Langmuir 1986,2, 780-786
780
0
'
2 103x~
Figure 5. Energy of interface formation vs. concentration curves at constant temperature: (1)303.15; (2) 298.15 K. of state which often accounts well for the behavior of nonideal monolayers: T ( A - A') = RT (8) where Ao is known as the excluded area. When this equation is applied to the x and A values, at large area the l / vs. ~A plot is observed to exhibit a linear relation which extrapolates to the origin. This fact suggests that the lateral interaction between cholesterol molecules is negligible in the gaseous film and the state is effectively treated as an ideal two-dimensional gas. The entropy and energy of interface formation of the gaseous film decreases gradually with increasing bulk concentration but never becomes negative as shown in Figures 4 and 5. These results indicate that the transfer of cholesterol from the solution to the interface is accompanied by small negative changes in partial molar entropy and energy. The negative change of partial molar energy suggests that the adsorption a t the interface is an energetically favorable process. However, the thermodynamically unfavorable process of the interface formation is not compensated sufficiently by the gaseous film formation at the interface. The expanded state differs somewhat from the gaseous state in thermodynamic quantities. The entropy of interface formation is almost zero, the positive value of a
pure bemenelwater interface being compensated by the negative change in the partial molar entropy of cholesterol. The energy of interface formation still exhibits a large positive value. When eq 8 is applied to the expanded film, the Ao value is estimated to be 0.42 nm2 by the extraporation of 1 / vs. ~ A plots. Then the lateral interaction between cholesterol molecules a t the interface cannot be negligible in the expanded state, but the linear plot of l / ~ vs. A suggests that the behavior is not so different from that of an ideal two-dimensional gaseous state. It seems important to notice that the above behavior of cholesterol is extremely different from the behavior of 1-octadecanol in the aspect that a phase transition does not take place in the adsorbed film at the benzene/water interface. It is seen from Figure 4 that the condensed state is characterized by a large negative value of the entropy of interface formation. Furthermore, the energy of interface formation also has a large negative value which is shown in Figure 5. This fact may indicate that the cholesterol molecules are closely packed together in the condensed film. However, it should be noted that these As and ALL values are one-tenth of the corresponding ones of l-octadecanol at the hexanelwater interface, respectively. By making use of the values given in Table I, the thermodynamic quantities associated with the phase transitions can be evaluated numerically. It is found that the absolute values determined for the expanded/condensed phase transition are remarkably larger than those for the gaseous/expanded phase transition. It is important to note that they differ significantly from the corresponding values of 1-octadecanol at the hexane/water interface. The distinction in the interfacial behavior between cholesterol and 1-octadecanol leads us to the conclusion that the lateral interaction of cholesterol molecules is different from that of hydrocarbon chains. It is of great interest that the cohesive interaction between cholesterol molecules exerts a striking condensing effect when the interface density is relatively high, while it does not when the interfacial density is small. Further, cholesterol and 1-octadecanol molecules may be predicted to form a nonideal mixed film. Registry No. Cholesterol, 57-88-5;benzene, 71-43-2.
ESR Spectroscopic Investigation on the Exchange Kinetics of Surfactants in Water/AOT/Isooctane Microemulsions Alessandra Barelli and Hans-Friedrich Eicke* Institut fur Physikalische Chemie, Universitat Basel, CH-4056 Basel, Switzerland Received May 13, 1986. I n Final Form: August 7, 1986 This investigation is concerned with the surfactant exchange kinetics in three-component water/ AOT/ismtane, thermodynamically stable W/O microemulsions, using ESR spectroscopy. Nitroxide-labeled spin probes of different chemical nature were solubilized in the aqueous nanodroplets, yielding valuable information on dynamic (exchange kinetics) and static (solubilization sites) properties of water-in-oil microemulsions.
