Micellar and Salt Effects on the Binuclear Complex Formation between

linear over more than three half-lives. .... As in the case of the kinetic salt effects, the shift depends ... expected and, thus, a MLCT blue shift w...
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Langmuir 1996, 12, 4090-4094

Micellar and Salt Effects on the Binuclear Complex Formation between Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+ Amalia Rodrı´guez, Marı´a del Mar Graciani, and Marı´a Luisa Moya´* Departamento de Quı´mica Fı´sica, Universidad de Sevilla, C/ Profesor Garcı´a Gonza´ lez s/n, 41012 Sevilla, Spain Received March 7, 1996. In Final Form: May 30, 1996X The binuclear complex formation between Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+ (en ) ethylenediammine, pzCO2 ) pyrazinecarboxylate) has been studied in several aqueous electrolyte solutions and in sodium dodecyl sulfate, SDS, micellar systems in the presence and absence of various salts at 298.2 K. In dilute salt aqueous solutions the main factor operating on reactivity is the influence of the electrolyte concentration on the approaching process between the two ionic reactants of opposite sign, whereas in concentrated salt solutions the direct interaction cation-iron(II) complex seems to be the determining factor controlling changes in the reaction rate upon changing salt concentration. In the case of the SDS micellar systems, kinetic micellar effects have been rationalized by considering the influence of interfacial electrical potential variations, with changing surfactant (or salt) concentration, on the process studied.

Introduction Kinetic studies in micellar systems have been in the scope of interest of many researchers for a long time. The rates of enzymatic, organic, and inorganic reactions have been investigated in the presence of micelles formed by a great variety of surfactants. In regard to inorganic reactions most of the studies have been done on redox reactions, both thermal and photochemical. However, works on substitution processes are more scarce in comparison. We are interested in the study of salt and cosolvent effects on substitution reactions with complexes of the type Fe(CN)5Ln-.1-7 The reaction of Fe(CN)5(4CNpy)3- (4-CNpy ) 4-cyanopyridine) has also been studied in AOT-oil-water microemulsions as a function of the nature of the oil phase, of the surfactant concentration, and of the molar ratio of [H2O]/[AOT].8 Our idea was to study this reaction or a related substitution process in micellar systems. In this regard, and taking into account the solubility problems found by other authors in the study of substitution processes with pentacyanoferrate(II) complexes in micellar and microemulsion systems,9 the substitution reaction between Fe(CN)5H2O3- (aquopentacyanoferrato(II)) and Co(en)2(2-pzCO2)2+ ((bis(ethylenediammine)2-pyrazinecarboxylato)cobalt(III)) ions was chosen to be studied. This reaction renders a binuclear complex characterized by a high molar extinction coefficient ( ) 9.5 × 103 mol dm-3 cm-1 at 635 nm10), and therefore low reactant concentrations can be used. Be* Author to whom all correspondence should be directed. X Abstract published in Advance ACS Abstracts, August 1, 1996. (1) Rodrı´guez, A.; Moya´, M. L.; Lo´pez, P.; Mu´n˜oz, E. Int. J. Chem. Kinet. 1990, 22, 1017. (2) Moya´, M. L.; Barrio, A.; Graciani, M. M.; Jime´nez, R.; Mu´n˜oz, E.; Sa´nchez, F.; Burgess, J. Trans. Met. Chem. 1991, 16, 115. (3) Barrio, A.; Graciani, M. M.; Jime´nez, R.; Mu´n˜oz, E.; Sa´nchez, F.; Moya´, M. L.; Alsheri, A.; Burgess, J. Trans. Met. Chem. 1992, 17, 231. (4) Tejera, I.; Rodrı´guez, A.; Sa´nchez, F.; Moya´, M. L.; Burgess, J. J. Chem. Soc., Faraday Trans. 1991, 87, 2573. (5) Graciani, M. M.; Sa´nchez, F.; Rodrı´guez, A.; Moya´, M. L. Inorg. Chim. Acta 1993, 208, 213. (6) Moya´, M. L.; Burgess, J.; Sa´nchez, F. Int. J. Chem. Kinet. 1993, 25, 469. (7) Rodrı´guez, A.; Pe´rez-Tejeda, P.; Jime´nez, R.; Sa´nchez, F.; Lo´pez, P.; Calvente, J. J.; Moya´, M. L. Int. J. Chem. Kinet. 1995, 27, 807. (8) Lo´pez, P.; Rodrı´guez, A.; Go´mez-Herrera, C.; Sa´nchez, F.; Moya´, M. L. J. Colloid Interface Sci. 1993, 159, 53. (9) Burgess, J.; Patel, M. J. J. Chem. Soc., Faraday Trans. 1993, 89, 783. (10) Toma, H. E. Inorg. Nucl. Chem. 1975, 87, 785.

