Langmuir 1995,11, 3007-3012
3007
Synthesis of Surfactants by Micellar Autocatalysis: NJV-DimethyldodecylamineN-Oxide Paul R. Kust and James F. Rathman" Department of Chemical Engineering, The Ohio State University, 140 West 19th Avenue, Columbus, Ohio 43210-1110 Received April 12, 1995. In Final Form: June 7, 1995@ The synthesis ofNJV-dimethyldodecylamineN-oxidein aqueous solutions by micellar autocatalysis was investigated. Micellarautocatalysisis a novel variation of conventionalmicellar catalysisin which surfactant micelles catalyze the reaction by which the surfactant itself is synthesized. The lipophilic reactant, dimethyldodecylamine,was initially solubilizedin micellar solutionsof the amine oxide surfactant, resulting in substantially higher reaction rates. Amine conversions of 90-100% were obtained within 2 h at 70 "C. "he effects of reactant concentrations, temperature, and initial surfactant concentration were studied. For systems with no surfactant at time zero, the system was initially an emulsion and reaction rates were low. A sharp increase in the rate was observed when enough surfactant had been produced to form micelles. Activation energy calculationsindicate that enhancements of the rate were due primarily to the localizedconcentration ofreactants in the micelle. A simplepseudophase model was used to model reactions under pseudo-first-order reaction conditions.
Introduction Background. Micellar autocatalysis is the use of surfactant micelles to catalyze a reaction by which the surfactant itself is synthesized. Investigation of "selfreplicating" micelles as models for cell reproduction and studies of prebiotic chemistry has been however, research into the application of micellar autocatalysis in synthetic schemes has not previously been published. The synthesisof surfactants in aqueous media via micellar autocatalysis is important for a number of reasons. First, micellar autocatalysis provides a method for synthesizing surfactants without employing volatile organic solvents in the reaction medium, providing potential economic and environmental benefits. Second, studies of micellar autocatalysis can refine and extend the understanding of other types of reactions in aqueous surfactant solutions. Despite a wealth of fundamental research in the area of conventional micellar c a t a l ~ s i s , ~ - ~ very few commercial synthesis processes currently employ this technology, in part due to an incomplete understanding of how to quantitatively model parameters such as high reactant concentrations and product effects. In order for this technology to become a viable industrial alternative, a number of key issues must be addressed. One of the most significant factors that affect the kinetics of reactions in aqueous media is the solubility or miscibility of each reactant in the reaction medium. The production of many chemicals can be made more efficient and economical by improving methods for combining waterinsoluble and water-soluble reactants. For example, the production of surfactants, compounds that contain hydrophobic and hydrophilic moieties within a single molecule, generally involves reaction of components which are not mutually soluble in a single solvent. In a typical
* To whom correspondence should be addressed. @
Abstract Dubfished inAdvance ACSAbstracts, August 1,1995.
(1)Bachmin, P. A.; Walde, P.; Luisi, P. L.; Lang, J. j.Am..Chem. SOC. ..
.
1990.112.8200. ---, -...
~-
(2) Bachmann, P. A.; Walde, P.; Luisi, P. L.; Lang, J. J.Am. Chem. Soc. 1991,113,8204.
(3)Bachmann, P. A.; Luisi, P. L.; Lang, J. Nature 1992, 357,57. (4)Menger, F. M.; Portnoy, C. E. J.Am. Chem. Soc. 1967,89,4698. ( 5 ) Fendler, J . H.; Fendler, E. J . Catalysis in Micellar and Macromolecular Systems; Academic: New York, 1975. (6)Bunton, C. A. In Kinetics and Catalysis in Microheterogeneous Systems; Gratzel, M., Kalyanasundaram,K., Eds.; Surfactant Science Series Vol. 38;Marcel Dekker: New York, 1991;p 13.