It appears striking in view of the intense research on the formation of the so-called water-in-oil (W/O) microemulsions, and particularly on the solubilization phenomenon, that relatively little is known about the kinetics of 0743-7463/86/2402-0780$01.50/0
these processes. The exchange of surfactants and cosurfactants between dispersion medium and aggregates, on the one hand, and the exchange processes within the aqueous nanometer-sized droplets (Le., nanodroplets) 0 1986 American Chemical Society
Exchange Kinetics of S u r f a c t a n t s i n Microemulsions
covered by a surfactant monolayer which forms the W/O microemulsion, on the other, are equally interesting. The main difficulty which hampers the investigation addressed above is the restricted time domain within which such exchange processes o c ~ u r . l - ~This leaves as one of the few possible techniques electron spin resonance (ESR) spectroscopy, the application of which immediately presents one to another problem, i.e., the synthesis of suitable spin probes. The latter turns out to be far more difficult in W/O microemulsions; the introduction of doxyl nitroxides, selected as the most stable and synthetically easily accessible radical^,^ into some of the proven surfactants for W/O microemulsions strongly affects their properties, e.g., their amphiphilicity and ~olubility.~ If suitable spin probes are synthesized, the ESR technique is able to provide, besides the information on solubilization sites, rate constants of surfactant exchange processes.6-11 The basic idea of this method as applied to microemulsion rests on the assumption that the introduction of a small amount of spin-labeled (SL) additives into the nanodroplets does not substantially alter the aggregational pattern of the system. It has been repeatedly confirmed (e.g., ref 12) that the mean size of the nanodroplets does not change with the amounts of additive used in this work. Hence, exchange rates of the SL molecules and some physical properties of the environment of the probe molecules (actually of the radical group) may be inferred from such experiments. Apart from the experimental difficulties encountered with this method, the evaluation of the data and their interpretation is not t r i ~ i a l . ~All J ~this might explain the lack of more kinetic information on W/O microemulsion. We have selected the thermodynamically stable threecomponent water/AOT/isooctane W/O microemulsion as a reference system to which SL compounds were added. In particular, we have synthesized two types of spin-labels represented by the structures I and 11. I
O\ ,N-0 CH3-C-R a ) R I (CH2),CH3
a) R = O H
b) R = CH,CH2CH(CH3)2
b) R. OS03Na
We are particularly concerned with the effect of the nanodroplet size (amount of bulk water as distinguished from bound water of the aqueous core) on the exchange rates of spin-labels between (i) nanodroplets and dispersion (1) Fendler, J. H.; Fendler, E. J. Catalysis in Micellar and Macromolecular Systems; Academic Press: New York, 1975.
(2) Aniansson, E. A. G.; Wall, S. N.; Almagren, M.; Hoffmann, H.; Kielmann, I.; Ulbricht, W.; Zana, R.; Lang, J.; Tondre, C. J.Phys. Chem. 1976, 80, 905. (3) Israelachivili, J. N.; MarEelja, S.; Horn, R. J. Quart. Reu. Biophys. 1980, 13, 121. (4) Keana, J. F. Am. Chem. Soc., Chem. Ref. 1978, 37. (5) Benatar-Barelli, A. Ph.D. Thesis, Basel University, Basel, 1985. (6) Spin Labeling: Theory and Applications; Berliner, L. J., Ed.; Academic Press: New York, 1976, 1979; Vols. I and 11. (7) Atherton, N. M.; Strach, S. J. J. Chem. SOC.,Faraday Trans. 2 1972, 374. (8) Fox, K. K. Trans. Faraday SOC.1971, 67, 2802. (9) Boglioni, P.; Ottaviani, M. F.; Martini, G; Ferroni, E. In Surfactants in Solution; Mittal, K. L., Lindman, B., Eds.; Plenum Press: New York, 1984; Vol. I. (10) Yoshioka, H.; Kazama, S. J. Colloid Interface Sci. 1983, 95, 240. (11) Schmidt, D.; Gahwiller, C. H.; Von Planta, C. J. Colloid Interface Sci. 1981, 83, 191. (12) Kim, V.; Hilfiker, R.; Eicke, H.-F. Colloids Surf. 1986, Oct.
Langmuir, Vol. 2, No. 6, 1986 781 Table I. Spectroscopic g Factors (fl X lo-*) and Hyperfine Coupling Constants (aN f 4 X for Different SL Compounds a t Room Temperature SL comaound solvent P factor uNlmT Ia isooctane 2.0062 1.403 Ib isooctane 2.0062 1.397 1.414 isooctane 2.0057 II(0,4)a water 2.0052 1.575 II(0,4)b water 2.0052 1.571 ~
medium and (ii) the amphiphilic monolayer and water core within the nanodroplets.
