Solvent effect on electrode reaction kinetics of transition metal salene complexes

J. Electroanal. Chem., 179 (1984) 187-199

187

Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

S O L V E N T EFFECT O N ELECTRODE REACTION KINETICS OF T R A N S I T I O N METAL SALENE C O M P L E X E S *

A N D R Z E J K A P T U R K I E W I C Z and B A R B A R A B E H R

Institute of Physical Chemistry of the Polish Academy of Sciences, 44/52 Kasprzaka st., O1- 224 Warszawa

(Poland) (Received 21st February 1984; in revised form 25th May 1984)

ABSTRACT One-electron oxidation of Co(salen) and reduction of Co(salen) and Cu(salen) have been studied in several organic solvents containing tetraethylammonium perchlorate T E A P in various concentrations at the Pt electrode. The effect of electrolyte concentration can be interpreted using the Frumkin correction for the diffuse double layer assuming weak specific adsorption of its ions. Thus, standard rate constants k s measured in the dilute electrolyte solutions are probably least affected by the double layer. These values, together with those published earlier, were used to discuss the effect of the solvent on electrode kinetics. The kinetic data for eight solvents: acetone (AC), acetonitrile (ACN, dimethylsulfoxide (DMSO), N-methylpyrrolidone (2) (NMP), dimethylformamide (DMF), dimethylacetamide (DMA), propylene carbonate (PC) and hexamethylphosphoracidtriamide (HMPT), were discussed using the Marcus theory without success. Similarly, no correlation of k s with donor D N or acceptor AN numbers could be found. However it has been found that the experimental data can be correlated with such properties of the solvents as viscosity, v/, or dielectric relaxation time, ~-. An approximately linear relationship between Gibbs energy of activation of an electrode process E a and activation energy of the relaxation of solvent molecules Eor has been found.

INTRODUCTION

Electron transfer reactions without breaking or forming chemical bonds have been extensively studied theoretically by many authors [1]. Generally, two different approaches have been used for the description of electrode reaction kinetics. The first one, which emphasises the electron tunnelling across the electrode-solution interface, was first introduced by Gurney [2] and later developed by Gerischer [3]. However, this concept is difficult to apply to the discussion of the solvent effect. On the other hand, the theoretical treatment in which the separation of the electron and solvent components of the Hamiltonian for the electron-solvent system have been used since the work of Marcus [4]. This treatment was later developed by Levich and Dogonadze [5]. * Salene is N,N'-bis (salicylidene) ethylenediamine dianion. 0022-0728/84/$03.00

© 1984 Elsevier Sequoia S.A.

188

For comparison with experiments the Marcus theory [6] is widely used. According to the concept of Marcus the standard rate constant k ° of an adiabatic electrode process is given by k~° = Z exp - ( E a / R T )

(1)

where Z is the frequency of collision and can be estimated from

Z = ( RT/Z~rM)I/2

(2)

where M is the molar mass of the reactant. The Gibbs energy of activation E a for the charge transfer process is the sum of two terms: the work of solvent reorganization X0 and the inner reorganization energy of the reactant X i. E a = (}k0 -t- )ki)/4

(3)

