J . Chem. SOC., Faraday Trans. I , 1985,81, 335-341 A Kinetic Study of the Formation and Dissociation of the Meisenheimer Complex Formed between 1,3,5Trinitrobenzene and the Hydroxide Ion in Micellar Dodecyltrimethylammonium Bromide Solution BY JOEL LELIEVRE* AND RENE GABORIAUD E.N.S.C.P., Laboratoire de Physicochimie des Solutions, 11 Rue Pierre et Marie Curie, 75231 Paris Cedex 05, France Received 29th March, 1984 The rates of formation and dissociation of the thermoadditioncomplex of 1,3,5-trinitrobenzene with OH- has been investigated in micellar solutions of dodecyltrimethylammonium bromide (DTABr) using a stopped-flow spectrometer. We have observed and measured three processes which proceed under different conditions of acidity: (a) formation of the r~ complex, (b) spontaneous (equilibrium) dissociation in less basic solutions and (c) dissociation via proton attack in acid solutions.These processes are influenced by the background KBr electrolyte. The interphase potential A$ = drn - &, depicted in the pseudophase model, allows us to explain the variations of the rate constants. The catalytic effect of cationic detergents, for example dodecyltrimethylammonium bromide solutions with water (DTABr), on nucleophilic aromatic substitution reactions has been studied by several workers.l-13 1 -Halogeno-2,4-dinitrobenzene compounds have often been used and reacted with the hydroxide ion, OH-, to form the corresponding phenate; the catalytic effect on such reactions is significant. In contrast, few investigations14-16 have been made on aromatic substitutions in which 0 addition complexes (or Meisenheimer complexes) are formed, e.g.when the aromatic substrate is a trinitrosubstituted benzene compound. The aim of this work is to investigate whether a cationic detergent (DTABr) has an active effect on the rate of formation of such complexes between 1,3,Strinitrobenzene (TNB) and the OH- ion. This aromatic substrate shows simple behaviour in sodium hydroxide solutions because there is no substitutable group (or leaving group) and the reaction is k , limited to k-1 TNB+OH- eTNBOH- (1) in which the 0 addition complex (TNBOH-) is stable in micellar solutions. We have already studied these reactions for several mixtures of water and methanol” and, depending upon the conditions of acidity used, we were able to describe several reaction processes.When the medium is strongly basic, the principal reaction is to give the o complex TNBOH- and only the rate constant k, may be determined experimentally. For a less basic range, the rate constant k-, is also accessible and corresponds to dissociation of the Meisenheimer complex, caused by thermal motion of solvent molecules. For strongly acid mixtures, dissociation occurs by direct attack and the reaction scheme is written as k,[H+l TNBOH--TNB + H,O. 3353 36 REACTION OF TRINITROBENZENE WITH OH - IN MICELLAR SOLUTION Fig. 1. Theoretical diagram showing the logarithm of the apparent constant A for the processes of ionization of TNB [reaction (I)] and dissociation of the TNBOH- complex [reaction (2)]. 3.= k,[OH-]+k-,+k,[H+]. The overall behaviour is summarized diagrammatically in fig. 1 . The effects of DTABr upon these three processes have now been investigated and this paper reports our results. EXPERIMENTAL DTABr detergent was used at a concentration of 2 x lo-' mol dm-3, higher than the critical micellar concentration (c.m.c. = I .56 x lop2 mol dm-3), and normally in the presence of an excess of background electrolyte in order to mask the secondary effects of the various buffer salts in solution, even though the background electrolyte is expected to affect the rates of reaction lo- 18. The rates of reaction were always high and a stopped-flow Durrum Gibson spectrometer was used with the measurement cells thermostatted at 20k0.1 "C, as were the two syringes for injection of reagents.The experimental kinetics curves were displayed on the screen of a Tektronix oscilloscope connected directly to the stopped syringe. However, the use of such a method with micellar solutions is difficult and tedious because many microbubbles are formed in the mixing jet during the injections of reagents. Thus, the different solutions had to be injected three or four times for each measurement in order to obtain reproducible results. RESULTS When TNB in micellar DTABr solutions is reacted with a range of concentrations of micellar sodium hydroxide solutions, we can calculate the rate constants k , and k-l, see fig. 2 and 3. On the other hand, the determination of the rate constant k , is experimentally more difficult. The reagents are mixed quickly in the mixing jet of the stopped-flow spectrometer, one being the preformed TNBOH- complex with an excess of OH- ions and the other being a weak acid solution chosen to obtain the desired final pH value.The two