J . Phys. Chem. 1992,96, 1994-1998
1994
anthracene-terminated poly(N-( 1-naphthylmethy1)methacrylate-co-acrylic acid) was greatly enhanced in aqueous solution. For the copolymer containing 7.6 mol % naphthalene and 0.25 mol % anthracene end groups, x was 0.70 in aqueous solution while it was about 0.15 in d i ~ x a n e .In ~ this system, the average energy-transfer time in aqueous alkaline solution was shown to be 6.4 ns, while it was 11.3 ns in d i ~ x a n e . ~ Extremely rapid energy transfer observed for poly(A/Np(59)lPy) as compared with these earlier results may be attributed to its structural feature that the chromophore groups are attached to the backbone through amide spacer bonds. A high rotational barrier and a tendency for hydrogen bonding of the amide bond would restrict the conformational freedom of the chromophore groups, and therefore their local free motion in the polymer side chain may be considerably reduced. Thus, in the aggregate of the aromatic chromophores in aqueous solution, the hydrophobic interaction may bring them into very close proximity, but an orientation of the Np chromophores for excimer formation and/or self-quenching may be disfavored owing to the restriction to the conformation. Apparently, the extremely rapid energy transfer observed for poly(A/Np(S9)/Py) cannot be satisfactorily explained by the FBrster dipole-dipole mechanism for noninteracting discrete chromophores. Therefore, we may have to evoke other mechanisms of energy transfer such as the exchange mechanism and the exciton mechanism.14 In order for photoexcitation energy to be delocalized over a large space as an exciton, chromophores are required to be arranged regularly with a sufficiently small (14) Philpot, M. R. J . Chem. Phys. 1975, 63,485.
distance and a similar orientation. These requirements may not be met by vinylic polymers bearing chromophores as the pendant groups. However, in some special cases,there may be the regular arrangement of the chromophores at least locally, which makes an exciton state be able to exist. This may be the case for the chromophore cluster formed by poly(A/Np(59)/Py) in aqueous solution. However, a quantitative understanding of the mechanism of the rapid energy transfer requires more detailed howledge of the interchromophore distance and orientation in the hydrophobic aggregate.
Conclusions Extremely efficient and rapid excitation energy transfers from N p chromophores to Py energy traps were found to occur in the amphiphilic terpolymers of AMPS, 1-NpMAm, and 1-PyMAm in aqueous solution, where the terpolymers form a microphase structure. For a terpolymer consisting of 59 mol % N p chromophores and 1 mol 9% Py traps, the efficiency of singlet energy transfer in aqueous solution was estimated to be practically unity. On excitation of the N p chromophores at 290 nm, a decay time of 19 ps and the corresponding rise time of 20 ps were observed for N p fluorescence and for Py fluorescence, respectively. The role of efficient energy migration contributing to the energy transfer to the Py traps was discussed. These remarkable "photon-harvesting" properties found in the present terpolymer system may be attributed to the polymer structure in which the chromophores were linked to the main chain through rigid amide bonds. Registry No. (1 -NpMAm)( 1-PyMAm)(AMPS) (copolymer), 138333-44-5.
Calorimetric Observations of the Sphere-Rod Transition of Sodium Dodecyl Sulfate: Effects of Electrolytes and Nonelectrolytes at 25 C Duy Nguyen*Yt and Gary L. Bertrand Department of Chemistry, University of Missouri-Rolla, Rolla, Missouri 65401 (Received: June 17, 1991; In Final Form: October 18, 1991)
Incremental calorimetric titrations of small amounts of alcohols into solutions containing 3.5% (w/w)sodium dodecyl sulfate and added electrolytes reveal a sharp break in the partial molar enthalpy of the alcohol. The end of this break corresponds to the micellar sphererod transition, as observed by light scattering and other techniques. Unlike other techniques, however, calorimetric measurements show a well-defined beginning and end of the transition in some solutions. This technique has been used to determine the effects of various electrolytesand nonelectrolytes on the transition at 25 O C . The effect of electrolytes and nonelectrolytes is generally complementary,but dioxane and urea retard the transition while benzyl alcohol and 2-pentanol have little effect. The effects of electrolytes are determined primarily by the cations, decreasing in the series K+ > NH4+ > Na+ > Li+. The effect of alcohols increases with chain length and decreases with branching.
