Br- in Acid

Micellar Effects on the Ionic-Redox Reaction BrO3-/Br- in Acid Medium. Ana Domínguez, and .... R. Yeşim Talman , Sinem Göktürk , Melda Tunçay. Colloid...
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Langmuir 1998, 14, 2677-2683

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Micellar Effects on the Ionic-Redox Reaction BrO3-/Br- in Acid Medium Ana Domı´nguez and Emilia Iglesias* Departamento de Quı´mica Fundamental e Industrial, Facultad de Ciencias, Universidad de La Corun˜ a, 15071-La Corun˜ a, Spain Received July 18, 1997. In Final Form: March 2, 1998 The redox reaction between bromate and bromide ions in aqueous acid medium was studied in the absence and presence of cationic and anionic surfactants. The rate equation in water was found to be v ) k3w[BrO3-][Br-][H+]2. The cationic surfactant, tetradecyltrimethylammonium bromide, catalyzes the reaction, using either HBr or HNO3 as the source of H+; nevertheless, the observed rate constant decreases slightly, goes through a minimum, and increases again as the tetradecyltrimethylammonium nitrate concentration increases. The anionic surfactant, sodium dodecyl sulfate, strongly inhibits the reaction when its concentration is increased, reaching a leveling off at high surfactant concentration; but hydrogen dodecyl sulfate catalyzes the reaction. These micellar effects are quantitatively explained on the basis of the pseudophase ion exchange model: tetradecyltrimethylammonium bromide increases the Brconcentration, while sodium dodecyl sulfate decreases the H+ concentration (through the exchange with Na+ counterions) in the water phase where the reaction takes place.

Introduction The modification of chemical reactions by incorporating reagent molecules into aqueous micelles and other organized assemblies has received considerable attention.1-3 Chemical reactivity,4 equilibria,5 and even stereochemistry6 have been significantly affected by micelles. Effects of micelles on rates, and on equilibria, are treated quantitatively in terms of the pseudophase model,7-9 in which micelles and water are treated as distinct reaction media and in which the overall rate is the sum of the rates in each pseudophase. Regarding unimolecular reactions, partitioning of only one substrate between both pseudophases must be considered. Associations of molecules are governed by hydrophobic and electrostatic interactions. Rate enhancements may be caused by ground-state destabilization or transition-state stabilization, possibly because of a change in the solvation of the micelle-bound reactant. Reaction rates between neutral and charged molecules depend on the uptake of both reactants by the aggregate. Simple electrostatic considerations lead to the conclusion that cationic micelles catalyze reactions be(1) (a) Fendler, J. H.; Fendler, E. J. Catalysis in Macromolecular Slystems; Academic Press: New York, 1975. (b) Fendler, J. H. Membrane Mimetic Chemistry; John Wiley & Sons: New York, 1982. (2) (a) Bunton, C. A.; Savelli, G. Adv. Phys. Org. Chem. 1986, 22, 231. (b) Bunton, C. A.; Nome, F.; Quina, F. H.; Romsted, L. S. Acc. Chem. Res. 1991, 24, 357. (3) Mittal, K. L., Ed. Micellization, Solubilization, and Microemulsions; Plenum Press: New York, 1977; Vols. 1 and 2. (4) (a) Garcı´a-Rio, L.; Iglesias, E.; Leis, J. R.; Pen˜a, M. E. Langmuir 1993, 9, 1263. (b) Ferna´ndez, A.; Iglesias, E.; Garcı´a-Rio, L.; Leis, J. R. Langmuir 1995, 11, 1917. (5) (a) Iglesias, E. J. Phys. Chem. 1996, 100, 12592. (b) Iglesias, E.; Leis, J. R.; Pen˜a, M. E. Langmuir 1994, 10, 662. (6) Bianchi, M. T.; Cerichelli, G.; Mancini, G.; Marinelli, F. Tetrahedron Lett. 1984, 25, 5205. (7) Menger, F. M.; Portnoy, C. E. J. Am. Chem. Soc. 1967, 89, 4698. (8) Quina, F. H.; Chaimovich, H. J. Phys. Chem. 1979, 83, 1844. Chaimovich, H.; Aleixo, R. M. V.; Cuccovia, I. M.; Zanette, D.; Quina, F. H. In Solution Behavior of Surfactants; Mittal, K. L., Fendler, E. J., Eds.; Plenum Press: New York, 1982; Vol. 2, p 949. Bunton, C. A.; Romsted, L. S. In Solution Behavior of Surfactants; Mittal, K. L., Fendler, E. J., Eds.; Plenum Press: New York, 1982; Vol. 2, p 975. (9) Martinek, K.; Yatsimirski, A. K.; Levashov, A. V.; Berezin, I. In Micellization, Solubilization, and Microemulsions; Mittal, K. L., Ed.; Plenum Press: New York, 1977; Vol. 2, p 489.