The value of X 0 is solvent dependent and is estimated on the basis of a simple Born model. The inner reorganization energy Xi corresponds to changes in bonds between the atoms of reactant. This m a y be estimated from the changes of force constants and lengths of bonds in the product versus the substrate of the electrode reaction. A point unclear in the literature is the role played by the solvent. Little systematic investigation of the dependence of the kinetic parameters of electrode reactions on the nature of the solvent has been carried out. Only recently some papers appeared dealing with this problem, but none of the presented results agree with the predictions of Marcus theory. Fawcett and Barafiski [7] have studied the electro-reduction of alkali metal cations to form amalgams. They have found a correlation between the rate of the electrode process and the donor number D N [8] of the solvent. A similar result was obtained by Elzanowska et al. [9] in the case of electroreduction of Eu (III) to Eu (II). Sahami and Weaver have studied the solvent effect on one-electron reduction of some complexes of cobalt (III) [10]. In this case also the solvent effect does not agree with the predictions of the Marcus theory. The solvent effect on one-electron reduction of p-dicyanobenzene and anthracene has recently been studied by Fawcett and Jaworski [11]. They found that the standard rate constant depends on the acceptor number A N [8] of the solvent. Generally, the experimentally observed solvent effect on electrode kinetics is greater and also the sequence of solvents is, as a rule, different from that predicted by the Marcus theory. The reaction mechanism of the electroreduction a n d / o r oxidation of transition metal salene complexes in aprotic solvents normally involves one-electron transfer to form a stable anion or cation respectively (see ref. 12 and refs therein). Since the mechanism of these reactions is solvent independent [12,13] they can be used as model systems for the study of the solvent effect on electron transfer kinetics, especially in the solutions where ion-pair formation is insignificant [12,14] and the effect of the electrolyte may be related only to the structure of the electric double layer. Recently, we have studied the electrode reaction kinetics of some transition metal salene complexes in A C N and D M F containing T E A P as supporting electrolyte [15]. We have concluded that the observed increase in reaction rate with

189

increasing electrolyte concentration may be interpreted in terms of the +2 effect in the presence of weak specific adsorption of the ions of the supporting electrolyte. In this case the "apparent" rate constant k s in the dilute supporting electrolyte solutions is only slightly affected by the potential drop across the diffuse layer ~2 and it can be regarded as closest to the " t r u e " rate constant k S° cf. refs. 16, 17. The purpose of the present investigation was to study the electrode reactions of Cu(salen) and Co(salen) in various non-aqueous solvents DMA, DMSO, PC and H M P T and to use these data together with the results obtained earlier [12,15] for the discussion of solvent effects on the kinetics of the electrode reactions studied. The electrode reactions of salene complexes are rather fast, thus, according to the Marcus theory, the term describing the solvent reorganization energy X0/4 (eqn. (3)) should be a significant part of the Gibbs energy of activation E a. In addition, the differences between the solvation energies of the product and of the substrate of the electrode reaction studies are probably slightly solvent dependent [12]. Therefore the study of these reactions in different solvents should indicate which of their parameters are significant for description of the solvent effect on the electrode reaction kinetics. EXPERIMENTAL

Materials The complexes studied were prepared according to the literature Co(salen) [18] and Cu(salen) [19]. The solvents used: dimethylsulfoxide (DMSO), dimethylacetamide (DMA), propylene carbonate (PC) and hexamethylphosphoracidtriamide (HMPT) were dried and purified for electrochemical use [20]. Analytical grade (C2Hs)4NC104 (TEAP), used as supporting electrolyte, was dried at 6 0 ° C in vacuum.

Apparatus The voltammetric curves were obtained, using for low scan rates (up to 100 mV s - l ) , a measuring system constructed from an EP-20 potentiostat, an EG-20 function generator (ELPAN-Poland) and a T R P X Y recorder (SEFRAM-France) or, for rapid scan rates, (1 to 10 V s - l ) , a OP-2 oscillopolarograph (TELPOD-Poland). Measurements were performed in a conventional three-electrode cell. A brightplatinum cylinder with an active area of ca. 0.1 cm 2 was used as working electrode. The electrode was polished mechanically and further was cycled in the studied solution to obtain a steady state curve. This procedure gave reproducible results. The other electrodes were: a platinum wire as counter electrode and saturated aqueous calomel electrode (SCE) with a salt bridge (0.1 M KC1 in aqueous solutions) as reference electrode. All potentials cited in this paper were referred to the internal reference redox system bis-biphenylchromium (I)/bis-biphenylchromium (0) (BBCr+/BBCr) [21]

190 using for recalculation our own data from independent experiments. The standard redox potentials of B B C r + / B B C r were - 0.730, - 0.635, - 0.780 and - 0.578 V vs. SCE in PC, D M A , D M S O and H M P T solutions containing 0.1 M T E A P respectively.