solutions of reagents were made up in such a way that, in every one, there were the same concentrations of surfactant (2 x mol dm-3) and background electrolyte (4 x lop2, 8 x lop2, 2 x 10-l or 4 x 10-1 mol dmP3 KBr). Such reagent conditions were necessary in order to ensure that there were no problems caused by viscosity differences in the mixingjet. The excess of salt enables not only the interphase potential but also the ionic surroundings of the micellar structure to be obtained. The preformed TNBOH- complex with an excess of OH- is stable in solutions of DTABr and KBr and the reaction observed after mixing with weak acid solutions corresponds to reaction (2). Indeed, in the measurementJ.LELIEVRE AND R. GABORIAUD 337 2 log [OH-] Fig. 2. Plot of the pseudo-first-order rate constant 2 against log [OH-]. [DTABr] = 2 x mol dmP3 and T = 20 "C. +, Without background electrolyte; 0, 4 x x , 8 x A, 2 x lo-' and .,4 x 10-' mol dmP3 KBr. t 2 1 5 7 9 11 13 PH Fig. 3. Plot of the pseudo-first-order rate constant 2 defined by kapp = A[OH-] against pH. Circles and crosses refer to KBr concentrations of 8 x and 4 x lo-' mol dm-3, respectively. [DTABr] = 2 x lop2 mol dmP3 and T = 20 "C. cell there occurs fast neutralization of the weak acid by the excess of OH- and formation of the buffer solution in situ, which sets the final pH value.Acetic acid and butylammonium, dibutylammonium or pyridinium salts were used in various proportions for obtaining the correct pH value over a range of ca. 6 pH units. The kinetic curves were recorded on the screen of the oscilloscope at a wavelength A = 460 nm, corresponding to the maximum absorbance of the CT complex. For reaction (2) we followed the decreasing absorbance at the same wavelength. In all experiments the concentration of TNB was very low ( 10-4-10-5 mol dm-3) and under338 REACTION OF TRINITROBENZENE WITH OH- IN MICELLAR SOLUTION Table 1. Variation of log mA (pseudo-first-order rate constant in s-l) with log [OH-] for different concentrations of KBr. CDTABr = 2 x lop2 mol dm-3, T = 20 "C; log A = log (k-.l + k,[OH-]) C,,,/mol dmp3 log [OH-] 0 4 x 10-2 8 x 10-2 2 x 10-1 4 x 10-1 - 0.7 - 1.0 - 1.3 - 1.6 - 2.0 -2.3 - 2.6 - 3.0 - 1.85 1.52 1.20 0.80 0.53 0.23 0.05 - 1.94 1.65 1.34 0.98 0.66 0.33 0.17 0.02 - 1.83 1.40 1.12 0.78 0.43 0.2 1 0.06 -0.03 1.60 1.20 0.88 0.54 0.29 0.07 - 0.02 -0.10 1.40 1 .oo 0.70 0.40 0.13 0.0 1 - 0.08 -0.10 such conditions the concentrations of OH- may be considered to be constant.Fig. 2 shows the variation of the pseudo-first-order rate constant, A, with log[OH-] (k = A/[OH-1; cf. caption to table 1) for solutions with and without KBr electrolyte (KBr = 4 x lo-,, 8 x lo-,, 2 x 10-1 and4 x 10-1 mol dm-3). All thedataare summarized in table 1, and fig. 2 leads to several conclusions. When the reaction occurs in the absence of electrolyte, the plot of log A against log [OH-] function is represented by a straight line whose slope is unity.The micellar catalysis is at a maximum compared with the results in the presence of KBr. On the other hand, when the reaction occurs in the presence of electrolyte, the rate constant is decreased in proportion to the increasing concentrations of KBr. For log [OH-] > - 2, the plots of log A against log[OH-] are represented by straight lines whose slopes are unity and for log [OH-] < - 2 these they tend towards a constant value corresponding to k1. The variations of the pseudo-first-order rate constants A with pH are illustrated (fig. 3) for solutions with two KBr concentrations, 8 x lop2 and 4 x mol dm-3. The form of the curves leads to several conclusions depending on the pH range.When pH > 11, the plot of log1 against pH is represented by a straight line whose slope is unity. The reaction is thus second order as expected. When 8 < pH < 11, log A is independent of pH. When 5 < pH < 8, as is seen in the third part of the kinetic diagram, the plot of logA against pH is a line whose slope is - 1, corresponding to the o complex dissociation reaction with the protons present in solution. In this case, we note that the proton reacts with the TNBOH- complex despite the repulsion of the positive micellar charges. Such a result has not previously been observed with cationic micelles, but the corresponding reaction with anionic surfactants (symmetrical trend) has been described for micellar dodecylsulphate in alkaline media.l3 When the concentration of the KBr electrolyte is increased, the rate constant k , invariably decreases.However, the k-, rate constant is modified very little by added salt. Finally, the dissociation rate, following reaction (2), is increased by adding KBr and we notice a comparable translation for the two lines of slope 1 and - 1 in the presence of salt. Thus the rate constants k , and k, are changed by almost the same absolute value. We were careful to check