commercial process, a water-insoluble reactant is first dissolved in a n organic solvent (e.g., methanol, acetone) and then mixed with a water phase containing the watersoluble reactant. After completion of the reaction, the organic solvent must be separatedfrom the product stream and either purified for recycle or discarded, thereby adding to the cost of the process and creating potential environmental hazards. A novel approach, described in this paper, is to use a surfactant to catalyze the reaction by which that same surfactant is produced. The autocatalytic scheme does not require a n organic cosolvent, and, in comparison to emulsion systems, can provide markedly increased reaction rates since the reaction solution is single-phase. Although conventional micellar catalysis has long been studied as a synthetic technique,'r8 the impetus for much of the previous research has been focused on using micelles as models for enzyme^.^ Indeed, the size of micelles is similar to that of many enzymes, and rate enhancements are comparable to those observedin enzymatic catalyses.1° Micelles are believed to affect reactions in several ways. The outer-core region of the micelle, commonly referred to as the palisade layer, may provide a medium of intermediate polarity that effects the energetics of transition-state f0rmation.l' In many systems, the primary influence of micelles is to concentrate all reactants in or near the micelles. Micellar catalysis exploits the ability of surfactant aggregates to solubilize significant amounts of reactants that are virtually water-insoluble.11J2 When ionic surfactants are employed, polar or ionic reactants that are freely soluble in water may also be concentrated near the micelles by electrostatic or dipole interactions. The degree to which reaction rates are enhanced depends on a number of factors, including the location of the solubilized reactant in the micelle: nonpolar compounds partition into the micelle core while more polar (7)Livneh, M.; Sutter, J. K.; Sukenik, C. N. J.Org. Chem. 1987,52, 5039. (8)Jursic, G. B. Tetrahedron 1988, 44, 6677. (9)Wennerstrom, H.; Lindman, B. Phys. Rep. 1979,52, 1. (10)Moroi, Y. Micelles: Theoretical and Applied Aspects; Plenum: New York, 1992. (11)Myers, D. Surfactant Science and Technology, 2nd ed.; VCH: New York, 1992. (12)Rosen, M. J. Surfactants and Interfacial Phenomena, 2nd ed.; John Wiley and Sons: New York, 1989.
0743-7463/95/2411-3007$09.0Q/O0 1995 American Chemical Society
3008 Langmuir, Vol. 11, No. 8, 1995
compounds (e.g., long-chain alcohols) are found closer to the micelle-water interface. The extent of solubilization, ionic charge of the micelle, and the size and shape of the micelle are also important factors. Micellar Autocatalysis. One area that has not been examined is the case in which the reaction product is itself a surfactant. More specifically, the use of micellar catalysis in the synthesis of surfactants and the effects of the product surfactant on the subsequent kinetics have not been studied. Since the formation of a surfactant results in an increase in the total concentration of surfactant in micellar form and since the reaction is catalyzed by micelles, these reactions are examples of micellar autocatalysis. Given the potential utility of this technology from both fundamental and practical points of view, this is a n area of special interest. The primary objective of this research was to develop a novel method for the production of synthetic surfactants which does not employ organic solvents. The second objective, based on studies of these same surfactant synthesis processes, was to initiate a systematic investigation of the effects of reaction products on micellar catalysis. Previous research has not determined the quantitative effects of reactants and products on the size and shape of the micelles; such effects must be understood in order to develop useful applications of this technology.
Experimental Section Materials. NJV-Dimethyldodecylamine N-oxide (DDAO) (CAS 1643-20-5) was from Fluka Chemika. Two grades were used in this study: purum (supplied as 30 wt % aqueous solution) and Biochemika (supplied as 98% active solid). NJV-Dimethyldodecylamine (DDA)(CAS 112-18-5)was from Fluka Chemika and Aldrich. Hydrogen peroxide (30% aqueous solution, CAS 7722-84-1) was from Mallinkrodt. Ethanol (CAS 64-17-5) was from Aaper Alcohol and Chemical Co. The solvents used for HPLC analysis were HPLC-grade water, methanol, and acetonitrile from Fisher. Aqueous solutions were prepared using deionizeddistilled water from a Barnstead MP-1 still. Titrations to determine the exact composition of DDAO received from Fluka Chemika were performed using a Corning 350 pWion analyzer. For aqueous titrations, a Cole-Parmer surfactant electrode (Catalog No. 27502-45) was used. The titrant used was 0.01 M sodium dodecyl sulfate in water. A carefully measured sample was placed in a