S0743-7463(96)00208-9 CCC: $12.00

sides, its mechanism has been well established in aqueous solutions.4,7,11,12 In the present work the substitution reaction between Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+ has been studied in several electrolyte aqueous solutions as well as in sodium dodecyl sulfate, SDS, micellar systems. Changes in the reaction rate upon addition of various salts to the micellar systems have also been studied. In all the experiments the temperature was maintained at 298.2 K. Experimental Section Materials. (Bis(ethylenediamine)2-pyrazinecarboxylato)cobalt(III) perchlorate, [Co(en)2(2-pzCO2)](ClO4)2, and sodium amminopentacyanoferrato(II) trihydrate, Na3[Fe(CN)5NH3]‚ 3H2O, were prepared from published methods11,13 and characterized by visible absorption spectra and CHN microanalysis. Sodium dodecyl sulfate, SDS, was obtained from Merck and purified before use.14 Its critical micellar concentration, cmc, obtained by use of conductivity measurements was in good agreement with literature data. All the electrolytes used were obtained from Merck, P. A. grade. Water was obtained from a Millipore Milli-Q water system; its conductivity was less than 10-6 S cm-1. The amminopentacyanoferrate(II) gives the aqua complex Fe(CN)5H2O3- instantly on dissolution (λmax ) 443 nm,  = 700 mol dm-3 cm-1 in pure water15). Therefore, our subsequent discussion will be based on the consideration that the reactant iron(II) species is the aquopentacyanoferrato(II) ions. Kinetics. The rate measurements were performed using a Hi Tech stopped-flow as well as a Hitachi UV-visible spectrophotometer with a manual mixing apparatus SFA-20 from Hi Tech, depending on the reaction time. Kinetics were followed in all cases at the wavelength of the maximum absorbance in the binuclear complex. The thermostatic system maintained the temperature in a range of (0.1 °C. The binuclear complex formed between Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+ goes through a subsequent electron transfer reaction in which the iron(II) center is oxidized and the Co(III) center is reduced. Nonetheless this electron transfer is so slow (ket ∼ 10-5 s-1 in 0.1 mol dm-3 LiClO411) that there is no coupling between the two processes. Kinetic measurements were always taken in fresh solutions. N2 was bubbled through the water before preparing the micellar (11) Malin, J.; Ryan, D.; O’Halloran, V. J. Am. Chem. Soc. 1978, 100, 2097. (12) Guardado, P.; Van Eldik, R. Inorg. Chem. 1990, 29, 3473. (13) Brauer, G. Handbook of Preparative Inorganic Chemistry, 2nd ed.; Academic Press: New York, 1965; Vol. 2, p 1511. (14) Kang, Y. J.; Kevan, L. J. Phys. Chem. 1994, 98, 7624. (15) Norris, P. R.; Pratt, J. M. J. Chem. Soc., Dalton Trans. 1995, 3643.

© 1996 American Chemical Society

Complexes of Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+

Langmuir, Vol. 12, No. 17, 1996 4091

Table 1. Kinetic Data for the Reaction Fe(CN)5H2O3- + Co(en)2(2-pzCO2)2+ in Aqueous Solutions (T ) 298.2 K) 104[Co(en)2(2-pzCO2)2+]/ mol dm-3

kobs/s-1

10-3k/ mol-1 dm3 s-1

5.6 12.5 5.6 12.5 2.0 2.0

1.49 3.21 1.51 3.20 0.62 0.61

2.6a 2.6a 2.7b 2.6b 3.1c 3.0d

a Reference 2, µ ) 0.2 mol dm-3 (LiClO ). b This work, µ ) 0.2 4 mol dm-3 (LiClO4). c Reference 6, µ ) 0.1 mol dm-3 (NaClO4).d This -3 work, µ ) 0.1 mol dm (NaClO4).