Experimental Section ( a ) Method. Electron spin resonance (ESR) spectra were recorded on a Varian E-9 X-band ESR-spectrometer equipped with a variable-temperature accessory. The temperature in the cavity was measured with a digital thermometer (Digitemp, No. 3995, West Germany) and was found to be stable within fl K. T h e ESR samples consisted of ternary mixtures of HzO, AOT (sodium bis(2-ethylhexyl) sulfosuccinate), and isooctane (2,2,4trimethylpentane, i-C8H18)a t constant AOT concentration and variable amounts of HzO. The samples also contained a minimal amount (about l.105mol dm-3) of doxyl nitroxide radicals as spin probes. The molar ratio of radical to AOT was about 3 x A suitable amount of the sample was then transferred to the ESR cell. The cell was a Pyrex glass tube of 2-mm internal diameter connected to a degassing device, whereas for ESR measurements in H 2 0 a glass capillary was used instead. The ESR probes were degassed by means of the freeze-thaw technique and the air was replaced by dry nitrogen gas. The water samples were deoxygenated by bubbling dry nitrogen through the solutions for 30 s. The spectroscopic g factors were determined by using 1,3bis(diphenylenyl)-2-phenylallylradical [g = 2.0026 f (1X lo4)] as a standard. The hyperfine coupling constants (aN) were measured between the central- and low-field lines of the ESR spectra. (b) Materials. HzO was deionized and doubly distilled. Isooctane was of highest quality (puriss. p.a.) from Fluka (Buchs, Switzerland) and was distilled over CaH, before use. The AOT (Fluka, 98% pure, pharmaceutical grade) was purified as described in ref 12. The purified AOT was spectroscopically identical with synthetic AOT.' Synthesis of Ia a n d Ib. These compounds were synthesized according to the general procedure described by Keana et al.4 from the respective ketones: 5-methyl-2-hexanone (Fluka, purum) and 2-dodecanone (Merck, for synthesis). Ia and Ib were identified by NMR spectroscopy of the respective N-hydroxy derivative^'^ after purification by preparative TLC on silica gel G using hexane-diethyl ether (l:l).5The g factors and uN values of Ia and Ib are given in Table I. Synthesis of II(0,4)a a n d II(0,4)b. These compounds were synthesized on the basis of the procedure described by Waggoner e t al.14 from ethyl 5-ketohexanoate (96% pure, Aldrich). The
2-methyl-2-[4-(ethoxycarbonyl)propyl]-4,4-dimethyloxazolidineN-oxy1 intermediate was eluted through a column of activated alumina with hexane-diethyl ether (1:l)and was identified by IR and NMR of the N-hydroxy derivative.13 UV (i-C8HI8): 441 ( E 5.4), 230 (e 2367.8), shoulder a t 209 nm. Mass spectrum, m / e 244 (molecular ion).5 Reduction of this nitroxide ethyl ester with 1 equiv of LiAlH4 in diethyl ether yields II(0,4)a as a yellow oil. After elution through a column of activated alumina with diethyl ether, II(0,4)a was identified by IR and NMR of the respective N-hydroxy d e r i ~ a t i v e . ~UV-visible J~ (i-C8H18):439 (t 5.8), 212 ( t 1666.7), shoulder at 226 nm. Mass spectrum, m / e 202 (molecular ion).5 II(0,4)b was synthesized from II(0,4)a with sulfur trioxide-pyridine5J5 in dry dioxane and pyridine. After neutralization of the reaction mixture with NaOH and evaporation of the solvents a yellow-orange solid was obtained. The compound was washed several times with diethyl ether, dissolved in ethanol, (13) Lee, T. D.; Keana, J. F. J. Org. Chem. 1975, 40, 3145. (14) Waggoner, A. S.; Kingzett, T. J.; Rottschaefer, S.; Griffith, 0. H.; Keith, A. D. Chem. Phys. Lipids 1969, 3, 245. (15) Sisler, H.; Andrieth, L. F. Inorg. Synth. 1946, 2, 173.
782 Langmuir, Vol. 2, No. 6, 1986
C
--
-
-
Barelli and Eicke
F-
05?1
I
Figure 1. ESR spectra of SL compounds in the AOT/H,O/iCJ-Ilssystem at T = 293 K and Wo= [H,O]/[AOT] = 5.2, [AOT] = 4.8 X lo-' mol dm-3. [SL compound] = 1.5 X mol dm-3. SL compounds: a = Ia, b = II(0,4)a, c = II(0,4)b.
20
w0
LO
60
Figure 3. Ratios of the relative amplitudes of the central- to the high-field line, h o / k l (izO.03) against Wofor different SL compounds in the AOT/H20/i-C8H18system at T = 293 K. SL compounds: (m) Ia; (0)Ib; ( 0 )II(0,4)a;(A)II(0,4)b. (Above W , = 15 the splitting of II(0,4)aprevents one from an unambiguous determination of hO/kl.)