beaker with 50 mL of water. The solution was acidified with 4-5 mL ofO.O1M HCl, and 1mL of Triton X-100 was added in accordance with the electrode instructions to preserve electrode life. For nonaqueous titrations, a Corning high-performance combination pH electrode (Catalog No. 476390) was used. The titrant used was a 0.1 M perchloric acid solution in dioxane, and titrant and flow rate were standardized against a 0.1 M solution of sodium hydroxide in water. A carefully measured sample was placed in a beaker with 50-100 mL of acetone or acetonitrile and titrated using methods described in the next section. The HPLC pump was a P-2000 binary gradient pump from Thermo Separation Products. Samples were injected via a Rheodyne 7025 injection valve, and separation was achieved on a 150- x 4.6-mm reversed-phase column with a 10- x 4.6" guard column, both columns packed with 5-pm particles of polymer-& from YMC Corp. The detector used to determine DDAO concentrations was a n Evaporative Analyzer Model 750/ 14 from Applied Chromatography Systems. The detector signal was processed and integrated on a Spectra Physics SP-4400 Chromjet integrator. Methods. Nonaqueous and aqueous titrations were employed to characterize DDAO received from Fluka Chemika. The characterization was necessary to accurately determine the DDAO content of the supplied material and to determine if these contained any residual tertiary amine. In the nonaqueous scheme, solutions of DDAO in acetonitrile were titrated with 0.10 M perchloric acid in dioxane. As described by Benzinger et
Kust and Rathman Table 1. Binary Gradient Pump Solvent Program for HPLCa time, min
solvent A, vol 7%
solvent B, vol 7%
0 2 3 8 12 15
50 50 90 90 50 50
50 50 10 10 50 50
aSolvent A 100% methanol. Solvent B: 15/15/70 (vol %) methanoYacetonitrile/water. al.,13 these titrations yield two endpoints, the first endpoint corresponding to one-half of the total amine oxide present and the second to half the total amine oxide plus any amine that may be present. The results showed that both the purum and Biochemika grades of DDAO contained at least 99 wt % amine oxide and a maximum of 1wt % residual amine. HPLC analysis showed that the primary difference between the purum and Biochemika grades was that the amine oxide in Biochemika is essentially 100% DDAO, while the purum grade contains ca. 60 wt % DDAO and the rest being a n amine oxide with a longer alkyl chain, most likely dimethyltetradecylamine oxide. Aqueous titrations were performed to accurately confirm the DDAO concentration in prepared stock solutions. This method takes advantage of the cationic-nonionic nature of amine oxide surfactants, which are nonionic in neutral or basic solution but are protonated in acidic solution to form cationic surfactants.14 Using a surfactant-selective electrode, cationic amine oxide can be titrated in aqueous media using a solution of anionic surfactant as titrant, eliminating many of the problems associated with the handling and use of nonaqueous solvents. Aqueous sodium dodecyl sulfate (SDS) (0.01 M) was used as the titrant and was standardized using 0.05 M Hyamine 1622 solution. DDAO solutions were acidified with HC1, and 1mL of Triton X-100 was added as a n ionic strength adjuster to optimize the electrode performance. In order to determine the range of tertiary amine concentrations to be used in reaction experiments, the solubility of DDA in aqueous DDAO solutions was estimated by dissolving a measured amount of amine into a concentrated DDAO solution and then adding water until phase separation occurred. Between each incremental addition of water, solutions were agitated continuously and maintained at 35 "C. Synthesis reactions were carried out in 125-mL Erlenmeyer flasks. During reaction, the flasks were sealed and continuously stirred while immersed in a constant-temperature water bath. After allowing the initial solution to come to thermal equilibrium, the desired amount of hydrogen peroxide was added-this was considered to be time zero for the reaction. As the reaction proceeded, small (0.1-0.5-g) samples were withdrawn by pipette and diluted with 0.1% sodium hydroxide at room temperature to 15-20 times the volume of the sample. Dilution, reduced temperature, and addition of sodium hydroxide to decompose the remaining peroxide all act to stop the reaction. Samples were then analyzed by HPLC to determine the amine oxide concentration. The HPLC pump was connected to two solvent reservoirs, one containing pure methanol (solvent A) and the other containing a mixture of 15% methanol, 15% acetonitrile, and 70% water by volume (solvent B). Solvents A and B were mixed by the gradient pump according to the program shown in Table 1. The total solvent flow rate was constant at 0.8 m u m i n . The evaporative column temperature of the detector was 55 "C, and dry air at inlet pressure 25 psig was used for sample nebulization.