and aqueous solutions to prevent oxidation of the iron(II) complex. All the media studied were unbuffered, since the reaction is pH independent in the range 5 < pH < 9.11,12 In all cases the Co(III) complex concentration was in large excess in order to work under pseudo-first-order conditions. Pseudo-first-order rate constants were obtained from the slopes of the ln(A∞ - At) against time plots, where At and A∞ are the absorbances at time t and at the end of the reaction, respectively. Under the working conditions the first-order kinetic plots were linear over more than three half-lives. For the slower reactions at least six determinations contribute to each rate constant value. For the stopped-flow measurements no less than twelve determinations contribute to each kobs value. The rate constants were reproducible within a precision of about 5% or better. Addition of Co(en)2(2-pzCO2)2+ to an SDS solution of low concentration caused a precipitate to form, presumably a cobalt(III) complex-dodecyl sulfate compound, which redissolved on stirring when the SDS concentration was increased enough. This fact put a limitation on the SDS concentration range available. An SDS concentration lower than 0.03 mol dm-3 could not be used. This behavior has been observed for other cationic species.16 This solubility problem does not allow determination of the cmc value in the presence of the cobalt(III) complex. The possibility of studying the reaction in hexadecyltrimethylammonium bromide, CTAB, micellar systems was also considered. However, solubility problems with the cobalt(III) complex in these systems did not permit us to carry out this study.

Figure 1. Plot of log(kobs/s-1) against log([Co(III)]/mol dm-3) for the reaction [Fe(CN)5H2O3-] + [Co(en)2(2-pzCO2)2+] in [SDS] ) 0.1 mol dm-3 in the presence of NaCl, 0.1 mol dm-3. T ) 298.2 K.

Results To test our data, various experiments were carried out under the same working conditions used by other authors. One can see in Table 1 that results obtained in this work are in good agreement with published data. In this table k is the second-order rate constant obtained by dividing kobs by the Co(III) complex concentration present in the working medium. The dependence of the pseudo-first-order rate constant on Co(III) complex concentration was also investigated, in aqueous and in micellar solutions. In all cases a secondorder rate law was found, first order with respect to each of the reactants. Figure 1 shows the plot for [SDS] ) 0.1 mol dm-3 and [NaCl] ) 0.1 mol dm-3. Plots of kobs (s-1) against [Co(III)] (mol dm-3) are straight lines and do not exhibit meaningful intercepts. In SDS micellar systems the wavelength of the maximum absorbance and the molar extinction coefficient of the d-d band in the visible region of the spectrum corresponding to the cobalt complex are similar to those values shown in water. In regard to the iron(II) complex and to the binuclear species formed during the reaction, the molar extinction coefficients of the d-d band corresponding to the former and of the metal-to-ligand charge transfer (MLCT) band corresponding to the latter hardly change in the SDS micellar systems. In both cases the

Figure 2. Dependence of the pseudo-first-order rate constants, kobs/s-1, on surfactant concentration for the reaction [Fe(CN)5H2O3-] + [Co(en)2(2-pzCO2)2+] in SDS micellar systems. T ) 298.2 K.

wavelength of the maximum shifts a little to shorter wavelengths. For this reason, kinetics were followed in dodecyl sulfate solutions at 630 nm. Figure 2 shows the dependence of the observed rate constant on SDS concentration. Figure 3 shows the variations of kobs upon changing electrolyte concentration, at [SDS] ) 0.1 mol dm-3, for various salts. Discussion Table 2 shows that kinetic salt effects for the reaction under study are negative for all the electrolytes studied in dilute as well as in concentrated salt solutions. Kinetic results and activation parameter values corresponding to this and related substitution reactions4,11,12,17 show that the process studied takes place preferentially through a dissociative mechanism, D; that is, k1

Fe(CN)5H2O3- {\ } Fe(CN)53- + H2O k

(1)

-1

(16) (a) Meisel, D.; Matheson, M. S.; Rabani, T. J. Am. Chem. Soc. 1978, 100, 117. (b) Kratohvil, S.; Shinoda, K.; Matijeric, E. J. Colloid Interface Sci. 1979, 72, 106.