Results and Discussion The ESR spectra line shapes of the SL compounds solubilized in the AOT/H20/i-C8H18mixtures depend on the chemical nature of the SL solubilizates. Figure 1 shows the ESR spectra of different SL compounds solubilized in the AOT/H20/i-C8Hl8 system a t room temperature and a t Wo= ([H,O)/[AOT]) = 5.2 (mol/mol). A comparison of the three spectra shows an obvious increase of the anisotropy from spectrum a to c. Since the hyperfine coupling constant aN of the SL compounds depends on the polarity of the immediate environment of the radical group, an analysis of the aNvalues of the spectra (Figure 1)provides information on the location of the SL compounds within the aggregates. In general, the difference of the aNvalues of nitroxide radicals between a polar and a nonpolar environment is typically 0.1-0.2 mT. Figure 2 shows a plot of the aN values from Figure 1 as a function of Wo. The hyperfine coupling constants of
water-insoluble compounds like Ia and Ib show a weak Wo dependence and the magnitude of the aN values in the whole W , range indicates that the radical group of these molecules is located (on the average) in a nonpolar environment. On the other hand, hydrophilic compounds like II(0,4)a and II(0,4)b show strongly Wo-dependent aN values. These spin-labels tend to move toward a polar environment with increasing amounts of added water. Isooctane-insoluble probes like II(0,4)b will approach the aNvalues found in pure water (see Table I). An analysis of the location of the different SL compounds in AOT aggregates should not be based exclusively on the hyperfine coupling constant. In fact, the anisotropy of the spectra of solubilized probes depends not only on the particular SL compound but also on Wo. Figure 3 displays the Wo dependence of the ratio h0/h-, (peak height ratio of the central- to the high-field line) of the systems from Figure 2. This ratio can be related to the rotational motion of the radicals if a possible exchange of the SL compounds between different environments is not taken into account;17 see also ref 5 . An increase of the ho/h-lratios indicates a decrease of the rotational motion of the radical. In the limits within which this relation is valid, a decrease of the tumbling rate of the radical produces an unequal broadening of the hyperfine lines, as that observed in Figure lb,c. Up to about W,, = 6, the water in the aggregated system is bound in the primary hydratation shell of AOT; hence, amphiphilic compounds like II(0,4)a and II(0,4)bwill be advantageously located in interstitial positions between the AOT molecules in the amphiphilic monolayer. As Wo increases from 0 to about 6, the SL compounds permeate through the monolayer toward the water core. This is seen from the increase of the ho/h-l and aNvalues in this W o range (see Figures 2 and 3). Above Wo= 6, a hydrocarbon-insoluble compound like II(0,4)b penetrates into the water core which gradually becomes less rigid with increasing amounts of solubilized water. Simultaneously, a decrease of the ho/h-lratio and increase of aNare observed. SL compounds like II(0,4)a, which are sufficiently soluble in hydrocarbon and water, show only intermediate h o / h _ and l a Nvalues up to about
(16) Barber, M.; Bordoli, R. S.; Elliot, G. J.;Sedwick, R. D.; Tyler, N. Anal. Chem. 1982,54, 645A.
Acad. Sei. U.S.A. 1967, 57. 1198.
Figure 2. Hyperfine coupling constants, (*0.004, mT), against Wofor SL compounds in the AOT H20/i-C8H18 system at T = 293 K, [AOT] = 4.8 X lo-' mol dm- , [SL compound] = 1.5 X mol dm-3. SL compounds: (m)Ia; (0) Ib; ( 0 )II(0,4)a;(A)(0,4)b.
/
and filtered. Evaporation of the solvents under reduced pressure gave II(0,4)b. Mass spectrum, m / e [fast atomic bombardment (FAB)]? 281 (base peak, molecular ion). Anal. Found: C, 39.27; H, 6.17; N, 4.56; S,10.66. Calcd for CloH1906N-NaS:C, 39, 47; H, 6.29; S, 10.54. The g factors and uN values of II(0,4)a and II(0,4)b are given in Table I.
(17) Waggoner, A. S.; Griffith, O.,H.; Christensen, C. R. Proc. Natl.
Langmuir, Vol. 2, No. 6, 1986 783
Exchange Kinetics of S u r f a c t a n t s i n Microemulsions
- - - -H*0
(0
- - - - -- -
20 wo 30
io
so
Figure 5. ESR spectrum of II(0,4)a in the AOT/H20/i-C8Hi8 system at W, = 60.3 and T = 303 K; [AOT] = 4.8 X lo-’ mol dm- ; mol dm-3. [II(0,4)a] = 1.5 X
Figure 4. Activation energy, E, (kJ mol-’) against W, for II(0,4)b in the A O T / H 2 0 / X 8 H 1 8 system. Broken line: E, value for II(0,4)b in aqueous solution: [AOT] = 4.8 mol dm-3. [II(0,4)b] = 1.46 X
X
lo-’ mol dm-3,