Introduction Micellar transitions at surfactant concentrations above the critical micelle concentration (cmc), sometimes called a second cmc, have been reported by a number of ~ 0 r k e r s . l - I ~Light scattering measurements1-I2 indicate a rapid increase in the average micellar molecular weight beyond this transition. This has been confirmed by other experimental methods.l3-I9 In the cases of sodium dodecyl sulfate (SDDS)2-7and tetradecyltrimethylammonium bromide ('ITMAB),'ov12the results have been interpreted as a sphere-rod transition but some authors2v3have Person to whom correspondence should be addressed. 'Present address: Betz Paperchem, Inc., 7510 Baymeadows Way, Jacksonville. FL 32256.
0022-365419212096- 1994$03.00/0
suggested that microcrystals may be formed in some cases. Added electrolytes and cosurfactants are generally found to facilitate the (1) Ozeki, S.; Ikeda, S. J. Colloid Interface Sci. 1982, 87, 424. (2) Hayashi, S.; Ikeda, S.J . Phys. Chem. 1981, 85, 106; 1980, 84, 744. (3) Corti, M.; Degiorgio, V. In Solution Chemistry of Surfactants; Mittal, K.L., Ed.; Plenum Press: New York, 1979. (4) Mazer, N. A.; Benedek, G. B.; Carey, M. C. J. Phys. Chem. 1976,80, 1075. (5) Young, C. Y.; Missel, P. J.; Benedek, G. B.; Carey, M. C. J . Phys. Chem. 1978,82, 1375. (6) Mazer, N. A.; Carey, M. C.; Benedek, G. B. In Micellizafion,Solubilization, and Microemulsion;Mittal, K. L., Ed.; Plenum Press: New York, 1977: Vol. 1. (7) Missel, P. J.; Mazer, N. A.; B:nedek, G. B.; Young, C. Y.; Carey, M. C.J . Phys. Chem. 1980,84, 1044.
0 1992 American Chemical Society
The Journal of Physical Chemistry, Vol. 96, No. 4, 1992 1995
Sphere-Rod Transition of Sodium Dodecyl Sulfate
TABLE I: PartiPl Molar Enthalpies of Solution (kJ/mol) of Benzyl Alcohol (e) in Aqueous Sodium Chloride with 3.5%SDDS at 25 O C [NaCI], m
[Bl, ma
0.10
0.20
0.25
0.30
0.35
0.40
0.50
0 0.01 0.03 0.05 0.08 0.10 0.12 0.14 0.16 0.18
(-1.61) -1.45 -1.17 -0.90 -0.53 -0.32 -0.12 0.05 0.20 0.34
(-1.51) -1.35 -1.05 -0.76 -0.38 -0.15 0.05 0.23 0.39 0.52
(-1.48) -1.32 -1.03 -0.76 -0.39 -0.16 0.05 0.25 0.42 0.58
(-1.58) -1.44 -1.15 -0.88 -0.49 -0.24 0.00 0.23 0.46 0.67
(-1.70) -1.62 -1.39 -1.11 -0.62 -0.28 0.05 0.34 0.58 0.74
(-2.15) -1.99 -1.63 -1.24 -0.63 -0.24 0.12 0.42 0.67 0.82
(-2.49) -2.17 -1.57 -1.06 -0.42 -0.06 0.24 0.50 0.72 0.92
Data smoothed with polynomial equations using least-squares regression. TABLX Ik Partial Molar Enthalpies of Solution (kJ/mol) of l-Pentnnol (P) in Aqueous Sodium Chloride with 3.5%SDDS at 25 O C
[NaCI], m -5
[PI, ma
0.10
0.20
0.25
0 0.005
(-5.43) -5.35 -5.20 -5.07 -4.94 -4.83 -4.73 -4.64 -4.5 1 -4.44, 4.38
(-5.24) -5.22 -5.16 -5.10 -5.05 -4.99 -4.93 -5.01 -5.03 -4.92 -4.74
(-5.60) -5.33 -5.12 -5.17 -5.35 -5.49 -5.47 -5.39 -5.09 -4.88 -4.66
0.015 0.025 0.035 0.045 0.055 0.065 0.080 0.090 0.100
b I I
0X y -5 3
0
m
Ob
0.30 (-5.1 6) -5.29 -5.53 -5.75 -5.95 -5.74 -5.46 -5.18 -4.77 -4.51 -4.25 (-7.1)
0.40
0.50
-6.38 -6.61 -6.11 -5.66 -5.27 -4.93 -4.64 -4.31 -4.16 -4.05 (-7.5)
-6.59 -5.91 -5.38 -4.97 -4.67 -4.44 -4.27 -4.06 -3.92 -3.74 (-7.0)
'Data smoothed with polynomial equations using least-squares regression. Extrapolated from data beyond transition.