Procedure All measurements were carried out at 25 + 0.2°C. The solutions were deoxygenated with pure argon presaturated by bubbling through the solvent used. Two concentrations of metal complexes were employed: 1 and 2 m M . Diffusion coefficients D were determined f r o m c u r r e n t - v o l t a g e curves at potential scan rates of 5 to 50 mV s-1. U n d e r these conditions the reversible behaviour was observed and diffusion coefficients could be calculated from the Randles and Sev~ik equation [22] for a one-electron process. Standard heterogenous charge-transfer rate constants k s were evaluated f r o m the observed differences in cathodic and anodic peaks potentials (100 to 180 mV) on the cyclic voltammetric curves according to the relationship given by Nicholson [23]. A detailed discussion of the method as applied to determination of k S of the order 0.01 to 0.4 cm s -1 is given by Du~ [24,25] In a few cases rate constants k s and transfer coefficients a or fl were determined from differences in the potentials of cathodic or anodic peaks and standard potentials (200 to 400 mV) using the Nicholson and Shain equation [26]. The error of the estimation of rate constants k s is smaller than _+ 20 and 30% for the slowest and fastest reactions studied respectively, and that of a and fl is ca. 10%. RESULTS

Oxidation of Co(salen) As was mentioned above, Co(salen) can be oxidized, in n o n - a q u e o u s solvents, to the Co(salen) + cation to according to eqn. (4). Co(salen) - e - ~ Co(salen) +

(4)

The rates of this reaction have been measured in D M S O and PC solutions containing T E A P as supporting electrolyte. The estimated values of standard rate constants k s and transfer coefficients/3 are presented in Table 1. The standard rate constant k S as well as the transfer coefficient fl are nearly independent of electrolyte concentration in a given solvent, as was observed in the case of D M F and A C N solutions [15].

Reduction of Co(salen) Co(salen) can also be reduced to the Co(salen)- anion according to eqn. (5) Co(salen) + e - ~ C o ( s a l e n ) -

(5)

191

The rate constants of this reaction in DMSO and PC solutions decrease with decreasing TEAP concentration much more strongly than in the case of other systems. They are presented in Table 2.

Reduction of Cu(salen) Cu(salen) can be reduced, like Co(salen), to the Cu(salen)- anion according to eqn. (6) Cu(salen) + e - ~ Cu(salen)-

(6)

The rates of Cu (salen) reduction have been measured in DMSO, DMA. H M P T and PC solutions containing TEAP as supporting electrolyte. They are presented in Table 3. The apparent rate constants are slightly dependent on TEAP concentration and their values are close to, or converge on, the value of k s for oxidation of Co(salen) in a given solvent. In some cases for Cu(salen) reduction the transfer coefficients a have also been evaluated and they are presented in Table 4. Unlike the k s values, the transfer coefficients a seem to be more strongly

TABLE 1 O x i d a t i o n of Co(salen), s t a n d a r d rate c o n s t a n t s k s, transfer coefficients fl and s t a n d a r d redox p o t e n t i a l s Eo° vs. B B C r + / B B C r at various T E A P c o n c e n t r a t i o n s c E

CE/M

DMSO

0.05 0.10 0.20 0.50

EoOx/W

a

PC

102ks/cm s- 1

/3

0.46 0.60 0.67 0.59 -}-0.717

0.20 0.18 0.19 0.22

102k 2 / c m s - 1 0.13 0.10 0.10

fl 0.30 0.38 0.38

+ 1.013

o = const. + 5 rnV). a E°x values are nearly i n d e p e n d e n t of T E A P c o n c e n t r a t i o n (Eo~

TABLE 2 R e d u c t i o n of Co(salen), s t a n d a r d rate c o n s t a n t s k s a n d s t a n d a r d r e d o x p o t e n t i a l s vs. various T E A P c o n c e n t r a t i o n s c E

CE/M

102ks//Cm S -1 DMSO

0.05 0.10 0.20 0.50 Er°ed/V a

0.62 1.20 1.20 2.10 -- 0.442

PC 0.44 0.90 1.80 -- 0.535

a ErOed are n e a r l y i n d e p e n d e n t o f T E A P c o n c e n t r a t i o n (Er°ed = const. _+ 5 mV).