tween hydrophobic neutral substrates and anionic species and inhibit reactions with cations, while the opposite holds for anionic micelles. Rate increases between nonhydrophobic ions of oppositely charged surfaces are the consequence of reactant localization by electrostatic attractions and transition-state stabilization. In this work we study the micellar effects of anionic and cationic surfactants on the redox reaction between bromate and bromide ions in acid medium to generate bromine (and tribromide ion if the reaction is carried out with a great excess of bromide). Though reasonably soluble in water, bromine is very toxic and unpleasant to work with; furthermore, its ready evaporation from aqueous solutions implies that precautions have to be taken against loss during reactions/kinetic runs or with stored solutions. To avoid these problems in bromination kinetic studies, the method generating bromine in situ within the reaction cell is sometimes preferred. Reaction of bromate ion (by using the primary standard KBrO3) with bromide ion in acid aqueous solution yields bromine quantitatively.10,11 The stoichiometry of the reaction is stated in eq 1.

BrO3- + 5Br- + 6H+ f 3Br2 + 3H2O

(1)

We decided to use this method of generating bromine in situ to investigate micellar effects on bromination reactions. We started then investigating the aforementioned reaction in acid medium in water and in aqueous micellar solutions of the cationic surfactants tetradecyltrimethylammonium bromide (TTABr) and nitrate (TTANO3) and of the anionic surfactants sodium dodecyl sulfate (SDS) and hydrogen dodecyl sulfate (HSD), to have a better understanding of the behavior of this reaction in micellar medium. Experimental Section The SDS and TTABr surfactants were Sigma products of the highest purity commercially available. HDS was prepared through the ion exchange of a SDS solution, using a Dowex-50W (10) Young, H. A.; Bray, W. C. J. Am. Chem. Soc. 1932, 54, 4282. (11) Wand, T. X.; Kelley, M. D.; Cooper, J. N.; Beckwith, R. C.; Margerum, D. W. Inorg. Chem. 1994, 33, 5872.

S0743-7463(97)00809-3 CCC: $15.00 © 1998 American Chemical Society Published on Web 04/14/1998

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Domı´nguez and Iglesias

cationic resin, and subsequent elution with distilled water. The Kolthoff reaction was used to check the absence of Na+ ions from HDS. This point was further confirmed with the use of a flame photometer. HDS solutions were used shortly after preparation to avoid any interference from acid hydrolysis.12 TTANO3 surfactant solutions were prepared by using an Amberlite IRA400 anionic resin previously saturated with NO3- ions. The precipitation reaction of Ag+ with Br- ions was employed to check the absence of Br- ions from the TTABr. The KBrO3 and NaBr were Merck products of the highest purity, and the concentration of the acids used (HBr and HNO3) was determined by titrations with standard NaOH solutions (also Merck). Among the possible oxidizing agents present in the reaction medium, nitrate or perchlorate ions have no effect on the bromate-bromide redox reaction. The redox potential corresponding to the reaction being studied, o(BrO3-/1/2Br2) ) 1.48 V, is higher than the standard redox potential of any of the reduction processes which, in acid medium, could occur with the NO3- or ClO4- ions (for example, o(NO3-/1/2N2) ) 1.25 V or o(ClO4-/1/2Cl2) ) 1.39 V);13 that is, BrO3- is the stronger oxidizing agent. In addition, even though in the absence of BrO3- the oxidation of Br- (to Br2) by nitrate or perchlorate ions is thermodynamically favorable (o(1/2Br2/Br-) ) 1.08 V), the reaction rate must be extremely slow. Therefore, we used the acids or the salts of NO3- or ClO4- ions to acidify the reaction medium or to control ionic strength, respectively. Reaction kinetics were observed with a Kontron Uvikon 942 spectrophotometer with notation of the change in absorbance due to the appearance of Br2 (at λ ) 390 nm) or Br3- (at λ ) 266 or 271 nm). The reaction cell was thermostated at 25 ( 0.1 °C. Kinetics were studied by following the integrated method. The initial concentration of BrO3- (1.1 × 10-5 or 5.5 × 10-4 mol dm-3) was always much smaller than that of the other reagents. The integrated first-order rate equation:

Ab ) Ab∞ - (Ab∞ - Ab0)e-k0t

(2)

where Ab, Ab0, and Ab∞ are the absorbance readings at times t, 0, and infinity, was fitted to the kinetic data by using a nonlinear regression analysis program to give the pseudo-first-order rate constants, k0. The good agreement obtained between the experimental and optimized Ab∞ values confirmed that the reaction was of first order with respect to bromate ions.