-6
0
.05
.I
mi p e n t a n o i l Figure 1. Partial molar enthalpy of solution of 1-pentanolin 3.5% SDDS for different concentrations of sodium chloride at 25 'C.
transition, but Missel et al?O found that urea retards the transition of SDDS in 0.8 M sodium chloride. Our first observation of the effects of this transition occurred in a study of the enthalpy of solution of benzyl alcohol in aqueous solutions containing 3.5% SDDS and added sodium chloride. There exists a linear relationship between the concentration of benzyl alcohol and the partial molar enthalpies of solution, allowing extrapolation to infinite dilution to obtain the standard enthalpy of solution. The standard enthalpy of solution (Table I) shows (8) Mazer, N. A. In Dynamic Light Scattering, Pecora, R., Ed.; Plenum Press: New York. 1985. (9) Ikeda, S.;Ozeki, S.;Tsunoda, M. A. J. Colloid Interface Sci. 1980, 73,-27. (IO) Debye, P.; Anacker, E. W. J. Phys. Colloid Chem. 1951, 55, 644. (1 1) Imae, T.; Kamiya, R.; Ikeda, S.J. Colloid Interface Sci. 1985, 108, 215. (12) Imae, T.; Ikeda, S.J. Phys. Chem. 1986, 90, 5216. (13) Liana, P.; Zana, R. J. Phys. Chem. 1980,84, 3339. (14) Staples, E. J.; Tiddy, G. J. J . Chem. SOC.,Faraday Trans. I 1978, 74, 2530. (15) Franses, E. J.; Davis, H. T.; Miller, W. G.; Scriven, L. E. J. Phys. Chem. 1980,84, 2413. (16) Almgren, M.; Lofroth, J. E. J. Colloid Interface Sci. 1981.81, 486. (17) Birdi, K. S.;Dalsager, S.U. J. Chem. SOC.,Faraday Trans. I 1980, 76, 2035. (18) Henriksson, U.; Odberg, L.; Eriksson, S. C.; Westman, L. J. Phys. Chem. 1911,81,16. (19) Vikholm, I.; Douheret, G.; Backlund, S.;Hoiland, H. J. Colloid Interface Sci. 1987, 116, 582. (20) M i d , P. J.; Mazer, N. A.; Carey, M. C.; Bcnedek, G. B. In Solution Behauior of Surfactants; Mittal, K. E., Fendler, E. J., Ed.; Plenum Press: New York, 1982; Vol. I .
a sharp exothermic break beginning between 0.25 and 0.3 m sodium chloride and ending between 0.4 and 0.5 m, while the standard enthalpy of solution of benzyl alcohol in sodium chloride solutions without SDDS increases fairly linearly with salt concentration to a t least 1 m. Mazer et aL4 have reported that the sphererod transition for 2% SDDS occurs at 0.45 M sodium chloride, corresponding to the end of the break shown in Table I. In order to confirm that this break in the enthalpy of solution curve is a property of the micellar system rather than the solute, the study was repeated with 1-pentanol as the solute. In the solution containing 0.1 m sodium chloride, the partial molar enthalpy of solution increased smoothly with concentration of pentanol, as was observed for benzyl alcohol. However, for higher concentrations of salt, an exothermic break is observed in the partial molar enthalpy of solution curve (Figure l), becoming more pronounced and occurring at lower pentanol concentrations as the salt concentration is increased. Below 0.3 m sodium chloride, these curves show a linear region of positive slope before the break, a negative slope through the break, then a positive slope decreasing with concentration of pentanol after the break. In 0.4 m sodium chloride, the initial region of positive slope is not observed. In 0.5 m sodium chloride, the break is not observed and the partial molar enthalpy of solution of pentanol increases smoothly with decreasing slope as the concentration of pentanol increases. For solutions with salt concentrations below 0.3 m,the initial region of positive slope was extrapolated to obtain the standard enthalpies of solution given in the top row of Table 11. For higher concentrations of salt, the later region of positive slope was extrapolated with nonlinear equations to obtain the values shown in the bottom row. While considerably different from the results for benzyl alcohol in both shape and magnitude, the break for 1-pentanol occurs in the same region of concentration of sodium chloride. The sharp break in the standard enthalpy of solution as a function of salt concentration is thus the same for two different solutes and appears to be a property of the aqueous surfactant plus salt solution.