BBCr+/BBCr at

192

dependent on electrolyte concentration. As was observed in D M F solutions, the transfer coefficients a increase with electrolyte concentration ',cf. ref. 15). DISCUSSION

Effect of electrolyte concentration We have reported previously [15], that the effect of electrolyte concentration on electrode kinetics of transition metal salene complexes must be related in the first place to the structure of the electric double layer. The observed changes of reaction rates with electrolyte concentration may be interpreted in terms of the ~b2 effect Frumkin correction. However in the case of solid electrodes and organic solvents this correction may present many problems because of the lack of reliable experimental data on the electric double layer. Although there are no data concerning the electric double layer on Pt in TEAP solutions in organic solvents, the potentials of zero charge can be estimated by comparison of the data for Hg electrode in organic solvents [27,28] and the data for Hg and Pt electrodes in aqueous solutions, using Frumkin's considerations [29]. The

TABLE 3 R e d u c t i o n of Cu(salen), s t a n d a r d rate c o n s t a n t s k~ at various T E A P c o n c e n t r a t i o n s c E a n d s t a n d a r d redox p o t e n t i a l s vs. B B C r + / B B C r

CE/M

102ks/cm s 1 DMSO

0.05 0.10 0.20 0.50 Ereo d / / V

0.56 0.75 0,75 0,75 - 0.406

a

HMPT 0.11 0.11 0.10 - 0.505

DMA 0.53 0.54 0.50 1.10 - 0.583

PC 0.28 0.33 0.37 - 0.498

a ErOa are nearly i n d e p e n d e n t of T E A P c o n c e n t r a t i o n ( E ° a = const. __+_5 mV).

TABLE 4 R e d u c t i o n of Cu(salen), transfer coefficients a at various T E A P c o n c e n t r a t i o n s ¢E CE/M

ot DMSO

HMPT

DMA

PC

0.05 0.10 0.20 0.50

0.28 0.32 0.34 0.44

0.23 0.27 0.55

0.40

0.28 0.34

193 potentials of zero charge of Pt electrodes should lie in the region + 0.7 to + 1.0 V vs. B B C r + / B B C r internal reference redox system bisbiphenyl chromium(1)/bisbiphenyl chromium(0) in all organic solvents studied. The standard potentials of Co(salen) oxidation in PC and D M S O are near the estimated potentials of zero charge. Similarly, the standard potentials of Co(salen) oxidation in other organic solvents AC, ACN, DMA, DMF, N M P and H M P T also lie in this region (cf. data from refs. 12 and 15). This could be the reason why the rate of this reaction in a given solvent is almost independent of electrolyte concentration. Therefore the rate constants of Co(salen) oxidation are probably not affected by the potential drop in the diffuse double layer and they may be regarded as the " t r u e " rate constants. The " a p p a r e n t " rate constant of the reduction of Co(salen) and Cu(salen) increases with electrolyte concentration in most cases, as one would expect on the basis of the electric double layer effect for an uncharged reactant being reduced at a negatively charged electrode. F r o m our earlier theoretical considerations [15] it might be expected that the rate constants of electrode reactions of transition metal salene complexes would be slightly dependent on the nature of the central metal ion. It seems that the effect of the diffuse layer is the smaller the lower the T E A P concentration because the rate constants of the reduction of Co(salen) and Cu(salen) converge to values similar to k s for the oxidation of Co(salen) in a given solvent with decrease in electrolyte concentration. This agrees with the behaviour expected in the case of the presence of weak specific adsorption of ions of the supporting electrolyte. It is very probable in view of the literature data which report the specific adsorption of the (C2H5)4 N+ cation on a Hg electrode from aqueous [30] and organic solutions [31,32]. The reduction of Co(salen) is accelerated much more by the electrolyte concentration than the reduction of Cu(salen) in both D M S O and PC solutions. Similar behaviour has been observed in the case of A C N and D M F solutions and this has been interpreted in terms of possible interactions between the Co(salen)- anion and the adsorbed cations (C2H5)4 N+ [15]. The present results can also be explained using this concept. The consequence of our conclusions about weak specific adsorption of electrolyte ions is also that O~2/OE- 0 ( E - - e l e c t r o d e potential) and that this factor cannot influence the value of the transfer coefficient. However ct can depend strongly on the difference between the reaction plane x r and the outer Helmholtz plane x2; the bigger the difference x 2 - x r the lower is the transfer coefficient [33-35]. In most cases the transfer coefficient increases with the electrolyte concentration. It m a y suggest that the reaction plane xr shifts further from the electrode if the specifically adsorbed charge increases. F r o m a model of the M(salen) molecule we can estimate that in the flat position it can easily be accommodated within the Helmholtz layer (x r < x2). Only a small deviation from the flat position could justify the observed changes of the transfer coefficients. For the oxidation of Co(salen), which probably takes place at an almost uncharged electrode, this effect is not observed which is understandable. The transfer coefficients are also dependent on the nature of the solvent. This can