Results and Discussion Reaction in Water. The spectrum of the reaction mixture 1.1 × 10-5 mol dm-3 of KBrO3 and 2.5 × 10-2 mol dm-3 of HBr revealed an absorption band centered at λ ) 266 nm whose intensity increased with time (see Figure 1a). This band has been ascribed to the absorption of Br3- ion formed from Br2 in the presence of excess Braccording to the equilibrium process stated in eq 3.

Br2 + Br- h Br3-,

Kc

(3)

The corresponding equilibrium constant, Kc, has been reported in the literature as 16.1 ( 0.3 mol-1 dm3 at 1.0 mol dm-3 of ionic strength11 or as 16.8 mol-1 dm3 obtained at an ionic strength of 0.54 mol dm-3.14 Tribromide ion absorbs more strongly than bromine, and at a shorter wavelength: λmax ) 266 nm,  ) 40 900 ( 400 mol-1 dm3 cm-1,11 or λmax ) 267 nm, and  ) 38 000 mol-1 dm3 cm-1 for tribromide ion,15 and λmax ) 390 nm,  ) 175 mol-1 dm3 cm-1, for the case of bromine.11 The absorbance value measured at the end of the reaction as Ab∞ ) 0.381 at 266 (12) Garnett, C. T.; Lambie, A. J.; Beck, W. H.; Liter, M. J. Chem. Soc., Faraday Trans. 1 1983, 79, 953. (13) Bard, A. J., Parsons, R., Jordan, J. Eds. Standard Potentials in Aqueous Solution; IUPAC; Marcel Dekker, Inc: New York, 1985. (14) Scaife, D. B.; Tyrell, H. J. V. J. Chem. Soc. 1958, 80, 386. (15) (a) Daniele, G. Gazz. Chim. Ital. 1960, 70, 1585. (b) Raphael, L. In Bromine Compounds Chemistry and Applications; Price, D., Iddon, B., Wakefield, B. J., Eds.; Elsevier: New York, 1988; pp 369-384.

Figure 1. Spectra of the reaction mixture of 1.1 × 10-5 mol dm-3 of KBrO3 and 0.025 mol dm-3 of HBr: scans 1 to 9, at 5 min interval; scan 10, at infinite time (a) in water and (b) in the presence of 0.050 mol dm-3 of TTABr.

nm (Figure 1a) is in total agreement with the values reported in the literature for the equilibrium constant and the absorption coefficient of tribromide ion. Therefore, the total bromine concentration generated at the end of the reaction, that is [Br2]free + [Br3-], can be quantitatively determined as [Br2]t ) 3[BrO3-] (eq 1). The kinetics of the reaction in water has been studied by several authors; therefore, we report here a brief summary of our results for comparison purposes with those obtained in micellar media. The influence of Br- concentration on the reaction rate of bromine generation from the reaction between bromate and bromide ions was investigated at 0.04 mol dm-3 of HBr, with the ionic strength controlled with NaClO4 kept constant at 0.46 mol dm-3. The results, displayed in Figure 2a, show a linear increase of k0 with the concentration of bromide ion, that is, k0 ) k1[Br-], with the slope of the line being equal to k1 ) (3.94 ( 0.08) × 10-3 mol-1

Micellar Effects on Ionic Redox Reactions

Langmuir, Vol. 14, No. 10, 1998 2679

ionic strength at 0.27 mol dm-3 (controlled with KBr) shows the same behavior trend. The value of k2 determined as 0.83 ( 0.02 mol-2 dm6 s-1 is in good agreement with the previous one. (Notice that [Br-] is included in the values of k2.) These results mean that the rate constant for the loss of BrO3- is given by eq 4. The values of k3w obtained in the several experimental conditions used are reported in Table 1. The influence of ionic strength, controlled with NaClO4, on k3w was also investigated.

k0 ) k3w[Br-][H+]2

Figure 2. (a) Influence of [Br-] in the pseudo-first-order rate constant of the redox reaction of bromate-bromide in the presence of (b) [HNO3] ) 0.04 mol dm-3, I ) 0.60 mol dm-3, and (2) [HBr] ) 0.04 mol dm-3, I ) 0.46 mol dm-3. (b) Influence of the nitric acid concentration on the pseudo-first-order rate constant of the redox reaction bromate-bromide at [Br-] ) 0.20 mol dm-3 and I ) 0.30 mol dm-3.