(17) Abu-Gharib, E. A.; Ben Ali, R.; Blandamer, M. J.; Burgess, J. Trans. Met. Chem. 1989, 12, 371.

4092 Langmuir, Vol. 12, No. 17, 1996

Rodrı´guez et al. Table 2. Pseudo-First-Order Rate Constants, kobs/s-1, for the Reaction Fe(CN)5H2O3- + Co(en)2(2-pzCO2)2+ and Wavelength of the Maximum Absorbance, λmax, for the MLCT Band Corresponding to the Binuclear Complex [(en)2Co(µ-pzCO2)Fe(CN)5]- in Various Aqueous Salt Solutions (T ) 298.2 K) [salt]/mol dm-3

Figure 3. Dependence of the pseudo-first-order rate constants, kobs/s-1, on electrolyte concentration for the reaction [Fe(CN)5H2O3-] + [Co(en)2(2-pzCO2)2+] in [SDS] ) 0.1 mol dm-3. T ) 298.2 K. k2

Fe(CN)53- + Co(en)2(2-pzCO2)2+98 [(en)2Co(µ-pzCO2)Fe(CN)5]- (2) Applying the steady state approximation, the rate law can be written

ν)

k2k1 [Fe(II)][Co(III)] k-1

(3)

When electrolyte concentration is increased, the observed rate constant, kobs ) (k1k2)/k-1[Co(III)] (with the Co(III) complex concentration staying constant through all experiments), is decreased, which can be explained by (i) a decrease in k2 due to a retardation of the approaching process between two ionic species of opposite sign by increasing salt concentration and (ii) the formation of the Fe(CN)53- species being more difficult with increased electrolyte concentration because of the electrostatic interactions between the Fe(CN)5H2O3- complex and cations which come from the background electrolyte. In regard to this second point, it is interesting to point out that the dissociative character of the substitution reactions with Fe(CN)5Ln- complexes has its origin in the special characteristics of the cyanide ligands. These strong field ligands weaken the Fe-L bond and favor a dissociative mechanism. Taking into account the electronic density on the cyanide ligands in the species Fe(CN)5H2O3-, it seems reasonable to expect the cations, which come from the salt, to interact preferentially with the iron(II) complex in the vicinity of the cyanide ligands.18 Because of the electrostatic interactions, the presence of the cations would diminish the electronic density on the iron(II) metal center, reinforcing the Fe-L bond and, therefore, making release of the water molecule from the inner-coordination shell to the bulk more difficult. Kinetic salt effects observed show a specificity in the nature of the salt present in the reaction medium. For a given ionic strength, the observed rate constant depends on the nature of the cation and is practically independent of the anion nature of the salt. This could be explained on the basis of the reaction mechanism mentioned above. Retardation in the release of the water molecule in path (18) (a) Warner, L. W.; Hoq, M. F.; Myser, T. K.; Henderson, W. W.; Shepherd, R. E. Inorg. Chem. 1986, 25, 1911. (b) Moya´, M. L.; Rodrı´guez, A.; Sa´nchez, F. Inorg. Chim. Acta 1991, 188, 185.

103kobsa/s-1

λmaxa/nm

0.1 6.0

NaNO3 96 17

630 613

0.1 3.0

NaCl 92 28

630 621

0.1

NaClO4 90

631

0.1 0.2 3.0

LiClO4 102 81 21

630 629 616

0.1

LiCl 103

630

0.1 3.0

CsCl 99 32

632 625

0.5 1.0 2.0 3.0 4.0 5.0 6.0

LiNO3 108 39 30 23 18 15 13

627 626 621 616 613 610 607

1.0 2.0 3.0

Mg(NO3)2 27 12 3.6

611 600 591

2.0

Mg(ClO4)2 12

601

1.0 3.0

Ca(NO3)2 10.4 4.5

617 605

1.0 2.0 3.0

Sr(NO3)2 13 7.1 4.7

610 614 607

a k -1 obs(pure water) ) 8.7 s ; λmax(pure water) ) 635 nm. All the kinetic data were obtained at reagents concentrations [Fe(CN)5H2O3-] ) 3 × 10-5 mol dm-3 and [Co(en)2(2-pzCO2)2+] ) 3 × 10-4 mol dm-3.