194 be explained b y the differences in the thickness of the H e l m h o l t z layer a n d b y the differences of the adsorbed charges.

Effect of the solvent T h e most i m p o r t a n t c o n c l u s i o n from the discussion presented above using the a s s u m p t i o n of weak specific a d s o r p t i o n of s u p p o r t i n g electrolyte ions is that the rate c o n s t a n t in dilute T E A P solutions is closest to the " t r u e " rate c o n s t a n t ks° a n d these values should be used for the study of the solvent effect o n electrode kinetics. O n the other h a n d only the Co(salen) oxidation a n d the Cu(salen) r e d u c t i o n m a y be considered because the rate constants of the Co(salen) r e d u c t i o n can be affected b y interactions in the outer Helmholtz layer also in the dilute s u p p o r t i n g electrolyte solutions. T h e effect of electrolyte c o n c e n t r a t i o n on the kinetics of r e d u c t i o n a n d o x i d a t i o n of Ni(salen) has b e e n studied previously in A C N a n d D M F solutions [15]. The observed increase in reaction rate with increasing electrolyte c o n c e n t r a t i o n has also been interpreted in terms of the simple qJ2 effect in the presence of weak specific a d s o r p t i o n of its ions. The rate constants of Ni(salen) oxidation, as well as of Ni(salen), r e d u c t i o n are n o t very different from the c o r r e s p o n d i n g values of k S for Co(salen) oxidation in a given solvent c o n t a i n i n g 0.1 M T E A P as s u p p o r t i n g electrolyte [12]. T h u s these values can also be used for the study of the solvent effect o n electrode kinetics. I n T a b l e 5 the rate constants of all four reactions are presented. O n the basis of the d a t a from T a b l e 5, it is readily a p p a r e n t that the rate c o n s t a n t varies considerably with solvent n a t u r e a n d less with the central metal a t o m in a complex molecule.

TABLE 5 Standard rate constants k s of electrode processes of transition metal complexes in organic solvents containing 0.1 M TEAP Solvent

102ks/cm s- 1 Oxidation Co(salen

AC ACN DMF DMA DMSO MMP PC HMPT a Ref. 12. b Ref. 15. c This work.

2.90 a 2.20 b 0.60 b 1.40 c 0.60 ~ 0.44 a 0.13 c 0.13 ~

Reduction Ni(salen)

1.87 b 0.90 a 0.34 a 0.43 a

Cu(salen)