dm3 s-1. The influence of [Br-] at 0.04 mol dm-3 of nitric acid, at an ionic strength of 0.60 mol dm-3 controlled with NaClO4, was also investigated. The slope of the corresponding straight line was equal to k1 ) (5.07 ( 0.09) × 10-3 mol-1 dm3 s-1. The influence of the acidity was studied at [Br-] ) 0.20 mol dm-3 with the ionic strength kept constant at 0.30 mol dm-3 with NaClO4, and with varying [HNO3] in the (1-10) × 10-2 mol dm-3 range. Typical results of the observed pseudo-first-order rate constant k0 as a function of [H+] are displayed in Figure 2b. Plots of k0/[H+] against [H+] are linear, which means that k0 ) k2[H+]2, with k2 ) 0.627 ( 0.006 mol-2 dm6 s-1 obtained from the slope of the corresponding straight line. The reaction carried out in the presence of HBr at [Br-] ) 0.27 mol dm-3 and with

(4)

As one can see from the data reported in Table 2, at low ionic strength, that is, approximately below 1 mol dm-3, the fourth-order rate constant decreases as the ionic strength increases, whereas a positive salt effect was observed at ca. ionic strength higher than 1 mol dm-3. The theoretical interpretation of the salt effects on k3w can be done by considering the Bro¨nsted equation that relates the variation of the rate constant with the activity coefficients of the reactants and the activated complex16 and with the activity coefficients being calculated by using the Davies equation17 or an extended Debye-Hu¨ckel equation.18,19 This method has been followed by Burgos et al.20 in the interpretation of salt effects in this reaction in aqueous electrolyte solutions of high ionic strength. Following the same procedure, we obtained a value of 6.9 mol-3 dm9 s-1 for (k3w)o, that is, the rate constant extrapolated to I ) 0. Reaction in the Presence of Anionic Surfactants. The influence of the anionic surfactant SDS was investigated for different concentrations of the surfactant (in the range 0-0.25 mol dm-3) at fixed concentrations of HBr (see Table 3). Reactions were monitored by following the increase in absorbance due to Br2 formation at 390 nm, and working with [BrO3-] ) 5.5 × 10-4 mol dm-3. The reaction was not followed at 266 nm (absorption band due to Br3-) because an apparent “induction period” is observed, which becomes more important as the [SDS] increases. This “induction period” is a consequence of the strong association of bromine to the SDS micelles.21 Typical experimental results of the variation of the pseudofirst-order rate constant with the [SDS] are plotted in Figure 3a. Addition of SDS inhibits the reaction. A quantitative interpretation was made by using the pseudophase ionexchange (PPIE) model proposed by Romsted,22 which treats micelles as a pseudophase uniformly distributed in the bulk aqueous phase. Recognizing that the SDS micellar surface is negatively charged, one is led by simple electrostatic considerations to the conclusion that the BrO3-/Br- redox reaction can only occur in the aqueous phase. (16) (a) Perlmutter-Hayman, B. Prog. React. Kinet. 1971, 6, 239. (b) Laidler, K. J. Chemical Kinetics, 3rd ed.; Harper Collins: New York, 1987. (17) Davies, C. W. Prog. React. Kinet. 1961, 1, 163. (18) Brezonik, P. L. Chemical Kinetics and Process Dynamics in Aquatic Systems; Lewis Publishers: Boca Ratan, FL, 1994. (19) (a) Capone, S.; De Robertis, A.; De Stefano, C.; Sammartano, S.; Scarcella, R. Talanta 1987, 34, 593. (b) Rodrı´guez, A.; Bejarano, M.; Ferna´ndez-Boy, E.; Graciani, M.; Sa´nchez, F.; Moya´, M. L. J. Chem. Soc., Faraday Trans. 1992, 88, 591. (20) Burgos, F. S.; Graciani, M. M.; Mun˜oz, E.; Moya´, M. L.; Capita´n, M. J.; Gala´n, M.; Hubbard, C. D. J. Solution Chem. 1988, 17, 653. (21) Domı´nguez, A.; Iglesias, E. Results to be published. (22) Romsted, L. S. In Surfactant in Solutions; Lindman, B., Mittal, K. L., Eds.; Plenum Press: New York, 1984; Vol. 2, p 1015. (b) Romsted, L. S. J. Phys. Chem. 1985, 89, 5107.

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Domı´nguez and Iglesias

Table 1. Experimental Conditions and Values of the Rate Constants k1, k2, and k3w Obtained in the Kinetic Study of the Redox Reaction Bromate/Bromide in Water at 25 °C [H+]/mol dm-3

[Br-]/mol dm-3

I/mol dm-3

0.27 0.20 variable variable

0.27a

variable (HBr) variable (HNO3) 0.04 (HBr) 0.04 (HNO3) a

k1/mol-1 dm3 s-1

0.30b 0.46a 0.60b

(3.94 ( 0.08) × 10-3 (5.07 ( 0.09) × 10-3

k2/mol-2 dm6 s-1

k3w/mol-3 dm9 s-1

0.83 ( 0.02 0.627 ( 0.006

3.07 3.13 2.46 3.17 3.30c

Controlled with NaBr. b Controlled with NaClO4. c Taken from ref. 10.