Results and Discussion pH Effects on DDA, DDAO, and Hz02 in Aqueous Solutions. Amine oxide surfactants are produced by reaction of a tertiary amine with hydrogen peroxide. For DDAO, the reaction is (13)Benzinger, N. N.; Galpern, G.D.; Ivanova, N. G.; Semechkina, G. A. Z h . Analit. Khim.1968,23,1538. (14)Rathman, J. F.; Christian, S. D. Langmuir 1990, 6, 391.
Langmuir, Vol. 11, No. 8, 1995 3009
Synthesis of Surfactants by Micellar Autocatalysis
This reaction system is complicated somewhat in that all components exhibit pH-dependent dissociation eq~ilibria:'~
H 2 0z z H+ + OH-
pK,
H202= H+ + HOODDAH' DDAOH'
* DDA + H+ t DDAO
+ H+
M
(2)
14
6
8
O 9
6 4
2 0
Q
2
4
6
60 62 64 66 68 70 72
a i 0
wt% DDAO in water
pKa x 11.6 pKa x 11.0 pKa M 5.0
(3)
(4) (5)
The only compounds present in the reaction systems studied here were H20,H202,DDA, and DDAO; no attempt was made to control the solution pH in order to avoid the addition of any component not desired in the final product. Solutions were observed to be slightly basic (pH 8-91, the exact value depending on the initial composition and varying only slightly over the course of a reaction. Within the observed pH range, the formation of DDAOH+ and the dissociation of HzOz can be neglected; however, the formation of DDAH+must be considered since this species is freely water-soluble and in fact may behave as a cationic surfactant. Essentially all tertiary amine in aqueous solution a t pH < 9 is present as DDAH+. The pH dependence of the total amount of tertiary amine that can be dissolved in water can be approximated by
+
+ 10PKa-PH) (6)
[DDAH+laq [DDAlaq= [DDAlaq(1
where [DDA],, is extremely low at all pH levels. From eq 6, it is apparent that due to the low solubility of DDA, appreciable concentrations of tertiary amine, primarily in the form DDAH+, can be achieved only a t pH levels well below the pK,. Solubilization of DDA in DDAO Solutions. A comparison of amine solubility in aqueous solutions of DDAO and methanol is illustrated in Figure 1. Although nearly insoluble in water, the solubility of DDA increases significantly in micellar solutions of DDAO. The purpose of determining DDA solubility was not to construct a detailed phase diagram but rather to identify the concentration range in which single-phase solutions can be prepared for reaction experiments. The solubilities determined from this technique are therefore only approximate. Over a wide range of surfactant concentrations, micellar solutions of DDAO are capable of solubilizing DDA up to a 2:l (mol/mol) DDMDDAO ratio. Solubilities in methanovwater solutions are shown for comparison since methanol is employed in some conventional synthesis processes for amine solubilization.16 Simple First-Order Model. Traditionally, micellar catalysis is discussed in terms of the pseudophase model, in which the micelles and surrounding bulk solution are treated thermodynamically as separate phases. Micellar solubilization is then viewed as an equilibrium partitioning of solubilized component between the two pseudophases:
(7)
A+M*MA
where A is the water-insoluble reactant, M is the surfactant micelle, and MA is the micelle containing solu~
~~
~~
(15)Maisonneuve, B. In Kirk-Othmer Encyclopedia of Chemical Technology, 4th ed.; John Wiley and Sons: New York, 1992; Vol. 2, p 357.
