight into the mechanism behind the processes of micellization and adsorption at the solid/liquid interface. LA047954D