Table 2. Fourth-Order Rate Constants, k3w, for the Ionic-Redox Reaction BrO3-/Br- Obtained in Aqueous Acid Medium at 25 °C at Different Ionic Strengths I/mol dm-3

k3w/mol-3 dm9 s-1

I/mol dm-3

k3w/mol-3 dm9 s-1

0.0 0.08 0.10 0.15 0.23 0.33 0.44

6.9a/8.4b

0.55 0.70 0.85 1.0 2.0 3.0 4.0

2.96 3.06 3.00 2.79 3.39 4.29 5.78

a

5.02 4.08 3.74 3.22 3.25 2.79

Extrapolated at I ) 0. b Reference 20.

Table 3. Experimental Conditions, Values of β and cmc Used To Fit the Experimental Data to Eq 6, and Values of the Rate Constant k3w Obtained for KNaH ) 0.75 in the Study of the Bromate-Bromide Reaction in Acid Medium in the Presence of the SDS Surfactant Micelles, and Values of k3w and r ()1 - β) Obtained in the Study of the Reaction in the Presence of HDS by Fitting the Results to Eq 8 [HBr]/ mol dm-3

cmc/ mol dm-3

[BrO3-]/ mol dm-3

β

k3w/ mol-3 dm9 s-1 RMSD

0.030 0.040 0.060

SDS 3.0 × 10-3 5.5 × 10-4 0.75 2.5 × 10-3 5.5 × 10-4 0.72 2.0 × 10-3 5.5 × 10-4 0.70

5.67 4.50 4.45

0.064

HDS 3.5 × 10-3 5.5 × 10-4 0.26a

4.53

a

0.042 0.050 0.023

Correspond to the value of R.

Because the reaction requires an acid medium in order to proceed, that is, k0 ) k3w [H+]2[Br-], the decrease of k0 as the SDS concentration increases must be due to a reduction of [H+] in the bulk aqueous phase resulting from the competition between H+ and Na+ surfactant counterions for binding to the micellar surface. The corresponding ionic-exchange equilibrium is given by eq 5, where the subcripts w and m refer to the aqueous and micellar pseudophases, respectively, and the KNaH, the equilibrium constant, has been measured to vary between 0.6 and 1.23

H+w + Na+m h H+m + Na+w

(5)

Starting with the assumption that the reaction proceeds mainly in the aqueous phase, thus BrO3- and Br- ions are excluded from the SDS negatively charged micellar surface, then k0 ) k2[H+]2w; with [H+]w ) [H+]t - [H+]m. The H+ concentration in the micellar phase is defined as mH ) [H+]m/[SDS]m, where [SDS]m ) [SDS]t - cmc. Then the observed rate constant is given by eq 6

k0 ) k3w[Br-] ([H+]t - mH([SDS] - cmc))2

(6)

where the value of mH at each [SDS] can be determined by solving the second-order expression in mH shown in eq 7 by fixing the values of β, the micelle neutralization (23) Pe´rez-Benito, E. Rodenas, E., J. Colloid Interface Sci. 1990, 139, 87.

Figure 3. (a) Influence of [SDS] on the pseudo-first-order rate constant of the reaction between bromate and bromide ions in aqueous acid medium at [HBr] ([) 0.060, (2) 0.040, and (b) 0.030 mol dm-3, solid lines fit eq 6. (b) Influence of [HDS] at [HBr] ) 0.064 mol dm-3 and [BrO3-] equal to (b) 5.5 × 10-4 and (2) 2.2 × 10-4 mol dm-3, solid line fits eq 8.

degree, and of KNaH. Different values of both parameters were attempted in order to get the best estimated values of k3w.

Micellar Effects on Ionic Redox Reactions

(

m2H + mH

[H+]t + KH Na[SDS]t (KH Na - 1)[SDS]m

)

Langmuir, Vol. 14, No. 10, 1998 2681

-β β[H+]t

(KH Na - 1)[SDS]m

) 0 (7)