that this was not due to pH variation in the micellar solution.J. LELIEVRE AND R. GABORIAUD 339 DISCUSSION Two theories are able to show the logical development of the effects of ionic surroundings on micellar catalysis, one being based on the interphase potential of the two phases A# = #M - #w (& and dw are the potential of the micellar and bulk phases, respectively) and the other on ionic exchange on the micellar surface.These two models have been applied by different workers and from time to time they have been considered as mutually exclusive. However, we believe that they are compatible but that the interphase potential model is able to explain our experiments more simply. In this paper we have chosen to show the influence of the addition of a salt on the micellar catalysis of a model reaction [reactions (1) and (2)]. The presence of the salt enable us to set the interphase potential and for that reason we used different concentrations in order to experiment with several values of A#. It is known that the presence of salt generally involves a decrease in the rate of the reaction: either by competition between two anions (Br-, the common anion of the surfactant and the background electrolyte, and the OH- anion) or by decreasing the interphase potential and correspondingly the rate constant. The three processes of the reaction are influenced differently by increasing concentration of electrolyte: (a) a decrease of the rate constant k,, as has been described by several authors, (b) an unmodified value of the rate constant k-, and ( c ) an increase of the rate constant k,.A qualitative interpretation becomes obvious if we use the interphase potential expression. We have already 24 an expression for the interphase potential using the pseudophase model : where si and bi are the parameters of each species of ion which are determined experimentally for kinetic results, [ai] is the ionic concentration in the bulk phase and A& is a constant depending on amphiphiles, counter ion, solvent and temperature.Using eqn (i) we are able to foresee the effect of a variation of electrolyte concentration and we believe, in accord with other 26 that an excess of salt will give rise to an interphase potential difference. If we develop eqn (i), taking account of the anions present in solution, we use a cationic amphiphile (DTA+, Br-), a reactive anion OH- and a fixed excess of back- ground electrolyte (K+, &-), whose anion is common with the counter ion of the surfactant. Consequently eqn (i) may be written as Eqn (ii) explains the variation of A# with varying concentration of background electrolyte. We do not need the values of si and bi for explaining qualitatively our kinetic results, nevertheless we have dete~mined~~9 2 4 9 27 these coefficients using other kinetic results and also for shifts of the deprotonation equilibria. Thus, choosing the OH- ion as a reference sOH- = 1, we have obtained sBr- = 25.Consequently the micellar structure has a selectivity which is 25 times larger for Br- than the OH-. We are in agreement with the selectivity magnitude for these ions found by Bunton et aZ.10v21 The contribution of the [OH-] term in eqn (ii) is thus very small in com- parison with that for the Br- ion and eqn (ii) may therefore be simplified to340 REACTION OF TRINITROBENZENE WITH OH- IN MICELLAR SOLUTION The charge transfer between two phases involves the development of the transfer rate concurrently with Ab: Ad (iv) F logmk = constant-- 2.3RT where "k is the kinetic micellar rate constant (second order) and the constant is the sum of several constants: one being the value of the kinetic rate constant in water and the other the ratio of transfer activity coefficients of several species in solution.We mentioned previously that we have determined experimentally the pseudo- first-order rate constant mA, which can be reproduced by taking account of the expressions for log "k and Ad: log sBr-[Br-]. When the salt concentration increases, the interphase potential decreases and consequently the rate of the TNB ionization reaction decreases (slope + 1 in fig. 2 and 3). The thermal dissociation process (rate constant k - l ) cannot be modified by a A4 potential variation, as is observed.On the contrary, for the dissociation of the complex caused by proton attack (slope -1 in fig. 3), we obtain an inversion of reactivity, the rate constant k, increasing with the increasing electrolyte concentration. Again, a decrease in Ad causes a reduction in ionic repulsion between the proton and the micelles and reaction is promoted. This is a very interesting experimental example of an increase of rate constant obtained by increasing the electrolyte concentration, the opposite effect having always been described for cationic amphiphiles previously. 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