1996 The Journal of Physical Chemistry, Vol. 96, No. 4, 1992
Zana2'*22studied the sphere-rod transition of TTMAB in mixtures of 1-pentanol and potassium bromide by light scattering techniques and found that the two have cooperative effects for inducing the transition. Through comparison of the e n t h a l p i ~ ~ ~ and light scattering studies on this system, we have concluded that the end of the break observed calorimetrically corresponds to the sphere-rod transition observed with light scattering. Light scattering responds to the average size or molecular weight of the micelles, while calorimetry responds directly to molecular changes. As the first small rods are formed, there is an immediate enthalpic effect but the light scattering response is very small and does not become apparent until a substantial number of large rods have been formed. The enthalpy effect appears to be different for spheres and rods but apparently is not greatly affected by the size of the rods. An abrupt change is thus observed over a region of compositions from the first appearance of rods until the fraction of surfactant molecules in spherical micelles is very small. The break in the partial molar enthalpy of solution of 1-pentanol vs pentanol concentration has been observed in both SDDS and TTMAB23 solutions containing a variety of electrolytes and nonelectrolytes. Similar effects were observed for 3-methyl- 1butanol and 1-hexanol but not for benzyl alcohol or 2-pentanol in solutions containing SDDS and salts. Attempts to observe the transition by adding concentrated sodium chloride solution to SDDS solutions containing 0.4-0.5 m sodium chloride were unsuccessful. However, when this experiment was repeated with a small amount of 1-pentanol present, a slight break was observed. From these observations, we conclude that the observed enthalpy effect is not due to different enthalpic states of SDDS molecules in spheres and rods but due to changes associated with the partition of some cosurfactants between the bulk solution phase and the spherical and rodlike forms of the micellar phase.
Experimental Section Materials. Sodium dodecyl sulfate (BDH Chemicals, 99+%) was twice recrystallized from ethanol and dried under vacuum. Urea and salts were from Fisher (ACS Certified Reagent Grade, when available). 1-Pentanol (99%) was from Sigma Chemical. 1-Hexanol (98%), 3-methyl-1-butanol (98%), 2-pentanol (98%), 1-hexylamine (99%),pdioxane (99+%),and benzyl alcohol (99%) were from Aldrich. Preparation of Solutions. Solutions of electrolytes and urea were prepared in 0.5-1 .O-L quantities by weight such that molalities could be calculated to better than 0.1%. These solutions were used to prepare SDDS solutions of approximately 100 mL containing 3.50 f 0.01% (w/v) of surfactant. These solutions were weighed into the reaction vessel, and then any cosurfactants to be added were pipetted directly into the reaction vessel. Calorimetry. Isoperibol measurements were performed on Tronac Models 450 and 550 titration calorimeters with the bath at 25.00 f 0.01 OC, maintained constant to 0.0002 OC. Successive increments (0.05-0.2 mL) were titrated into 100-mL reaction vessels from 2.5-mL Gilmont micrometer burets readable to O.OOO1 mL. The heater was operated exactly as in a measurement of heat capacity, and the net enthalpy effect was calculated as the sum of the electrical work input, the temperature displacement multiplied by the effective heat capacity, and the small correction for the difference in temperature between the solution and the bath. In the case of endorthermic measurements, the calibrating heater was used to compensate for the thermal effect such that the net tempersture change was very small. In the case of exothermic measurements, there is an accumulation of temperature increases, and in some cases it was neceSSary to remove the reaction vessel from the bath for cooling after several increments had been added in order to keep all measurements within the range of 25.0 f 0.2 O C . Contraction of the solute in the buret during this cooling period causes mixing of solute and solution in the delivery tube