wt% MeOH in water
Figure 1. Solubility at 35 "C of DDA in (a, left) DDAO/H20 solutions and (b, right) methanol/HzO solutions. bilized reactant. The chemical reaction of interest is assumed to proceed irreversibly in both phases; for a simple first-order reaction A P,
-
k'w
A-P
MAk'"M+P
(9)
where K', and K', are the fist-order reaction rate constants in the water and micellar phases, respectively. The derivation and evaluation of various models based on the pseudophase approach are summarized by Bunton.6 Current models for micellar catalysis do not adequately address systems in which the lipophilic reactant is present at high concentrations, i.e., concentrations on the same order ofmagnitude as the surfactant. This is a n important area of investigation because use of micellar catalysis in synthesis processes is realistic only if high reactant concentratians can be employed. The overall and micellar concentrations of the solubilized reactant can be defined in terms of the total and micellar volumes: ntot
[AItot = TI
" tot
(10) where ntot and nmicrepresent the amount of total and micelle-bound (solubilized) amine, respectively. The micelle volume (Vmic)is dependent on the amount of surfactant in micellar form: vmic =
(
[SIto, - cmc @M
)vtot
(11)
where @M is the molar density of the micelle. If a reactant is virtually insoluble in water, then the number of moles of reactant solubilized in micelles is essentially equal to the total number of moles of reactant present (nmic ntOt). The micellar concentration of reactant can then be calculated:
(12)
Also, if all the lipophilic reactant is solubilized in micelles, the reaction rate is determined only by the reaction taking place in the micelles:
3010 Langmuir, Vol. 11, No. 8, 1995 0.7
3cl
0.6
3
0.5
E
Kust and Rathman
1 I
I
3
i
s
0.7
9
0.4
' [DDAO],
0.3
* 0
100
50
150
200
250
300
Figure 2. Conversion of DDA vs time at 70 "C for reaction with no DDAO initiallypresent. Initial concentrations: [HzOZIO = 1.0 M, [DDAlo = 1.0 M, [DDAOlo = 0.
The pseudo-first-order rate constant, k,, is most easily related to k', from the pseudophase model since it is assumed that the reaction takes place only in the micelles. Because the reactant in this case is reacting to form the solubilizing surfactant, a mass balance can be written as [Altot+ [Slht = [AI0 [Slo,where the subscript 0 indicates the initial concentrations. The rate equation becomes
+
-d[S1tot- K&[Al0 + [SI, - [SI,,,) dt
([SI,,, - ) @M
cmc
(14)
and now the rate equation is in terms of one variable ([SItot) and other parameters ([Alo, [Slo, @M, kv, cmc) that can be assumed constant for a given reaction system. The integrated rate expression is
The surfactant monomer concentration is expected to vary somewhat with time as the amount of solubilized reactant decreases as the amount of micellar surfactant increases. As discussed in the following section, this variation is not significant in this study since [SI0 and [SIhtx=- cmc. Reactions with No Surfactant Present at Time Zero. Initial experiments were designed to be similar to methods described in the patent 1iteraturel6-l8 for the synthesis of amine oxide surfactants. At time zero, solutions contained only tertiary amine, hydrogen peroxide, and water. Neither organic solvent nor surfactant was used to solubilize the tertiary amine, so the system was two-phase, and vigorous mechanical agitation was provided to emulsify the amine. As expected, this emulsion was initially very unstable and could only be maintained by continuous agitation. No catalysts or other additives were used. Figure 2 displays the progress of a typical reaction using stoichiometric amounts of the reactants. The initial reaction rate was extremely low and a sharp increase was observed between 110 and 120 min. During the first 100 min, formation of surfactant (DDAO) resulted in fairly efficient emulsification of the tertiary amine; however, the reaction rate remained low during this period. Despite the fact that emulsification of the amine by the surfactant (16)Shehad, N.S.;Dussack, L. L.; van Hemelrijk, D. US.Patent 5,254,765, 1991. (17)Smith,K. R.;Borland, J.E.; Sauer, J. D. U.S. Patent 5,120,469, 1992
(18)Borland, J. E.; Impastato, F. J.; Smith, K. R. U S . Patent 5,254,735, 1993.
= 0.53
M
0
[DDAO], = 0.41 M
A
[DDAO], = 0.19 M
0.1 0 0
time (min)
1
20
40
60
EO
100
120
140
time (min)
Figure 3. Effect of initial DDAO concentration on conversion of DDA at 70 "C for reactions with initial concentrations of [HzOzlo = 0.31 M and [DDAIo = 0.16 M.
