J . Chem. SOC., Faraday Trans. I , 1984,80, 1651-1655 Kinetics of Reactions Involving Three Organic Substrates in Microheterogeneous Systems BY MICHAEL J. BLANDAMER,* BARBARA CLARK AND JOHN BURGESS Department of Chemistry, The University, Leicester LE 1 7RH AND JOHN W. M. SCOTT Department of Chemistry, Memorial University of Newfoundland, St. John's, Newfoundland, Canada Received 25th November, 1983 The rate constants have been obtained for reactions of Crystal Violet, Malachite Green and 2,4-dinitrochlorobenzene with hydroxide ions in two aqueous microemulsions. The rates of reaction are significantly faster than in the corresponding aqueous solutions. For the reaction involving the chlorobenzene, kinetic evidence is reported for a two-state reaction. The catalysis is assigned to interphase effects in the microheterogeneous system.The rates of reactions1 of two low-spin iron(1r) complexes with hydroxide ions in certain aqueous mixtures are significantly faster than in the corresponding aqueous solutions. This acceleration was interpreted in terms of microheterogenity in these aqueous systems combined with catalysed reaction at an interphase. In this paper we report kinetic data for reaction in these same systems between hydroxide ions and Crystal Violet (CV), Malachite Green (MG) and 2,4-dinitrochlorobenzene (DNCB); the kinetics of the aquation reaction for CV and MG in these systems are also commented on, together with qualitative observations on the corresponding reactions with cyanide ions. Again, for reactions involving hydroxide ions, a marked increase in the rate of reaction is observed over that in aqueous solution.The reactions involving CV and MG are first order in organic substrate but the dependence of rate constant on [OH-] is not linear. The dependence of [DNCB] on time is analysed using an equation involving two exponential functions of time. This observation is discussed in terms of the mechanism of reaction. EXPERIMENTAL MATERIALS The kinetics were examined in two aqueous microemulsions prepared as described previously.' Medium A comprised 60 mol % 2-butoxyethanol+20 mol % decane and 20 mol % water. Medium B comprised 40 mol % hexane + 45 rnol % propan-2-01 + 15 mol % water. DNCB, CV and MG were purified as described KINETICS The rates of reaction were measured at 298 K using the method described previously.2 The reactions involving CV and MG were monitored by following the disappearance of reactant at 579 and 625 nm, respectively. The reaction involving DNCB was monitored by following the appearance of the corresponding phenol at A = 358 nm.In these experiments the computer-controlled spectrophotometerl was instructed to collect data at two sets of time 16511652 KINETICS IN MICROHETEROGENEOUS SYSTEMS intervals. The first set comprised data at short time intervals; the second set described longer time intervals. This procedure was necessary in order to delineate two time constants characterising the overall reaction (uide infra). KINETIC ANALYSIS As a preliminary to the presentation of results we consider the kinetics of the following reaction scheme k,*[OH-] k3 X X OH -products.(1) k2 In the limit [OH-] 9 [XI, the product k,*[OH-] can be replaced by a first-order rate constant k,. The algebra is simplified if we set [XI = a, [x - OH] = b and [products] = c. For the experiments reported here, at t = 0, a = a", b" = 0 and co = 0. Hence4 and c/ao = 1 Y 2 - Y 1 -(a)exp(-y,t)-(")exp(-y,t) Y1-YYz (3) where y1 and y, are the two roots of a quadratic equation' in k,, k, and k,. follows first-order kinetics The dependence of c on time is given by In the limit that k, 9 k , (i.e. a consecutive reaction), the dependence of a on time (4) a = ao exp (-k,t). k3 ciao= l-(-)eXP(-klt)- k3 - kl (A) exp(-k3t). ( 5 ) In the limit that k, 9 k, and either k, $= k, or k3 9 k,, the dependences of a and c on time are determined by first-order rate constants.For the case where the appearance of products follows either eqn (3) or (9, the dependence on time conforms to the following general equation : (6) For both cases, the condition is that a,/a, = p2;/p1. However, the meanings attached to these quantities differ. If the quantity c is determined by the absorbance at