Nguyen and Bertrand TABLE 111: Partial Molar Enthalpies of Solution (W/md) of Alcohols (A) in Aqueous Sodium Chloride with 3.5%SDDS at 25 OC [A], 3-methyl2111ma 1-butanolb pentanolb hexanolb hexanolC hexanold 0.005 0.015 0.025 0.035 0.045 0.055 0.065 0.080 0.090
0.100
-5.33 -5.25 -5.24 -5.28 -5.33 -5.36 -5.38 -5.18 -5.04 -4.91
-8.49 -8.44 -8.39 -8.34 -8.29 -8.24 -8.19 -8.12 -8.07 -8.02
-4.59 -5.44 -6.30 -5.63 -4.89 -4.13 -3.38
-4.39 -4.14 -3.95 -3.80 -3.81 -4.01 -4.39 -4.29 -4.03 -3.65
"Data smoothed with polynomial equations using least squares regression. [NaCI] = 0.25 m. [NaCI] = 0.20 m. d[NaC1] = 0.10 m. TABLE I V Partial Molar Enthalpies of Solution (kJ/mol) of 1-Pentanol (P) in 0.25 m Salt Solutions with 3.5%SDDS at 25 O C [P1,ma NaBr NaF NaNOI NaSCN NaCl 0.005 0.015 0.025 0.035 0.045 0.055 0.065 0.080 0.090 0.100
-4.93 -5.1 1 -5.27 -5.4 1 -5.53 -5.63 -5.64 -5.24 -5.02 -4.19
-5.20 -5.08 -5.14 -5.29 -5.43 -5.48 -5.33 -5.02 -4.78 -4.54
-5.19 -5.08 -5.08 -5.13 -5.21 -5.27 -5.28 -5.07 -4.87 -4.64
-5.18 -5.12 -5.16 -5.25 -5.35 -5.40 -5.39 -5.10 -4.86 -4.62
-5.33 -5.12 -5.17 -5.35 -5.49 -5.47 -5.39 -5.09 -4.88 -4.66
Data smoothed with polynomial equations using least-squares regression.
so the measurement immediately following this period was discarded. Incremental enthalpies of solution were divided by the number of moles of solute added (calculated from the volume increment using densities from Timmerman~~~). These enthalpy values have been shown25 to closely approximate the partial molar excess enthalpy of the solute at the average of the concentrations before and after the addition. Results Smoothed results for partial molar enthalpies of solution of benzyl alcohol and 1-pentanol in aqueous sodium chloride solutions containing 3.5% SDDS are given in Tables I and 11. These results were discussed earlier in this paper. For comparison of the effects of different cosurfactants (Table 111), the 0.25 m solution of sodium chloride was chosen, since this concentration gives a welldefined start and end of the break for 1-pentanol. The break for 3-methyl-1-butanol begins and ends at slightly higher alcohol concentrations than that for 1-pentanol, and no break was observed for 2-pentanol, showing that branching of the alcohol reduces the ability of the alcohol to induce the transition. 1-Hexanol exhibits an earlier, more pronounced break than 1-pentanol, showing that the longer alcohol induces the transition more effectively. The effect of different monovalent eo-ions at a concentration of 0.25 m is shown in Table IV. The shapes of the enthalpy of solution curves vary somewhat in the transition region for the different sodium salts, but the concentrations of 1-pentanol at which the transition begins and ends are essentially the same for all of these salts, indicating that a monovalent eo-ion has little effect on the transition. The addition of 0.125 m sodium sulfate causes the partial molar enthalpy of solution of 1-pentanol to flatten out through and beyond the transition region observed with the monovalent salts at the same concentration of sodium ion and then essentially follows the same curve as the monovalent salts ~~
(21) Candan, S.; Zana, R. J . Colloid Interface Sci. 1981, 84, 206. (22) Zana, R.; Yiv, S.; Strazielle, C.; Lianos, P. J . Colloid Interface Sci. 1981, 80, 208. (23) Duy Nguyen and Gary L. Bertrand, University of Missouri-Rolla, Research in progress.
-4.71 -4.78 -5.42 -6.63 -5.94 -5.08 -4.28 -3.28 -2.83
~
~~~
(24) Timmermans, J. Physico-chemical Constants of Pure Organic Compounds; Elsevier: New York, 1950. (25) Taylor, E. L.; Bertrand, G. L. J . Solution Chem. 1914, 3, 419. (26) Kunin, R.; Myers, R. J. Ion Exchange Resins; John Wiley and Sons: New York, 1950. (27) Stilbs. P. J . Colloid Interface Sci. 1982, 87, 385.