1 will be greater the higher the polarization power of the cation. Therefore, the trends expected would be kobs(Cs+) > kobs(Na+) > kobs(Li+) and kobs(Sr2+) > kobs(Ca2+) > kobs(Mg2+). This effect is important in concentrated salt solutions in which the approaching process (path 2) will not be substantially affected by changes in salt concentration due to screening effects (considering only Coulombic interactions19). And in fact, in concentrated salt solutions the trends considered above are found. In more dilute salt solutions salt effects do not show a clear specificity, as was expected if the factor controlling reactivity is the effect of changes in salt concentration on the approaching process, path 2. Table 2 also shows the wavelength of the maximum absorbance of the MLCT band in the binuclear complex in the different electrolyte aqueous solutions. In all cases the metal-to-ligand charge transfer band from the iron(19) (a) Dunn, M. H.; Kozak, J. J. J. Chem. Phys. 1986, 76, 984. (b) Khokhlova, A. I.; Shishin, L. P. Russ. J. Phys. Chem. 1983, 57, 1653. (c) Bruhn, H.; Nigan, S.; Holwarth, J. F. Faraday Discuss. Chem. Soc. 1982, 74, 129.

Complexes of Fe(CN)5H2O3- and Co(en)2(2-pzCO2)2+

(II) center to the pyrazinecarboxylato ligand goes through an hypsochromic shift with increasing salt concentration. As in the case of the kinetic salt effects, the shift depends on the nature of the background electrolyte, in particular, on the nature of the cation of the salt. Considering the studies made on the LMCT (ligand-to-metal charge transfer) bands corresponding to various FeIII(CN)5Lncomplexes,18a,20 it is expected that the MLCT transition in the binuclear complex occurs with a decrease in dipole moment in the excited state with respect to that in the ground state. Taking into account that the polarity and polarizability of the medium (measured through the statical and optical bulk dielectric constants) decrease with increasing electrolyte concentration, if the influence of the salt on the MLCT had been through changes in the properties of the solvent affecting solvation of the solvatochromic species, a higher destabilization of the more polar ground state would be expected in relation to the less polar excited state; that is, a bathochromic shift would be expected. Given that the experimental trend is the opposite, the influence of the salt has to work through a direct interaction. Considering that no substantial influence of the anion nature on λmax is observed, the results can be explained as above. That is, the cations from the salts are expected to interact preferentially in the vicinity of the cyanide ligands as well as of the pyrazine π donor ligand and of the carboxylate group. Taking into account that the transition studied is metal-to-ligand, the optical electron transfer results in a decrease in the electronic density of the cyanide ligands and an increase of the π donor ligand in the excited state with respect to the ground state. Since the solvent molecules cannot change their orientation during the electronic transition, a greater cation binding of the cyanide ligands for the ground state than for the excited state in the binuclear complex is expected and, thus, a MLCT blue shift with increasing salt concentration, as in fact, is observed. Some contributions from the interaction with the pyrazine ligand and from the changes in the solvation of the cobalt(III) complex end of the binuclear species can also be expected. However, the experimental trends λmax(Mg2+) < λmax(Ca2+) < λmax(Sr2+) and λmax(Li+) < λmax(Na+) < λmax(Cs+) found in concentrated salt solutions seem to indicate that, at least for high electrolyte concentrations, the interactions between cations and cyanide ligands were the main factor responsible for the solvatochromic shifts observed. After consideration of the mechanism followed by the reaction and the factors controlling the kinetic salt effects in aqueous solutions, kinetic data in micellar systems will be taken into account. Figure 2 shows that in the concentration range from 0.03 to 0.25 mol dm-3 of SDS (which is expected not to alter significantly the aggregation number and the shape of the micelles21) the observed firstorder rate constant increases upon increasing [SDS] at low surfactant concentration, reaching a plateau at about 0.15 mol dm-3 of SDS. Assuming that no changes in the mechanism of the reaction take place, the influence of the SDS concentration on paths 1 and 2 of the reaction will be considered. The negatively charged Fe(CN)5H2O3complex is expected to be mainly in the bulk aqueous phase. Taking into account a degree of counterion dissociation in the range 0.2-0.3 for the SDS micelles,22 (20) Shepherd, R. E.; Hoq, M. F.; Hoblack, N.; Johnson, C. R. Inorg. Chem. 1983, 23, 3249. (21) (a) Mysels, K. J.; Dorion, F.; Gaboriaud, R. J. Chim. Phys. 1984, 81, 187. (b) Bernas, A.; Grand, D.; Hautecloque, C.; Giannotti, C. J. Phys. Chem. 1986, 90, 6189. (22) (a) Shedlovski, L.; Jakob, C. W.; Epstein, M. B. J. Phys. Chem. 1963, 67, 2075. (b) Shanks, P. C.; Frances, E. I. J. J. Phys. Chem. 1992, 96, 1794. (c) Sasaki, T.; Hattori, M.; Sasaki, J.; Nukina, K. Bull. Chem. Soc. Jpn. 1975, 48, 1397.