greatly increases the contact surface area between the amine and aqueous phases, these results indicate that emulsification alone has little effect on the reaction rate. This suggests that the abrupt increase in reaction rate observed near 110 min is due to the production of enough surfactant to exceed the cmc in the aqueous phase so that micelles capable of solubilizing amine are then present. It is therefore the reaction occurring in the micellar pseudophase, and not a t the emulsion droplet interface, that is responsible for the observed catalysis. For the reaction system in Figure 2, a t approximately 180 min, the solution became single-phase, indicating sufficient surfactant had been produced to solubilize all remaining amine. The cmc of DDAO is 0.0019 M in HzO a t 25 "C and decreases slightly with increasing temperature.14 The presence of significant quantities of solubilized DDA is expected to reduce this value. For experiments in which the amine was emulsified during the initial reaction period, partitioning of the surfactant into the emulsion droplets and adsorption of surfactant a t the droplet/ solution interface also may be significant. For these reasons, an exact determination of the amount of surfactant in micellar form at any given time during a reaction is difficult. The cmc of DDAO was not measured a t all reaction conditions explored in this study, but since the total DDAO concentration was a t least 50 times greater than the maximum possible cmc (0.0019 M) in all experiments, errors in the calculated monomer and micellar surfactant concentrations are minimal. Reactions at Various Initial Surfactant Concentrations. The pseudophase model, discussed previously, predicts that the observed rate constant is dependent upon the surfactant concentration. A series of experiments was performed to determine the relationship between reaction rate and the initial surfactant concentration. In these experiments, the initial concentrations of amine and peroxide reactants were the same in each experiment: 0.16 and 0.31 M, respectively. The initial amine oxide concentration was varied from 0.092 M, which is the minimum amine oxide concentration needed to completely solubilize the amine, to 0.652 M; the reaction mixtures were single-phase a t all times. Figure 3 displays the progress for several of these reactions, and Figure 4 illustrates the integrated second-order rate equation, wherefA represents the fractional conversion of DDA and ROis the ratio of initial reactant concentrations ([HzOzld [DDAlo). Both figures show that the.reaction rate increases with a n increase in the initial DDAO concentration. The overall second-order rate constant, kz = k d J [ H ~ 0 ~ 1 , was calculated for these reactions. As shown in Figure
Synthesis of Surfactants by Micellar Autocatalysis
Langmuir, Vol. 11, No. 8, 1995 3011
u$
8
*2
0.9
dn
1.4 1.2
.-
ff
1.0
E
0.8
-
0.6
E
[DDAO],= 0.53 M
e !.
0.8
z
0.2
0.0 0
20
40
60
80
100
120
140
-
0.5 -
0.4 -
0.3
0.4
. .. . . . . .
4
0.7 0.6
-
0.2-
0.1
-
0.0
4
T = 3OoC
. ..
.*
Figure 4. Integral plot of second-order rate equation for experimental data in Figure 3. fA is the fractional conversion of DDA, and the ratio of the initial reactant concentrations is Ro = [Hz02ld[DDAlo.
0
T=42OC
A
T = 5ooc
I
0
time (min)
.
..
.
1.o
1.8 1.6
20
40
60
80
100 120 140 160 180 200
time (min)
Figure 6. Conversion ofDDAvs time at differenttemperatures with excess hydrogen peroxide. Initial concentrations: [HzOZJO = 2.0 M, [DDAIo = 0.1 M, [DDAOlo = 0.24M.
0.16
T = 3OoC
0.04
0.02 0.00 0.0
.' 0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
PDAOIo (MI
0
20
40
60
80
100
120
time (min)
Figure 6. Effect of initial DDAO concentrationon the overall second-order rate constant, kz,for reactions at 70 "C. Initial concentrations: [HzOzlo = 0.31 M, [DDAlo = 0.16 M.
Figure 7. Integral plot of pseudo-first-orderrate equation for experimental data in Figure 6. f~is the fractional conversion of DDA.