0.70 b 0.52 ~ 0.73 c 0.33 ~ 0.11 c

Ni(salen) 3.07 a 3.23 b 1.90 b 1.52 a 1.22 a 0.55 a 0.43 ~ 0.40 a

195

The reaction rates could be considered using Marcus theory in which the estimate of Gibbs energy of activation is based on a continuum model for the solvent. It corresponds to the difference in the reciprocals of the static dielectric constant c and optical dielectric constant c 0, in a given solvent, multiplied by the factor, characterizing the reactant, which depends also on its distance from the electrode and its orientation in the double layer [36]. Assuming that this factor, in the cases of the reactions studied, does not depend on the solvent, the changes of activation energy should correspond to the changes in the values of the terms 1 / ~ - 1/n 2 given in Table 6.These changes do not exceed +10% in X0 and correspondingly less with respect to the total energy of activation E a. The differences found experimentally are much greater. Also the sequence of experimentally obtained rate constants does not agree with the sequence of the term 1 / c - 1/n 2. Thus the simple Marcus theory which allows the outer solvation shell effects to be estimated does not describe correctly the solvent effect in the presented cases. This conclusion was also drawn by the authors cited previously [7,9-11]. Also, no simple correlation with donor D N or acceptor A N numbers, which should reflect the inner solvation shell effects, could be found, the sequence of rate constants being the same for both reduction and oxidation processes (cf. Tables 5 and 6). The regularity which has been found is the same sequence of rate constants and the solvent viscosity ~ (cf. Tables 5 and 6) The lower ~ the more rapid is the electrode reaction. For electrolyte concentrations differing from 0.1 M this sequence is also preserved for a given electrode reaction. Also the sequence of rate constants k s is the same as the dielectric relaxation time • of solvent molecules. The lower ~- the higher is k s. The times of relaxation r, found experimentally, are also presented in the Table 6. Theoretically the time of relaxation of solvent molecules can be calculated using

TABLE 6 P a r a m e t e r s of the solvent studied, 1 / c r e l a x a t i o n time ~-

1 / n 2 ~, d o n o r n u m b e r D N , acceptor n u m b e r A N , viscosity ~,

Solvent

1/c -1/n 2

DN

AN

1 0 2 ~ / g cm -1

AC ACN DMF DMA DMSO MMP PC HMPT

- 0.493 - 0.526 -0.462 - 0.456 - 0.437 -0.438 -0.481 - 0.437

17.0 14.1 26.6 27.8 29.8 27.3 15.1 38.8

12.5 18.9 16.0 13.6 19.3 13.3 18.3 10.6

0.316 0.345 0.766 0.927 1.980 1.663 2.530 3.470

S -1

1012"r/S

Ref. b

2.5 3.8 11.0 14.0 19.6 30.0 43.0 80.0

38 39 40 40 40 _ c 41 40

a T e r m from the M a r c u s theory, n = r e f r a c t i o n coefficient. b References for the values of r e l a x a t i o n times. c Value of ~" c a l c u l a t e d a c c o r d i n g to eqn. (7). Values of c, n a n d ~ f r o m ref. 42.

196

the Debye equation [37]

"r = 3 V ~ / k T

(7)

in which V is the volume of the solvent molecule. Thus in view of the above equation the correlations between k~ and ~/ or k~ and ~- are equivalent, for the solvents discussed in this paper where V increases monotonically with ,/. In consequence one could expect a correlation between the Gibbs energy of activation of the electrode process E a (cf. eqn. (1)) with the energy of activation of solvent molecules reorientation Eor which is related to ~- according to [37] (1/~-) = ( k T / h ) exp( - E o r / k T )

(8)

The relationship between E a and Eor is presented in Fig. 1. A roughly linear relationship between E a and Eor with the slope equal to unity is observed. The value E a for Eor = 0 should correspond to the energy of activation E a in the sense of the Marcus theory. This experimental result confirms our earlier theoretical considerations according to which Xi should not depend greatly on the solvent and on the central ion in the complex. The deviations from the statistical straight line can be ascribed to the accuracy both in the theoretical considerations and the experimental data. It is obvious that the correlation of this type can be expected only in the cases where the value Xi is much less dependent on the solvent than the value of Eor. Solvent effects were much more discussed in the literature in the case of

Ea / k c a l tool -~ O

10. 9.

~_ ~_

C3

,/5;

I

.

L)

OgAO i

7. ~ ~ o /

Co(salen) - e Isalen) e - Ni (salen) + e

v - Ni

Eor./kcal tool -~

4 0

Fig. 1. The relationship between the energy of activation of the electrode processes E a and the energy of the solvent reorganisation Eor (1 cal --- 4.184J).