Wo= 12. This suggests that we observe only mean values of h,/ h-l and of aNfor different environments. The curves of II(0,4)a in Figures 2 and 3 show a characteristic kink at about Wo = 12. Previous experiencelg allows us to suppose that the kink indicates the onset of the penetration of the SL compound into the water core. This coincides with the first appearance of so-called free water in the water For Wovalues well above 20, the ho/h-l ratios of solubilized II(0,4)b approach a constant value which is larger than that found in pure aqueous solution (1.17 f 0.03). This apparent restricted motion of the SL compound in the water core of the nanophase has to be ascribed to the structure of the core since the radical concentration is very small. The idea of a viscous water shell surrounding the water core of the AOT aggregate has already been proposed.20*21Additional results in support of this concept are obtained from the Wo-dependent activation energies of the rotational correlation times of solubilized II(0,4)b (Figure 4). These data are derived from temperature-dependent rotational correlation times at constant Wo. (For a detailed discussion of this evaluation, see ref 5 and 17.) In Figure 4 the activation energy passes a maximum at about Wo = 6, i.e., a t a value which coincides with the number of HzO molecules in the primary hydration layer of the sodium ion. Already a slight increase of Wo makes the activation energy drop steeply, while above Wo= 25 it approaches a constant value of 15.7 kJ mol-’ which is still significantly larger than the corresponding value in pure water (10.5 k J mol-l). SL compounds like Ia and Ib display ho/h-lvalues which are very close to those observed for nitroxide radicals tumbling isotropically in nonviscous liquids.6 A very slight increase of the ho/h-land aN values is, however, observed up to about Wo = 15 (see Figures 2 and 3). This small effect could be tentatively related to the @crease of the area covered by one surfactant molecule (fAOT) with the increase of the amount of added water,22which facilitates the penetration of the SL compound toward the water core. Figure 5 shows an example of the spectral splitting observed in solubilization experiments with II(0,4)a. Since the uNvalues and g factors depend on the environmental p o l a r i t i e ~ such , ~ ~ spectra can only be observed if a slow exchange (i.e., between lo4 and 10“ s) occurs between two sites of sufficiently different p o l a r i t i e ~ . ~Accordingly, ~.~~ (18) Wong, M.; Thomas, J. K.; Nowak, T. J. Am. Chem. SOC.1977,99,
20
40
wo
60
Figure 6. Line width of the right high-field line, 8-1(H20) (*0.004, mT), against W, for II(0,4)a in the AOT/H20/i-C8H18system a t different temperatures (*l,K): ( X ) 283; (A)293; (0)303; (0) 313. (Scattering of points is due to reading out the experimentally recorded ESR spectra.) the right-hand high-field line is assigned to a less polar environment (aggregate monolayer). An interesting example of the line-width behavior as a function of Wo is shown in Figure 6 where the line width of the right-hand high-field line is plotted against W,. The plot displays isotherms which pass through minima. In order to elucidate this unexpected line-width dependence on Wo,we refer to a kinetic treatment based upon a simple stoichiometric equilibrium, i.e., E M+R=MR (1) kL where M denotes the nanodroplet, R the II(0,4)a radical, and MR the radical nanodroplet “complex”, Le., the SL compound solubilized in the nanodroplet. Applying the mass action law to eq 1 yields
where s is the molar concentration ratio [MR]/[R]. Since sufficiently small amounts of reactants and products are assumed, the activity coefficients of all species are taken to be unity. The experimental determination of the rate constants and k can be accomplished by the exchange-modified Bloch equation, which reads for slow-exchange conditionsx
l / T z denotes the experimental line width, l / T z , ois the
47.70.
(19) Eicke, H.-F. Chimia 1982, 36, 241. (20) Hoffmann, H. Prog. Colloid Polymer Sci. 1978, 65, 140. (21) Zinsli, P. E. J. Phys. Chem. 1979, 83, 3223. (22) Eicke, H.-F.; Rehak, J. Helu. Chim. Acta 1976, 59, 2883. (23) Kawamura, T.; Mataunami, S.; Yozenawa, T. Bull. Chem. SOC. Jpn. 1976,40, 1111.
(24) Carrington, A.; MacLahlan, A. D. Introduction t o Magnetic Resonance; Harper & Row: New York, Tokio, 1969. Atherton, N. M. Electron Spin Resonance; Ellis Harwood: Chichester, England, 1973. Griffith, 0. H.; Waggoner, A. S. Acc. Chem. Res. 1969, 2, 17. (25) Griffith, 0. H.; Waggoner, A. S. Acc. Chem. Res. 1969, 2, 17.
784 Langmuir, Vol. 2, No. 6, 1986
Barelli and Eicke
Table 11. Experimental Line Width of the Right-Hand High-Field Line (b-,(a20)e/mT), Simulated Line Width of the Same Line (&l(H20)JmT), and the Difference between Them r(L - a_, )/mT1 for II(0.4)a” (6-10,0),
(7’ f 1 ) / K 293 298 303 313 318
f 0.0038)/mT
0.1088 0.0975 0.0938 0.0900 0.0863
6-1(H20),/
I
- a-18)/ mT 0.0022 0.0028 0.0034 0.0047 0.0062
(a-le
mT 0.1066 0.0947 0.0904
0.0853
o.oaoi
“ I n the AOT/H,0/I-C8H18 system at W , = 50.4 and different temperatures.
natural line width in the absence of any exchange (or at very low exchange), and 1/T2,ex is the line-width contribution of the exchange processes. T2,ex represents the mean lifetime during which II(0,4)a exists between exchanges. It is assumed that the proper exchange occurs instantaneously. The rate constants are related to 1/T25,ex, Le., (4) and
To evaluate the rate constants, l/T2,0must be known. This can be elucidated by simulating the experimental spectra with a polynomial expansion of the line width, d with respect to the nuclear spin quantum number of nitrogen, ml, i.e., 6,, = -= A + B m I + C m ?