5, there is a sigmoidal relationship between the secondorder rate constant and the initial concentration of surfactant, with the rate constant almost tripling invalue over a n initial surfactant concentration range of 0.3-0.5 M before leveling off. The pseudophase model predicts a similar relationship between the pseudo-first-order rate constant and the overall surfactant concentration.'j If the pseudophase model is considered in the limiting cases of very small and very large surfactant concentrations, it is apparent that the value of the overall pseudo-first-order rate constant (k,) will be approximately equal to the aqueous-phase first-order rate constant at low surfactant concentrations and the micellar first-order rate constant at high surfactant concentrations. The transition from aqueous to micellar rate constant begins at the cmc, above which the surfactant is able to form micelles that solubilize the reactant. Figure 5 may be interpreted in a similar manner; however, the transition region in Figure 5 occurs at surfactant concentrations several orders of magnitude higher than the expected cmc, particularly when the presence of solubilizate is known to lower the cmc somewhat. Figure 5 may indicate a different type of transition, possibly a transition from spherical to rodlike micelles. The value of the overall second-order rate constant before the transition would correspond to the second-order rate constant in spherical micelles, and the value aRer the transition would correspond to a modified second-order constant in rodlike micelles. If this is the case, then the increased reaction rate could be the result of greater solubilization of DDA by rodlike micelles. Effect of Temperature on Pseudo-First-Order Reaction Kinetics. Reactions of amine and hydrogen
peroxide were carried out in amine oxide solutions a t temperatures of 30,42, and 50 "C. A20-fold molar excess of hydrogen peroxide was used so that these reactions can be approximated by pseudo-first-order kinetics. Figure 6 shows some sample data taken a t different temperatures. Pseudo-first-order rate constants were determined from the slopes in Figure 7, and an Arrhenius activation energy was calculated to be 65.4 kJ/mol. It is interesting to note in Figure 7 that the pseudo-first-order rate constant is apparently constant during the reaction, even though the pseudophase models predict that the rate constant changes with surfactant concentration, which in micellar autocatalysis is changing with time. A similar set of experiments was conducted using ethanol to solubilize the amine instead of a micellar amine oxide solution; this reaction medium was assumed to be non-micellar. Reactions were conducted a t 25, 30, 35, and 40 "C.The data were subjected to the same analysis, and a n Arrhenius activation energy of 53.9 kJ/mol was obtained. A graphical comparison of how the reaction rate behaves with temperature in the two solvent systems is shown in Figure 8. It should be noted that the activation energy is of the same order of magnitude in both systems. This indicates that there is no dramatic change in reaction pathway or lowering of activation energy which is common in conventional catalysis; in fact, the activation energy in the micellar medium is only slightly higher, which could be due to increased diffusional resistance of the peroxide into the micelles. This finding gives further support to the claim6that in many cases micellar rate enhancements are caused primarily by concentration effects and not by actual changes in the reaction rate constant.
Kust and Rathman
3012 Langmuir, Vol. 11, No. 8, 1995 -2.5
-5.5
7
I
J
0.00305
0.00315
0.00325
0.00335
1IT (K')
Figure8. Arrhenius plot comparingthe effect of temperature on k, for the reaction of DDAwith hydrogen peroxide in DDAO/ HzO and in ethanol/HZO solutions.
E '1
0.8 T=42OC
no
0.6
2
0.5
f
0.4
-3
0.3
f
-I
T = 3OoC
0.7
0.2 0.1 0.0 0
40
a0
120
160
200
time (min)
Figure 9. Application of eq 15 to experimental data in Figure 6.
Figure 9 shows how the simple model (eq 15) fits the first-order data from Figure 6. The model generally fits the data well up to the first 50- 100 min of the reaction, after which there can be a considerable deviation. This deviation occurs when the reactant conversion has reached
about 90%. One probable cause for this deviation is the magnification of experimental error by the log function as the argument of the log function approaches zero at high conversions. Another possible cause of this deviation could be that at high conversions, some ofthe assumptions used to develop this model, specifically the assumption that all the lipophilic reactant is solubilized in micelles, become less valid.
Conclusions This work has demonstrated that the synthesis of dodecyldimethylamine oxide can be performed under conditions of micellar autocatalysis with excellent reaction rates and yields; conversions of 90-100% are obtained within 2 h at 70 "C. The reaction is characterized by a n activation energy of 65 kJ/mol, and the same reaction in a n ethanol solvent system exhibits an activation energy of 54 kJ/mol. This finding indicates that the presence of surfactant micelles does not markedly affect the reaction pathway or mechanism and that the rate enhancements are primarily due to the localized concentration of reactants near the micelle surface. Analysis of rate data for reactions without any surfactant present initially demonstrates an increase in reaction rate during the reaction. This indicates that initially a n emulsion reaction takes place until enough surfactant is formed to fully solubilize the amine, at which point the system exhibits much faster micellar-enhanced reaction kinetics. Also, reaction rate data collected for various initial surfactant concentrations indicate that there is a sigmoidal dependence of the second-order rate constant upon the initial surfactant concentration. The increase in the rate constant could indicate a possible transition from spherical to rodlike micelles, and the rodlike micelles apparently provide more rate enhancement compared to spherical micelles.
Acknowledgment. This work was supported by The Ohio State University and the National Science Foundation (Grant No. CTS-9308592). LA950292V