The experimental data (k0 and [SDS], and the calculated values of mH) were fitted in this manner, i.e., by a simulation procedure,24 to eq 6 to estimate the best values for KNaH, β, and k3w. The corresponding values obtained in the fitting process are reported in Table 3, along with the mean square root deviation (RMSD) of the experimental data from the model. Solid lines in Figure 3a were calculated from fitting of the experimental data to eq 6. The values of k3w (∼5 mol-3 dm9 s-1) agree quite well with those obtained in water (see Table 1) at the corresponding ionic strength. Therefore, the inhibition effect of SDS is qualitatively and quantitatively interpreted on the basis of a reduction of the [H+] in the bulk aqueous phase as a consequence of the exchange between the reactive H+ ions and the inert Na+ micellar counterions. To further confirm this, we studied the influence of HSD, in the presence of which the exchange process cannot take place; a catalysis of the reaction should occur, thus increasing the [HDS] and the total [H+] with it. Figure 3b shows the experimental data of the pseudo-first-order rate constant k0 as a function of [HDS]obtained at [HBr] ) 0.064 mol dm-3 and [BrO3-] ) 2.2 × 10-4 or 5.5 × 10-4 mol dm-3 (see figure legend). As it can be seen, the results show the expected experimentally behavior. Under these experimental conditions, [H+]t ) [H+]ad + cmc + R[HDS]m, then the expression for the observed rate constant is that of eq 8, where R is the micellar ionization degree and [H+]ad the proton concentration added to the reaction medium as HBr.

k0 ) k3w[Br-]([H+]ad + (1 - R)cmc + R[HDS]t)2

(8)

The fit of the experimental data to the above equation taking cmc ) 3.5 × 10-3 mol dm-3 gives R ) 0.260 ( 0.008 and k3w[Br-] ) 0.290 ( 0.007 mol-2 dm6 s-1. The value of R is in total agreement with that determined directly from conductivity measurements; in the same manner the value of k3w obtained as 4.53 mol-3 dm9 s-1 compares perfectly with those obtained in the previous sections (see Tables 2 and 3). The solid line in Figure 3b shows the theroretical fit. Reaction in the Presence of Cationic Surfactants. The bromate-bromide redox reaction was also investigated in the presence of the cationic surfactants TTABr and TTANO3. The spectra of the reaction mixture (1.1 × 10-5 mol dm-3 of KBrO3 and 2.5 × 10-2 mol dm-3 of HBr in the presence of 0.05 mol dm-3 of TTABr) are shown in Figure 1b. We can note (1) a strong increase in the intensity of the maximum absorption band and (2) a red shift of the maximum wavelength absorption in the presence of TTABr, when the comparison with that obtained in water is made (see Figure 1a). Both effects can be explained if one considers that in the presence of the positively charged TTABr micelles, the equilibrium Br2 + Br- a Br3- is completely shifted to the formation of tribromide ion, which by both specific and electrostatic interactions with the positively charged TTABr micelles, binds completely to the micellar surface. (24) Garcı´a-Rı´o, L.; Leis, J. R.; Pen˜a, M. E.; Iglesias, E. J. Phys. Chem. 1992, 96, 7820.

Ionic interactions in micellar systems follow the pattern proposed by Pearson25 of soft and hard ions or the more empirical lyotropic series.26 In this context, the polarizability of an ion, which can be assumed proportional to the number of electrons carried by it, determines its position in a lyotropic series. The higher an ion is in a lyotropic series, the more effective it will be the binding ability of the ion to the micellar surface. Tribromide ion is expected to be in a higher position in a lyotropic series than Br-. In other words, poorly solvated ions show the strongest binding; that is, they can closely approach the charged micellar sufaces. In this sense, the hydration number of Br- is 1 (i.e., the number of water molecules in the primary shell), whereas that expected for Br3- is 0 (as is the case of, e.g., I-).27 On the other hand, the tribromide ion formation equilibrium process is greatly influenced by the solvent,28 increasing the concentration of Br3- as the polarity of the medium decreases. Consequently, since the polarity of the microenvironment that experiences the Br3- ion in the micellar surface is lower than that corresponding to the water phase, a bathochromic shift is observed: λmax ) 266 nm in water shifts to 271 nm in the presence of TTABr micelles. In this sense, as the tribromide ion is more polarizable than the bromide ion, and is an unhydrated ion, the equilibrium exchange constant between them must be favorable for the Br3- ion, which associates completely to the micellar surface. In fact, extrapolating from Ab∞ ) 1.375 the total [Br2] generated (equal to 3[BrO3-]), we can determine the extinction coefficient of Br3- ion at 271 nm as 41 600 mol-1 dm3 cm-1, which is in good agreement with the published values.11 The same value of Ab∞ is readily obtained by working at 5 × 10-3 mol dm-3 of TTABr; that is, the association of tribromide ion with the cationic TTABr micelles is quite effective. Direct measurements of the increase of absorbance due to Br3- as the [TTABr] increases show that at [TTABr] g 2 × 10-3 mol dm-3 all of the [Br3-] is at the micellar surface. A value between 104 and 105 for the association constant of Br3- ion to TTABr micelles is reported in the literature.29 From kinetic results for bromination reactions in acid media21 we were able to estimate a value of 1.5 × 104 mol-1 dm3. The influence of TTABr concentration on the reaction rate of bromate-bromide redox reaction was examined for different surfactant concentrations, in the range of 0-0.2 mol dm-3. The addition of TTABr resulted in an increase of the pseudo-first-order rate constant k0, which increases linearly with [TTABr]. When HBr is used to acidify the reaction medium, the linear function of k0 vs [TTABr] displays a not negligible intercept at the origin (the value of k0 when [TTABr] ) 0); but when HNO3 is used to acidify the reaction medium, the values of k0 extrapolate to a zero-intercept (see Figure 4). However, when the reaction is carried out in the presence of TTANO3, at fixed HBr concentration, the observed rate constant first decreases slightly as [TTANO3] increases, passes through a minimum, and then increases again. This experimental behavior can be understood if one remembers that the Br- is a reagent. Increasing the TTABr concentration in the presence of a fixed amount of HBr also increases the total [Br-]; but when the reaction is carried out in the presence of nitric acid, all Br- comes (25) Pearson, R. G. J. Am. Chem. Soc. 1963, 85, 3533. Pearson, R. G. Science 1966, 151, 172. (26) Larsen, J. W.; Magid, L. J. J. Am. Chem. Soc. 1974, 96, 5774. (27) Israelachvili, J. N. Intermolecular and Surface Forces; Academic Press: San Diego, CA, 1992; p 55. (28) Ruasse, M. F. Advs. Phys. Org. Chem. 1993, 28, 207. (29) Cerichelli, G.; Grande, C.; Luchetti, L.; Mancini, G. J. Org. Chem. 1991, 56, 3025.