(6)
TZ(m,)
assuming Lorentzian line shapes.24 This approach considers the anisotropy of the spectra as an effect only of the natural line width. Under slow exchange conditions, the exchange line width contribution can be assumed to be the same for all spectral lines.24 This method to obtain the natural line widths differs considerably from that reported by Yoshioka and KazamalO applied to data from similar exchange kinetics experiments. These authors calculated the natural line width of their solubilized radicals as the average value between the line width in pure hydrocarbon solution and in micellar solution ( Wo= 0), without taking into account the anisotropy and hence line-width dependence on Wo. As an example, Figure 7 shows the experimental (digitalized) and simulated spectra of II(0,4)a in the AOT/H20/i-C8H18system at Wo= 50.4 and T = 303 K. Table I1 exhibits the experimental and simulated line-width values of the right-hand high-field line of the same system of Figure 7 a t different temperatures. The line width that was attributed to the exchange process was calculated by subtracting the simulated (=natural) line width from the experimental one (in frequency units and with 6 = 2/3If2T2). As has to be expected, the exchange line width increases with increasing temperature due t o the increase of the exchange rate. Thus, the minima of the 6-1(H,o)- W , plot follows from an initially stronger decrease of the natural line width compared to the increase of the exchange line width with W,. At larger weighed-in amounts of water, however, the effect of the exchange line width prevails. This result is again in line with the already discussed Wo-dependent properties of the water core of the nanodroplets.
Figure 7. ESR spectrum of II(0,4)ain the AOT/H20/i-C8H18 system at Wo= 50.4 and at T = 303 K: (a) instrument digitalized experimental spectrum; (b) simulation of the experimental spectrum. The weak satellite lines (from 13C nuclei in natural abundance) are not simulated; (c) superposition of (a) and (b).
The analysis of the forward reaction (eq 1)indicates that the inverse of the mean life time of II(0,4)a in the water phase ( l/T2,ex)R depends on the nanodroplet concentration [MI since the reaction is of second order (see eq 4). [MI can be obtained from geometrical considerations assuming that (i) the mean area of the water/oil interface covered by one AOT molecule, fAOT, is W , and temperature independent above W , = 20 and the value is approximately 0.54 nm2,22(ii) the AOT monomer concentration is negligible compared to [MI, i.e., all of the AOT molecules are accommodated in the interface, and (iii) the nanodroplets are spheres of radii r, where the radii at Wo= 34.8,40.0, 45.2, 50.4, and 60.3 are 6.9, 7.7, 8.6, 9.5, and 1.1 nm, respe~tively‘~ (assuming the AOT hydrocarbon tail length is 1.1nm). The total interface AM created by the nanodroplets is given by AM =
4.rrr2nM= fAoTIAOT]
(7)
where n M is the number of nanodroplets. Since the amount of “labeled” nanodroplets MR is assumed to be negligible compared to the amount of “nonlabeled” na; nophases M, one obtains n M = [MI. The calculated k values (eq 4) are given in Table 111. The backward reaction constant k can be calculated from (1/T2,ex)Rand the molar ratio s (see eq 5). Since the intensities of the mI = -1 components are proportional to the concentrations of II(0,4)a in the water core [R] and in the interface [MR], the s values can be calculated from the ratio of the line widths of the t y o high-field lines in the simulated spectra. The values of k and s are giyen in_Table I11 as functions of Woand temperature. The k and k data are reported only for high W , values because the line widths and s values for low Wovalues could not be accurately determined. The ( ~ / T P , and ~ ~ )( 1R/ T 2 , e x ) M R expressions were fitted according to Arrhenius plots, i.e.,
in order to obtain the activation energies of the forward
Langmuir, Vol. 2, No. 6, 1986 785
Exchange Kinetics of Surfactants in Microemulsions s-l mo1-I dm3) and Backward Tabl? 111. Forward @/lo6 s-l) Rate Constants of Exchange of II(0,4)aa i / l O l O s-l W, (5" 1 ) / K kc/106 s-l mol-' dm3 S 283 4.12 293 4.55 34.8 303 5.13 313 5.61 323 7.14 2.98 0.163 283 6.26 8.22 4.00 0.167 293 4.40 0.177 40.0 303 8.53 10.45 5.58 0.183 313 11.01 6.13 0.191 323 2.44 0.124 293 5.47 3.48 0.136 303 7.12 4.50 0.163 45.2 313 7.67 5.62 0.183 323 8.54 2.80 1.43 0.117 293 3.28 1.86 0.130 298 3.57 2.23 0.143 50.4 303 3.13 0.163 313 4.40 5.29 4.09 0.177 318 0.95 0.099 293 1.62 1.36 0.105 298 2.08 1.81 0.111 60.3 303 2.75 3.58 2.50 0.117 308
120
*
OIn the AOT/H20/i-D8H18system as a function of Woand temperature (T l / K ) ; s = [MR]/[R] (see text).