wavelength il where the extinction coefficient is E , the dependence of absorbance P on time t is given by c/ao = 1 - a, exp ( -p2 t) - a, exp ( -pl t). p = g,-g, exp ( - I 3 2 0 3 3 exp (-A 0. (7) Hence at time t, P is a function of five parameters: g,, g,, g3, p1 and p,. The computer analysis used the Wentworth method5 to calculate' by an iterative method the five unknown parameters starting out with initial estimates (or guesses). It is important to measure P over times where first p1 and then p2 dominate the time dependence of P, otherwise the least-squares analysis5 allows one exponential term to dominate, assigning residual scatter to the other exponential term.In certain cases the dependence of P on time could be fitted by a single-exponential function. These data were analysed in terms of a single first-order rate constant.M. J. BLANDAMER, B. CLARK, J. BURGESS AND J.W.M. SCOTT 1653 1.0 2.0 3 [ NaOH] / 1 0'3 mol dm-3 Fig. 1. First-order rate constant for reaction between Crystal Violet and hydroxide ions in medium B. RESULTS The kinetics of reaction in media A and B are complicated. Furthermore, where the rate of reaction is accelerated over that in more conventional aqueous solutions, the rate of reaction is often extremely fast and can therefore present a number of practical problems.In certain cases it proved impossible to lower [OH-] and thereby slow the rate to within manageable limits. The complexity of the systems (e.g. microheterogenity) was evident because at the other extreme the solutions became opaque at high concentrations of added hydroxide and cyanide. We confine our attention to systems which are clear and transparent, at least in the visible region. At 298 K, in medium B the reaction between CV and hydroxide ions follows first-order kinetics; k = 1.2 x 10-l s-l where [OH-] = 8.99 x lo-* mol dm-3, which corresponds to an effective second-order rate constant of 133.5 dm3 mol-l s-l.The corresponding second-order rate constant in aqueous solution is 2 x 10-l dm3 mol-l s-l. For a simple bimolecular process, the first-order rate constant should be a linear function of [OH-]. This is not the case for the reaction between CV and OH- in medium B (see fig. 1). The shape of the plot points to some degree of saturation in the catalysis. The reaction between CV and cyanide ions in this medium was also very fast; it was impossible to characterise the rate using the available apparatus. The acceleration in rate is comparable to that observed for the reaction with hydroxide. Only a modest change in the spontaneous rate of aquation3 was noted. A similar enhancement of rate was observed for the reaction between OH- and CV in medium A; e.g. k , = 0.74 x 10-1 s-l where [NaOH] = 5 x lop4 mol dm-3.The dependence of k , on [NaOH] is marked. A modest increase in [NaOH] to 5.5 x mol dm-3 produced an increase in k , to 2.5 x 10-1 s-l. Further increase in [NaOH] resulted in a rate of reaction which was too fast for us to measure, i.e. A pattern similar to that described above was observed for the reactions involving MG. In practice the rates of reaction involving OH- and CN- were extremely fast k, > 1.0 s-l.1654 KINETICS IN MICROHETEROGENEOUS SYSTEMS / / i I I / I I I I 0 1 .o 2.0 3.0 [NaOHl mol dm-3 Fig. 2. Dependence of rate parameters [PI(-) and p2(---)] on [OH-] for reaction with 2,4-dinitrochlorobenzene in medium A. and we can only estimate that the second-order rate constant for these reactions in media A and B are > 200 dm3 rnol-l s-l (cf.1.24 x and 5.9 x lo-* dm3 mol-1 s-l for reaction with OH- and CN-). In the experiments described above the kinetics of reaction could be accounted for by a single rate constant. Such was not the case for reaction between OH- and DNCB in medium A when [OH] < 7 x mol dm-3. Under these conditions the data were fitted to eqn (6). With increasing [OH-], both and p2 increase (see fig. 2). However, when [OH-] > 7 x mol dm-3, the kinetics of reaction revert to a simple pseudo- first-order process. The large limits on the second rate parameter in fig. 2 reflect the experimental problems in identifying the two rate parameters in a statistically meaningful fashion. The fact that the data points for p1 and k fall on a continuous curve