The Journal of Physical Chemistry, Vol. 96, No. 4, 1992 1997
Sphere-Rod Transition of Sodium Dodecyl Sulfate
TABLE V: Concentrations (molal) for Transition of 3.5%SDDS with Alcohols and Salts (Values in Parentheses Are for Start of Transition When ObseNed) [3-methyl-1 [NaCI] [ 1-pentanol]" [LiCI] [ 1-pentanol]" [NHICI] [ 1-pentanol]" [NaCI] [ 1-hexanol]" [NaCI] butanol]"
-
(0.070) 0.093* 0.200 0.250 0.300 0.400
0.205 0.080 (0.058) 0.055 (0.018) 0.035
0.500 0.550 0.700 0.900
0.065 (0.030) 0.055 (0.027) 0.037 0.018
0.100 0.150 0.200
0.073 (0.040) 0.058 (0.025) 0.037
0.100 0.200 0.250
0.072 (0.041) 0.035 (0.016) 0.025
0.25
0.064 (0.017)
0.010
"The titrant is the nonelectrolyte (alcohol) unless otherwise noted. *The titrant is the electrolyte (NaC1). TABLE VI. Concentrations (molal) for Transition of 3.5%SDDS with 1-Pentawl (P) and Two Salts at 25 OC
[PI 0.050 0.042 0.041 0.033
[NaCI] 0.238 0.225 0.219 0.208
[KCI] 0.012 0.025 0.031 0.042
[PI 0.048
[NaCI] 0.208
[NH4CI] 0.042
at concentrations of pentanol above 0.1 1 m. Graphs of the partial molar enthalpy of solution vs concentration of titrant were constructed, and the fairly linear portions of the curves after, during, and (when possible) before the transition were extrapolated to crossing points which were taken to be the composition at the end and/or start of the transition. Concentrations of the various additives at these transition points are given in Tables v-VII. Variation of the counterion has a pronounced effect on the partial molar enthalpy of solution of 1-pentanol (Table V). Much greater concentrations of lithium chloride relative to sodium chloride are required to cause the transition to occur at similar concentrations of 1-pentanol, indicating that the lithium ion is much less effective than sodium ion for promoting the transition. However, when sufficient quantities of these salts are added to cause the transition at the same concentration of pentanol, the partial molar enthalpy of solution curves are very similar, indicating that the enthalpy of transition observed in this manner is more dependent on the concentration of cosurfactant than on the concentration of electrolyte. Much lower concentrations of ammonium chloride relative to sodium chloride are required to cause the transition at similar concentrations of 1-pentanol. The shapes of the partial molar enthalpy of solution curves are somewhat different at the same concentration of pentanol, perhaps due to pH effects from hydrolysis of the ammonium ion. SDDS is not sufficiently soluble in potassium chloride to allow a similar study with this salt. To determine the relative effectiveness of the potassium ion relative to sodium ion, mixtures of sodium chloride and potassium chloride at a total concentration of 0.25 m were studied, with the concentration of the potassium ion ranging from 0.012 to 0.042 m (Table VI). The concentration of 1-pentanol required to cause the transition decreased substantiably with increased concentration of potassium, indicating that the potassium ion is considerably more effective than the sodium ion for promoting the transition. A similar study with mixtures of sodium chloride and ammonium chloride gave similar results, indicating that the ammonium ion is slightly less effective than the potassium ion for promoting the transition. The effects of added nonelectrolytes on the amounts of 1pentanol and 1-hexanol required to cause the transition are given in Table VII. Addition of urea and pdioxane increases the amount of pentanol required, with p-dioxane showing considerably more effectiveness than urea for retarding the transition. Addition of
1-hexylamine greatly reduces the amounts of both alcohols required, indicating that this cosurfactant is more effective than either of the alcohols for promoting the transition. Addition of 1-pentanol reduces the amount of hexanol required, while benzyl alcohol and 2-pentanol may slightly increase the amount of 1hexanol required, though the effect is barely a t our level of observation. In all of these measurements, added nonelectrolytes caused the partial molar enthalpy of solution of the alcohols to become less exothermic before the transition and reduced the magnitude of the effect of the transition. After the transition, the partial molar enthalpies of the alcohols as a function of titrant concentration were very similar to results without the added nonelectrolyte. Discussion
In this study, we were unable to observe the sphere-rod transition of SDDS without the combined presence of an electrolyte and a nonelectrolyte, thus preventing a direct comparison of the calorimetric results to light scattering studies involving only an electrolyte. However, extrapolation of our results to zero concentration of 1-pentanol allows this comparison. Unfortunately, this extrapolation is compromised somewhat by the nonlinear relationship observed between concentrations of salt and alcohol (Table V). Using quadratic expressions for the concentration of salt as a function of 1-pentanol concentration, we obtain the following molalities of salts which would cause the transition in 3.5% SDDS: NaCl, 0.455 f 0.005; LiCl, 1.16 f 0.02; NH4Cl, 0.35 f 0.05. This value for NaCl is in good agreement with the value of 0.45 M (0.455 m) for 2% SDDS reported by Mazer et ale4from light scattering techniques. Ikeda et a1.* observed that the sodium salts of other anions showed different aggregation numbers after the transition but had little effect on the concentration of salt at which the transition occurred. This is in qualitative agreement with our observation of somewhat different enthalpic effects during the transition for various sodium salts but little effect on the concentration of 1-pentanol at the end of the transition for a particular salt concentration. The effectiveness of cations in promoting this transition decreases in the order K+ > NH4+> Na+ > Li+, which is the order of increasing hydrated radius of ions or decreasing ionic crystal radius. This order was also observed by Missel et a1.20 The effectiveness of the nonelectrolytes studied decreases in the order of 1-hexylamine > 1-hexanol > 1-pentanol > 3-methyl-1-butanol > 0 benzyl alcohol 2-pentanol > urea > p-dioxane. These effects very closely parallel the effects of these additives on the cmc of SDDS, in that additives which lower the cmc also promote the sphererod transition.