Langmuir, Vol. 12, No. 17, 1996 4093

and considering a cmc of 8 × 10-3 mol dm-3, a maximum change in the ionic strength of the bulk from 0.015 to 0.085 mol dm-3 is expected. Actually, this can be considered an upper limit in the ionic strength change, since a decrease in the cmc of the SDS micelles is expected in the presence of a cationic species such as the Co(III) complex. Given that solubility problems precluded the experimental determination of the cmc in the presence of the reactant Co(III) cationic complex, its value in pure water has been taken as an approximation. If the change in the ionic strength of the bulk is of the order assumed above, no substantial influence in the formation of the species Fe(CN)53- by changing SDS concentration would be expected according to the kinetic salt effects observed in dilute electrolyte aqueous solutions. Besides, an increase in [SDS] would result in an increase in the ionic strength of the bulk, and therefore the release of the water molecule in path 1 would be more difficult, retarding the reaction. Since the trend observed in the variations of kobs upon changing [SDS] is the inverse, the conclusion is that the important factor controlling reactivity in SDS systems (at least for the lower SDS concentrations studied) is the influence of surfactant concentration on the approaching process path 2. From purely electrostatic considerations, the Co(III) complex can be assumed to reside predominantly in the Stern layer and to a lesser extent in the counterion diffuse layer surrounding the micelles. When the SDS concentration increases, the interfacial electrical potential, ∆Ψ, decreases.21a,b This decrease in ∆Ψ of a negatively charged micelle will favor the approaching of an anionic species, Fe(CN)53-, to the Co(III) complex residing in the micelle surface and therefore will increase the reaction rate. Another way of seeing this point is that an increase in surfactant concentration will increase the free Na+ counterion concentration which can neutralize or screen the micelle charge, favoring the approaching process. At the same time, Na+ can compete with the Co(III) complex molecules for surface sites, thereby lowering the amount of Co(III) molecules bound in the Stern layer and, thus, increasing the effective concentration of the cobalt reactant. The interfacial electrical potential changes more steeply upon changing [SDS] at low surfactant concentration, but at higher [SDS], changes in surfactant concentration scarcely affect the ∆Ψ values.21a Of course, other factors influencing reactivity can be operative such as possible structural changes and the influence of [SDS] on path 1. This could explain the near constancy of kobs for [SDS] > 0.15 mol dm-3. To shed more light on the influence of the interfacial electrical potential on the reaction rate, the effect of added electrolytes on kobs in SDS micellar systems was investigated. Figure 3 shows the variation in kobs, at [SDS] ) 0.1 mol dm-3, with salt concentration in the presence of LiCl, NaCl, and CsCl. In all cases kobs increases with increasing electrolyte concentration, following the trend kobs(Li+) < kobs(Na+) < kobs(Cs+). It is known that the addition of these types of electrolytes to SDS micellar systems decreases the interfacial electrical potential, as shown, for instance, in the data of Hartland et al.23 for SDS systems (0.01 mol dm-3 surfactant concentration) in the presence of several NaCl concentrations. It has been emphasized that this lowering of the ∆Ψ values is higher when the hydrated counterion radius is smaller, following the sequence ∆Ψ(Cs+) < ∆Ψ(Na+) < ∆Ψ(Li+).24 Therefore, on the basis of a reaction rate mainly determined by changes in ∆Ψ, the expected trend would be kobs(Li+) < (23) Hartland, G. V.; Grieser, F.; White, L. R. J. Chem. Soc., Faraday Trans. 1 1987, 83, 591. (24) Ferna´ndez, M. S.; Fromherz, P. J. Phys. Chem. 1977, 84, 566.