*
Table IV. Activation-Energies for the Forward (k',/kJ mol-') and Backward (EJkJ mol-') Reactions according to Eq 8 and 9 W, l?,/kJ mol-' r l?,/kJ mol-' P 34.8 9.9 0.975 0.945 0.975 10.5 40.0 13.6 0.934 0.994 11.2 45.2 21.9 18.5 0.986 31.0 0.998 50.4 0.999 0.999 39.9 60.3 47.6 O r
= least-squares correlation coefficient.
backward 2, reaction5 cf eq 1 (see Table IV). These data together with the k, k, and s values of Table I11 indicate an increasing tendency of II(0,4)a to stay in the polar water core with increasing W,. Since in the Wo range investigated the area per AOT molecule at the interface fAoT is constant,22the sole factor which can determine this behavior is the size of the aqueous nanodroplets. In view of the amphiphilic character of th_e spin prgbe, it appears interesting to plot AgHI(Wo)= Ea(Wo) - Ea(Wo)(see Table IV) as a function of the water core size which is determined by Wo = [H,O]/[AOT] (Figure 8). AgHI represents the transition enthalpy of the label between the two solubilization sites which is seen to pass through a pronounced maximum with increasing Wo. Independently of the AzHI(Wo)data, we have access to the entropic part of the free energy which governs the exchange process. Table I11 lists s values as functions of temperature and Wo,where s (see eq 2) denotes the molar concentration ratio [MR]/ [R] with MR the radical-nanodroplet complex (i.e., the spin label solubilized in the monolayer) and R the II(0,4)a radical. This s ratio is a measure of the equilibrium distribution of the label between the amphiphilic monolayer and the water core of the nanodroplet. Following this idea, the transition entropy may be obtained from the coexistence condition with respect to the radical in the two "states", i.e.,
Za and
dprm= dhIw
(10)
Both sides of eq 10 can be considered a total differential of the chemical potential as a function of temperature and
L
40
wo
50
60
Figure 8. Transition enthalpy AzHr(Wo)and entropy AgSr(Wo) of II(0,4)a against W,, = [H,O]/[AOT].
mole fractions, XImand XIw,respectively; thus eq 10 can be rewritten:
Since the fluidity of the water/surfactant (AOT) interface is strongly dependent on the amount of solubilized water and on temperature below about Wo = 20,22we are only interested in the situation above this Wovalue. We define the transition entropy as AgS, = Srw - S," and obtain from eq 11, according,
AZS, = -R
d In (Xrm/XIw) d In
T
(12)
AZSr as determined from eq 1 2 has also been plotted for comparison in Figure 8 against Wo. Like the transition enthalpy, AgS, also passes a maximum with growing core size of the nanophases. A consequence of this behavior is the relatively feeble minimum of the equilibrium constant around W , = 50 which can easily be inferred from Table 111. This Figure demonstrates that transition entropy and enthalpy are controlled by two counteracting effects. Keeping in mind the pronounced amphiphilicity of this particular label and its equal solubility in water and oil, we find steeply rising plots of A:&( Wo)and AEHr(Wo) with increasing amounts of solubilized water. This is understandable in view of the increasingly favorable tendency of the label to exchange between monolayer and water core of the growing nanophase. The exchange, however, is hampered (AgHI> 0) by the ordered surfactant layer and its viscous double layer20 due to the hydrated sulfosuccinate groups and sodium ions of AOT. This latter effect will decrease with larger water cores, probably due to an easier accommodation of the hydrophilic groups of the label between the AOT molecules, hence giving rise to a fuller hydration of the head groups of the radical. Also, larger water cores imply more "free" (bulk) water and thus an improved antagonism between water and oil domains. This starts t b confine the spin-label to the correspondingly better defined oil/ water interface and results in a reduction of AgS,. In conclusion we have shown that the electron spin resonance technique together with suitable spin labels can provide valuable information on dynamic (surfactant exchange kinetics) and static (solubilization sizes) properties of water-in-oil microemulsions. The physical picture of the W/O microemulsion which evolves from this investigation agrees very satisfactorily with the general views on
Langmuir 1986, 2, 786-788
786
such liquid-liquid dispersions which have been developed over the past decade.
Acknowledgment. We are greatly indebted to Prof. F. Gerson for making his ESR equipment available to us. A.B. thanks the "Centro de Formacidn y Adiestramiento
Petrolero y Petroquimico (CEPET)", Caracas, for a fellowship. Finally, we acknowledge financial support from the Swiss National Science Foundation. Registry No. Ia, 104393-98-8;Ib, 104393-99-9;IIa, 104394-00-5; IIb, 104394-01-6;AOT, 577-11-7;isooctane, 540-84-1.
Study of Thermal Behavior of Langmuir-Blodgett Films with an Emission Probe Tadahiro Murakata, Tokuji Miyashita, and Minoru Matsudah Chemical Research Institute of Nonaqueous Solutions, Tohoku Uniuersity, Katahira 2-1 -1, Sendai, 980 Japan Received May 23, 1986. I n Final Form: July 25, 1986 Fluorescence emission spectra of 1-pyrenedodecanoic acid (PDA) dispersed into barium stearate Langmuir-Blodgett (LB) films were measured at various PDA mole fractions. The emission intensity ratios of excimer to monomer (Ie/Im) in the LB films were considerably higher due to good chromophore orientations than those in the films prepared by casting. The temperature dependency of Ie/Imwas examined, and a sudden decrease was observed at a critical temperature related to the destruction of the layer structure in LB films. It was found that the emission probe method is a convenient and sensitive method to study the thermal behavior of LB films.