-
-
Conclusions
The mechanism of the sphere-rod transition via calorimetry is not fully understood. As clearly stated in the title, the purpose
TABLE VII: Concentrations (molal) for Transitions of 3.5%SDDS with Sodium Chloride (E) and Two Nonelectrolytesat 25 OC
[E] [I-pentanol] [p-dioxane] [El [I-pentanol] [urea] [E] [I-pentanol] [ 1-hexylamine] 0.40 0.040 0.200 0.40 0.020 0.100 0.25 0.042 0.0085 0.40 0.080 0.500 0.40 0.042 0.500 [El [I-hexanol] [ 1-pentanol] [El [ 1-hexanol] [benzyl alcohol] [E] 0.10 0.058 0.050 0.10 0.080 0.052 0.10 0.10 0.049 0.099 0.10 0.077 0.105 0.10 0.20
[E] 0.10
[I-hexanol] (2-pentanoll 0.075 0.099
[ 1-hexanol]
[ 1-hexylamine]
0.045 0.03 1 0.0135
0.0084 0.0168 0.0101
1998
J. Phys. Chem. 1992, 96, 1998-2006
of this paper is to present a new technique of observing this transition, a technique which is quicker, cheaper, and more direct than any other (e.g. light scattering). Also, we present the first evidence that this is not a sharp, single-point transition but one which occurs over a range of composition with a fairly distinct beginning and end (in some cases). Using this technique, we have been able to catalog the effects of a number of compounds on the transition. We briefly summarize four types of effects: counterions, co-ions, cosurfactants, and cosolvents (organic modifiers). While to some readers this may invite discussions of mechanisms and causal effects suggested by Ruckenstein, Ninham, and I~raelachvili,~*-~~ we are well aware (28) Nagarajan, R.; Ruckenstein, E. J. Colloid Interface Sci. 1977, 60, 221. (29) Nagarajan, R.; Ruckenstein, E. J. Colloid Interface Sei. 1979, 71, 580. (30) Israelachvili, J. N.; Mitchel, D. J.; Ninham, B. W. J. Chem. Soc., Faraday Trans. 2 1976, 72, 1525.
that these are observations on a single surfactant at a single concentration at a single temperature. Until we have convinced the scientific community of the validity of our observations, any discussion of mechanism would be speculation. Registry No. NaC1, 7647-14-5; NaBr, 7647-15-6;NaF, 7681-49-4; NaNO,, 763 1-99-4; NaSCN, 540-72-7; KCI, 7447-40-7; LiC1, 744741-8; NH,Cl, 12125-02-9;sodium dodecyl sulfate, 151-21-3; 1-pentanol, 71-41-0; 1-hexanol, 111-27-3; 3-methyl-I-butanol,123-51-3;2-pentanol, 6032-29-7; 1-hexylamine,11 1-26-2;p-dioxane, 123-91-1;benzyl alcohol, 100-51-6;urea, 57-13-6. (31) Ruckenstein, E.; Chi, J. C. J. Chem. SOC.,Faraday Trans. 2 1975, 71, 1690. (32) Ruckenstein, E. Chem. Phys. Lett. 1978, 57, 517. (33) Ruckenstein, E. Chem. Phys. Lett. 1983, 98, 573. (34) Israelachvili,J. N. In Physics of Amphiphiles: Micelles, Vesicles and Microemulsions; Proceedings of the International School of Physics "Enrico Fermi", 1985; Degiorgio, V., Corti, M., Eds.; North-Holland: Amsterdam,
1986.