indicates that the p2 component continues to increase.The kinetic data over the range 3 x < [OH-]/mol dm-3 < 7 x satisfied the criteria set by eqn (3) and (5). Over nine independent measurements, pl/& = 0.24k0.08 and al/a, = 0.23k0.06. Where the data fitted a single first-order rate constant, the calculated second-order rate constant is significantly faster than that for aqueous solution2 (e.g. 1.3 x lov3 dm3 mol-1 s-l at 318 K). However, this calculation can be misleading because the measured first-order rate constant in medium A is not a linear function of [OH-]. Again there is a trend towards saturation with increase in [OH-], as shown in fig. 2. Nevertheless the enhancement in rate of reaction is dramatic if difficult to quantify. DISCUSSION The extent to which the rates of the chemical reaction reported here are enhanced by conducting the reaction in a microemulsion is in itself remarkable. However,M.J. BLANDAMER, B. CLARK, J. BURGESS AND J.W.M. SCOTT 1655 quantitative treatment of microemulsion catalysis is not straightforward. Previously we explained the complicated kinetics of reaction between OH- and an iron(I1) complex in terms of catalysis at an interphase. In these terms, the microheterogeneity of microemulsions is central to their catalytic function. Various aspects of the data reported here support this conclusion. The non-linear dependence of the first-order rate constant on [OH-] for a bimolecular process points to the rate constant being a composite of parameters describing a number of different processes, e.g.competition for sites at a surface where reaction occurs.1 Previouslyf we found that the kinetics of reaction involving an iron complex in certain systems could be described by an equation which contained both zero- and first-order components. These observations pointed directly to a surface-catalysed reaction. For the reactions reported here the evidence is less strong but the combination of acceleration in rate and complexity of the dependence of k on [OH-] points to a similar surface-based reaction. In reviewing the kinetic data it is important to bear in mind how the kinetic parameters are obtained, i.e. the ‘reporter’ of the reaction. Thus for the reactions involving CV and MG, the rate constants were obtained by monitoring the dependence of reactant, a, on time [eqn (2) and (4)].In terms of the reaction scheme given in eqn (I), we can only exclude the conditions required by eqn (2). However, for the reaction between OH- and DNCB the reporter is the product of reaction, Where two kinetic parameters (e.g. PI and PZ) are identified, the data can be understood in terms of either eqn (3) or (5). In both cases the overall rate of reaction is determined by parameters describing the formation and decomposition of an intermediate. In other words, the reaction between DNCB and OH- proceeds via an intermediate. For reaction in aqueous solution the kinetics of reaction can be described by a single rate constant. However, there are good grounds for writing the mechanism of reaction as two stage;6 the first stage involves the formation of the [DNCB - OH]- carbanion where both C1 and OH are bonded to the aromatic ring through the same carbon atom. It is tempting to understand the kinetics of reaction reported here in terms of the production of this intermediate at the interphase. If this assignment is correct, kinetics of reaction in microemulsions may have an important r6le to play in mechanistic studies. The distribution of polar and non-polar reactants, products and intermediates between polar, non-polar and interphase regions could decouple rate parameters in an informative fashion. We thank the S.E.R.C. for a maintenance grant to B.C. ’ M. J. Blandamer, B. Clark, J. Burgess and J. M. W. Scott, J . Chem. SOC., Faraday Trans. I , in press. M. J. Blandamer and D. J. Reid, J. Chem. SOC., Faraday Trans. I , 1975, 71, 2156. K. Hillier, J. M. W. Scott, D. J. Barnes and F. J. P. Steel, Can. J. Chem., 1976, 54, 3312. M. J. Blandamer, E. Ralph, J. M. W. Scott and R. E. Robertson, J . Chem. Soc., Faraday Trans. 1, 1983, 79, 1289. W. E. Wentworth, J. Chem. Educ., 1965, 42, 96. C. A. Bunton and L. Robinson, J . Am. Chem. SOC., 1968,90, 3972. (PAPER 3/2095)