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Domı´nguez and Iglesias Table 4. Values of the Rate Constants k3w Obtained from the Intercept Values, and of k3m, Obtained from the Slope Values,a Corresponding to the Linear Plots of the Influence of TTABr Concentration in the Redox Reaction BrO3-/Br- in Acid Medium k3w/mol-3 dm9 s-1

k3m/mol-3 dm9 s-1

5.2 4.8 3.5

4.8 6.6 9.2

dm-3

[HBr]/mol 0.025 0.040 0.075 [HNO3]/mol dm-3 0.035 0.053 0.10

9.8 12.5 16.2

a Using R ) 0.24, see: Domı´nguez, A.; Ferna ´ ndez, A.; Gonza´lez, N.; Iglesias, E.; Montenegro, L. J. Chem. Educ. 1977, 74, 1227.

micellar surface and the reaction depends on [H+]2, we suppose that the reaction cannot occur in this region, but probably in the palisade layer where the ionic strength is higher than that in the bulk water phase; the second term represents this contribution. In an explanation of the different terms, we call the Stern layer to a small thickness of about 4-5 Å, surrounding the micellar core, which includes the headgroups and associated counterions of ionic micelles; but whereas some of the counterions are bound within the shear surface, along with water of hydration, most of them are located in the Gouy-Chapman electrical double layer, that is, the palisade layer, where the counterions are dissociated from the micelle and are free to exchange with ions distributed in the bulk aqueous phase.1a Taking into account that a great amount of BrO3(the limiting reagent) is at the micellar surface (an also that of Br-, the surfactant counterion), even though H+ is excluded from the cationic micellar surface, its concentration in the palisade layer would not be zero, decreasing gradually from water to the micellar surface and increasing with added surfactant and H+, then the reaction could occur in this region.

k0 ) k3w[Br-]ad[H+]2 + k3m[H+]2R[TTABr]

Figure 4. (a) Influence of [TTABr] on the pseudo-first-order rate constant of the reaction between bromate and bromide ions in aqueous acid medium at (2) [HBr] ) 0.040 mol dm-3 and (b) [HNO3] ) 0.035 mol dm-3. (b) Influence of [TTANO3] at [HBr] ) 0.060 mol dm-3. (Dotted line is a guide for the eye.)

from the addition of TTABr and, obviously, one cannot observe reaction in the absence of the TTABr surfactant. In the presence of TTANO3 an exchange between Br- and NO3- ions at the micellar surface occurs. The change in k0 is too small for quantitative treatment because of the high [Br-] used and of the low value of the ionic-exchange equilibrium constant, KBrNO3 ) 0.91.30 The results in Figure 4a may be quantitatively explained on the basis of eq 9. The first term represents reaction via the Br- added to the reaction medium as HBr, and of course, this term is not observed when nitric acid is used as the source of H+. As H+ is excluded from the cationic (30) Bartet, D.; Gamboa, C.; Sepu´lveda, L. J. Phys. Chem. 1980, 84, 272.