4094 Langmuir, Vol. 12, No. 17, 1996

Figure 4. Plot of ln(kobs/s-1) against the hydrated Stokes radii of the different cations, which come from the background electrolyte, present in the SDS micellar systems. The SDS and salt concentrations were 0.1 mol dm-3. T ) 298.2 K.

kobs(Na+) < kobs(Cs+), as in fact was found. Other studies have shown that the nuclear relaxation time, T1, of Cs+, Na+, and Li+ increases in SDS solutions, in respect to aqueous solutions, following the sequence ∆T1(Cs+) > ∆T1(Na+) > ∆T1(Li+).25 At the same time, the degree of dissociation, R, of dodecyl sulfate micelles (obtained from conductivity measurements) increases in the order R(LiSD) > R(SDS) > R(CsSD).26 All these experimental results have the same origin (considering that the electrostatic interactions are the main ones). The smaller the radius of the cation is, the stronger the electrostiction effect on water will be, resulting in a longer hydrated radius, which means that approaching the highly charged surface of the micelle cannot be so close and therefore a less firm attaching of the cation to the SDS micelles results. This causes a smaller decrease in ∆Ψ for more hydrated cations than for less hydrated cations. Given that the lowering in ∆Ψ, originated by the cations, is nearly proportional to their hydrated Stokes radii (at least for [SDS] ) 0.1 mol dm-3 and [salt] ) 0.1 mol dm-3 21a), Figure 4 shows the plot of ln(kobs/s-1) against the hydrated Stokes radii of the cations studied. In this figure the results for (25) Robb, I. A.; Smith, R. J. Chem. Soc., Faraday Trans. 1 1973 , 70, 287. (26) Mukerjee, P.; Mysels, K. J.; Kapanan, P. J. Phys. Chem. 1966, 71, 4166.

Rodrı´guez et al.

the tetraalkylammonium salts Me4NBr, Et4NBr, Pr4NBr, and Bu4NBr (where Me, Et, Pr, and Bu are methyl, ethyl, propyl, and butyl) also appear. For these quaternary salts, except for tetramethylammonium bromide a linear relation between ln(kobs/s-1) and the hydrated Stokes radii is also observed, although the slope has the opposite sign from that found for the hydrophylic electrolytes. In the case of the quaternary salts it seems that a hydrophobic bonding with the exposed hydrocarbon on the micelle surface overcomes the electrostatic contribution of the cation-head group interactions,27 this being responsible for the decrease in the cmc and in the dissociation counterion degree of the dodecyl sulfate micelles as the size of the alkyl group in the quaternaries increases.26 One can consider that the trend in the interfacial potential may be ∆Ψ(Bu4NBr) < ∆Ψ(Pr4NBr) < ∆Ψ(Et4NBr), which could explain the experimental trend observed in kobs. In the case of the tetramethylammonium salt, it is not clear if the hydrophobic interactions overcome the electrostatic ones, and in fact the behavior of this salt is different from that shown by the other tetraalkylammonium salts.27 This could explain the deviation shown by this salt in Figure 4. Resuming, kinetic salt effects in aqueous solution for the reaction Fe(CN)5H2O3- + Co(en)2(2-pzCO2)2+ can be rationalized, on the basis of a dissociative mechanism, by considering two ways of influence of the electrolyte on the reaction: (i) an indirect one through retardation of the approaching process between two ions of opposite sign due to screening effects and (ii) a direct one through the cation-cyanide ligands electrostatic interactions which make the release of the water molecule of the ion Fe(CN)5H2O3- from the inner shell to the bulk more difficult. In the case of kinetic micellar effects in sodium dodecyl sulfate micelles it seems that the main factor operating on reactivity is the surfactant concentration influence on the approaching process, which is favored upon increasing [SDS]. The increase in surfactant concentration decreases the interfacial electrical potential of the micelles and increases the free counterion concentration, therefore, screening the charge of the micelle and making the approaching process easier. Acknowledgment. This work was financed by the Consejerı´a de Educacio´n y Ciencia de la Junta de Andalucı´a and by the D.G.C.Y.T. PB92-0677. LA960208R (27) Almgrem, M.; Swarup, S. J. Phys. Chem. 1983, 87, 876.