Application of Tanford's Micellization Theory to Gel Filtration Chromatographic Data for Nonionic Surfactants Noriaki Funasaki,* HyangSook Shim, and Sakae Hada Kyoto Pharmaceutical University, Misasagi, Yamashina-ku, Kyoto 607, Japan (Received: March 19, 1991; In Final Form: September 25, 1991)
The micelle formation of hepta(ethy1eneglycol) decy1 ether ( C I A ) has been investigated by frontal gel filtration chromatographic (GFC) and surface tension methods at 25 'C and explained in the framework of the Tanford theory. From an analysis of the concentration dependence of the centroid volume of the GFC pattern, the monomer concentration and weight- (n,)and number- (n,) average aggregation numbers of CI0E7micelles are estimated as a function of the total CloE7concentration, The n, value increases rapidly above the critical micellization concentration (cmc) and levels off at higher concentrations. The ratio n,/nn has a maximum around the cmc and approaches unity at higher concentrations. The derivative GFC pattern suggests the formation of premicelles. The dimerization constant of CI0& is smaller than that of Cl&. These experimental results are quantitatively explained on the basis of many micellization models, originally considered by Tanford. These models are different in the configurations of the decyl and hepta(oxyethy1ene)chains, the penetration of water into the micelle core, the roughness of the micellar surface, and the micelle shape. A most probable CIOE7micelle of n, = 60, conforming to the present and literature data, seems to be an oblate ellipsoid whose minor semiaxis is the sum of the length of the fully extended decyl chain and that of the randomly coiled hepta(oxyethy1ene) chain. The micelle size distribution function is calculated on the basis of several micellization models.
Introduction It is well-known that a surfactant forms micelles above the critical micellization concentration (cmc) in aqueous solution, but little is yet known about the formation of premicelles (dimers and oligomers) and the relation of micelle size and shape to the chemical structure of the s~rfactant.I-~For the investigation of these problems, the Tanford theory is of most promise among a number of theories for micelle formation.'-3 Tanford related micelle size, the cmc, and the other micellar properties to a size-dependent free energy of micellization which is a function of the chemical structure of the surfactant. Then he used the definition of the cmc as the total concentration x at which the monomer concentration x, is 95% of x, although a better definition of the cmc would be x at which d3xl/dx3= 0.4-7 Furthermore, (1) Wennerstrom, H.; Lindman, B. Phys. Rep. 1979, 52, 1. (2) Tanford, C. J. Phys. Chem. 1974, 78, 2469.
(3) Tanford, C. The Hydrophobic Effect; John Wiley: New York, 1980; Chapters 2-8. (4) Ben-Naim, A,; Stillinger, F. H. J. Phys. Chem. 1980, 84, 2872. (5) Warr, G.G.;White, L. R. J. Chem. SOC.,Faraday Tram. 2 1985,81, 549.
0022-3654/92/2096-1998$03.00/0
he could reproduce weight-average micellar aggregation numbers
n,at the cmc on the basis of his theory. The agreement between theory and experiment for nonionic surfactants is poorer than that for ionic and zwitterionic surfactants. The values of n, and x , should both depend on x , but only those values at the cmc were compared with theoretical v a l ~ e s . ~He, ~considered that no premicelle forms below the cmc3 Recently we have developed a gel filtration chromatographic (GFC) method for the determination of n, and x l as a function of x.&Io For hexa(ethy1ene glycol) decyl ether (CloE6),we have shown that the concentration dependence of n, determined by GFC is in good agreement with that by static light ~cattering.~ For octa(ethy1ene glycol) decyl ether (CloE8),we have shown that premicelles (mainly dimers) form and that micelle size grows rapidly with increasing concentration above the cmc, as compared (6) Funasaki, N.; Shim, H.-S.; Hada, S. J. Chem. Soc., Faraday Trans. 1991, 87, 957. (7) Funasaki, N.; Hada, S . Bull. Chem. SOC.Jpn. 1991,64, 682. (8) Funasaki, N.; Hada, S.;Neya, S . J. Phys. Chem. 1988, 92, 7112. (9) Funasaki, N.; Hada, S.; Neya, S. J. Phys. Chem. 1990, 94, 8322. (IO) Funasaki, N.; Hada, S.; Paiement, J. J. Phys. Chem. 1991,95,4135, and references cited therein.
0 1992 American Chemical Society