(9)

Therefore, for the influence of TTABr studied in the presence of HBr, k3w in eq 9 was determined from the intercept at the origin, and k3m, which was determined from the slope of the corresponding straight lines, stands for the reaction via the ionized Br- ions introduced with the surfactant, with R being the ionization degree of the TTABr micelles and [H+] referring to the concentration in the bulk phase. The results obtained are reported in Table 4. As we can see, k3w values are similar to those obtained in the absence of surfactant; nonetheless, k3m values are higher than the former and increase with the acid concentration. We suppose that this rate constant corresponds to the reaction which occurs at the palisade layer. Then, the k3m values are not referred to the micellar volume due to the assumption made that the reaction does not occur in the Stern layer, thus only the fraction of ionized Br- ions of the surfactant is considered to take part in the reaction. A correct interpretation of the values of k3m requires to remain that this rate constant does not correspond to the rate-determining step of the reaction. On the contrary, depending on the reaction mechanism, k3m will be a function of equilibrium constants (Ki) of those rapid processes previous to the rate-determining step and to the rate constant (k) corresponding to the rate-limiting step. Consequently, the increase of k3m in the micellar region where the reaction occurs can be due to the higher

Micellar Effects on Ionic Redox Reactions

value of the ionic strength in this region, which can modify either any Ki or increase the k value. (The ionic strength in this region has been estimated31 to be around 4 mol dm-3.) For example, if the reaction goes through the attractively simple mechanism suggested by Hinshelwood,32 in which the bimolecular reaction between the conjugate acids of BrO3- and Br- ions is assumed to be the rate-determining step, an increase in k3w could also be related to a variation of the pKa of these strong acids in the reaction region. Another possibility would be the stabilization of the activated complex (if charge separation occurs) by the positively charge micellar surface with the concomitant increase of the rate constant corresponding to the ratedetermining step, k. This latter effect of relating the micellar charge effects to charge distribution in the transition state has been suggested by Bunton et al.2a as being applicable to the elucidation of mechanisms of reactions. In this sense, if k+/k- > 1 (where k+ is related to the reaction in the presence of cationic micelles and kis related to the reaction in the presence of anionic micelles) the transition state has a significant anionic character, as would be the present case, where k+/k-∼ 2.8. Also possible would be the stabilization of a hypothetical intermediate of the reaction, increasing in this manner the corresponding equilibrium constant. The knowledge of the reaction mechanism in water is very necessary to interpret unambiguously the micellar effects. Finally, the higher values of k3w obtained in the presence of TTABr could also be rationalized as due to salt effects. (31) Romsted, L. S. In Micellization, Solubilization, and Microemulsions; Mittal, K. L., Ed.; Plenum Press: New York, 1977; p 509. (32) Hinshelwood, C. N. J. Chem. Soc. 1947, 694. Barton, A. F. M.; Wright, G. A. J. Chem. Soc. A 1968, 1747.

Langmuir, Vol. 14, No. 10, 1998 2683

In fact, recently, Sa´nchez et al.33 have explained the micellar effects on reactions between charged species by using a formulation based on the Bro¨nsted equation. Conclusions The micellar effects on the ionic reaction studied here are quantitatively and qualitatively explained on the basis of the pseudophase ion exchange model. In summary, rate enhancements by TTABr micelles on the reaction between bromate and bromide ions in acid media can be attributed to the effect of the ionic strength resulting from the localization of the reactants in the palisade layer and to the most important effect: the transition state stabilization by the cationic micelles. The inhibition effect of SDS micelles on the reaction rate can be easily explained by relating it to the ionic exchange between H+ and the Na+ counterions; nevertheless, the rate constants k3w are increased by the effect of an increase in the ionic strength of the site where the reaction occurs. These results would be also interesting in the understanding of the micellar effects on the Belousov-Zhabotinsky oscillation reactions.34 Acknowledgment. Financial support from the Xunta de Galicia (projects XUGA 10302A93 and 10302A95) is gratefully acknowledged. Thanks also to the reviewers for their useful comments. LA970809Y (33) Lo´pez Cornejo, P.; Jime´nez, R.; Moya´, M. L.; Sa´nchez, F. Langmuir 1996, 12, 4981 and references therein. (34) (a) Maritato, M.; Nikles, J.; Romsted, L. S.; Tramontin, M. J. Phys. Chem. 1985, 89, 1341. (b) Vanag, V. K.; Hanazaki, I. J. Phys. Chem. 1995, 99, 6944. (c) Ungvarai, J.; Nagy-Ungvarai, Zs.; Enderlein, J.; Mu¨ller, S. C. J. Chem. Soc., Faraday Trans. 1997, 93, 69.