Introduction
It is known that molecular assemblies in which chromophores are regularly arranged show some characteristic optical properties1-12which are often different from those observed in solutions or in amorphous assemblies. The Langmuir-Blodgett (LB) technique provides a thin film in which the chromophore orientation is well-defined, and this film can be regarded as a quasi-crystalline state of lamellar structure. Pyrene excimer formation is a fundamental photophysical process. The process in fluid media is governed by the diffusion of pyrene13 and has often been used for the measurement of the microviscosity of the fluid media. However, excimer formation in rigid media has not been studied as much as in fluid media.14-16 In the LB films, excimer formation is expected to represent the packing and orientation of molecules. In this paper, we study the temperature effect on the excimer formation of pyrenyl groups of PDA dispersed into barium stearate LB (1) Vaidyanathan, S.;Patterson, L. K.; Mobius, D.; Gruniger, H. R. J . Phys. Chem. 1985,89,491. (2) Mobius, D. Ber. Bunsenges. Phys. Chem. 1978, 82, 848. (3) Mobius, D. Acc. Chem. Res. 1981, 14, 63. (4) Heesemann, J. J . Am. Chem. Soc. 1980, 102, 2167. (5) Heesemann, J. J . Am. Chem. Soc. 1980, 102, 2176. (6) OBrien, D. F. Photogr. Sci. Eng. 1974, 18, 16. (7) Nakahara, H.; Fukuda, K.; Kato, T. J. Colloid Interface Sci. 1976, 54, 430. (8)Nakahara, H.; Fukuda, K. J. Colloid Interface Sci. 1979, 69, 24. (9) Nakahara, H.; Fukuda, K. J . Colloid Interface Sci. 1981, 83, 401. (10) Nakahara, H.; Fukuda, K. J. Colloid Interface Sci. 1979,68,555. (11) Valenty, S . J. J. Colloid Interface Sci. 1979, 68, 486. (12) Fukui, T.;Saito, M.; Sugi, M.; Iizima, S . Thin Solid Films 1983, 109, 247. (13) Quina, F. H.; Whitten, D. G. J. Am. Chem. SOC.1977, 99, 877. (14) Birks, J. B.; Lumb, M. D.; Munro, I. H. Proc. Roy. SOC.London, A 1964,280, 289. (15) Doller, E.; Forster, T. 2. Phys. Chem. N . F. 1962, 34, 132. (16) Sisido, M.; Takeuchi, K.; Imanishi, Y. J . Phys. Chem. 1984, 88, 2893.
0743-7463/86/2402-0786$01.50/0
films and find that PDA is a good emission probe to investigate thermal properties of LB films.
Experimental Section 1-Pyrenedodecanoicacid (PDA) (Molecular Probes, Inc.) and stearic acid (Wako Chemical Co.) were used without further purification. Chloroform was of fluorescence spectroscopygrade. The surface pressure-area isotherms were measured with an automatically working Langmuir trough (Kyowa Kaimen Kagaku HBM-AP, using a Wilhelmy film balance). Fluorescence intensities and X-ray diffractions were measured by a Shimadzu RF 503A spectrofluorophotometer and a Shimadzu VD-1 X-ray powder diffractometer, respectively. Quartz slides on which monolayers were deposited were cleaned in boiling H2S04/HN03 (2/1) solution and made hydrophobic by treating them with a 2 % solution of dimethyldichlorosilane in chloroform. Prior to the deposition of PDAjbarium stearate monolayers, a few barium stearate layers were deposited onto the quartz slide as base layers to gain a homogeneous hydrophobic surface. The cast films on quartz slides were prepared by evaporation of solvent from a stock solution used for monolayer preparation.
Results and Discussion A chloroform solution of PDA and stearic acid was spread on the surface of water which contained 3 X M BaC1, and 5 x M NaHC03. The surface pressure-area ( F A ) isotherms for the monolayers were measured at 17 "C (Figure l). Stable mixed monolayers were obtained in the presence of barium salt ((a)-(d) in Figure 1). The mixed monolayers were compressed to a surface pressure of 20 dyn/cm and transferred onto the pretreated quartz slide in both down and up trips by the LB method. Emission and excitation spectra of PDA in the obtained LB films (Y type) were measured under N2 atmosphere (Figure 2). The former spectra were obtained by exciting PDA a t 280 nm and the latter spectra by monitoring the emission of PDA a t 375 and 460 nm. The emission bands with peaks a t 375, 395, and 425 nm are assigned to mo0 1986 American Chemical Society