197 h o m o g e n e o u s r e a c t i o n s kinetics t h a n of e l e c t r o d e reactions. T h e effect of viscosity as a factor d e p e n d i n g on solvent structure was m e n t i o n e d b y various a u t h o r s [43]. Especially interesting for electrochemistry are the w o r k s c o n c e r n i n g h o m o g e n e o u s charge transfer reactions. K o s o w e r et al. [44,45] have p r e s e n t e d e x p e r i m e n t a l results for a n u m b e r of reactions of h o m o g e n e o u s charge r e d i s t r i b u t i o n in o r g a n i c m o l e cules in various solvents. T h e y have f o u n d a c o r r e l a t i o n b e t w e e n rate c o n s t a n t s a n d dielectric r e l a x a t i o n times of the solvents. R e c e n t l y this p r o b l e m was s t u d i e d theoretically b y C a l e f a n d W o l y n e s [46,47]. T h e y have d e r i v e d the r e l a t i o n s h i p b e t w e e n rate c o n s t a n t s a n d dielectric r e l a t i o n times. This theoretical t r e a t m e n t explains the e x p e r i m e n t a l results for h o m o g e n e o u s charge transfer a n d can, p r o b a bly, also be a p p l i e d to heterogeneous e l e c t r o d e reactions. It is interesting to c o m p a r e the literature d a t a in terms of the c o r r e l a t i o n f o u n d in this work. T h e kinetic d a t a for the r e a c t i o n s Eu ( I I I ) / E u (II) [9], L i + / L i (Hg) [7] as well as C o ( e n ) 3 + / / C o ( e n ) 3 2 + [10] c a n n o t b e i n t e r p r e t e d in the s a m e w a y as the r e a c t i o n of salene complexes. T h e strong i n t e r a c t i o n b e t w e e n r e a c t a n t a n d solvent molecules can lead to significant changes in X i in d e p e n d e n c e o n the solvent. However, there exist some literature d a t a which can be discussed in terms of the r e l a x a t i o n times ~- of the solvent molecules. This is the case of o n e - e l e c t r o n r e d u c t i o n of p - d i c y a n o b e n z e n e [11] a n d o x i d a t i o n of ferrocene [49,50]. In T a b l e 7 these d a t a are p r e s e n t e d . T h e c o r r e l a t i o n b e t w e e n ~- a n d k ° can b e r e g a r d e d as satisfactory. In the case of p - d i c y a n o b e n z e n e the c o r r e l a t i o n of k ° with A N f o u n d b y F a w c e t t a n d Jaworski should b e checked b y a d d i t i o n a l m e a s u r e m e n t s for a solvent which d r a s t i c a l l y deviates f r o m the sequence of ~- as c o m p a r e d with an eg. H M P T .

TABLE 7 Solvent effect on the standard rate constants k s of the reduction of C6H4(CN)2 [11] and the oxidation of (C2Hs)2Fe [49,50], comparison with the relaxation time of the solvents ~'. Solvent

1012'r/s

k J c m s- 1 Reduction of C6Hs(CN)2

ACN DMF DMA DMSO PC EtOH NMF S

3.8 11.0 14.0 19.6 30.0 100 a 123 b 130 ~

3.52 4.39 2.17 2.45

Oxidation of (CsHs)2Fe 4.4 × 10- 2 3.3×10 2 3.8 × 10 -2 1.6 × 10 -2

0.60

" Value for ethanol from ref. 48. b Value for N-methylformamide from ref. 40. c Value for sulpholane calculated from eqn. (7).

0.3 × 10-2

198 CONCLUSIONS F o r the discussion of the solvent effect o n kinetics of electrode reactions the influence of the electric double layer structure should be taken into account. I n the case of the systems studied the a s s u m p t i o n of weak specific a d s o r p t i o n of ions of the s u p p o r t i n g electrolyte enables the i n t e r p r e t a t i o n of the observed effects of electrolyte c o n c e n t r a t i o n on k S a n d c~ or ft. The correction for the diffuse d o u b l e layer is p r o b a b l y the smallest in the case of the dilute electrolyte solutions. T h u s the values of k S in these solutions were used for the discussion of the solvent effect o n the electrode reaction kinetics. The data for eight solvents were c o m p a r e d with the prediction of the Marcus theory. It can be c o n c l u d e d that this theory c a n n o t describe correctly the solvent effect as was reported b y other authors [7,9-11]. Also n o correlations could be f o u n d with A N a n d D N . The correlation which we could find was that with the viscosity ~ a n d / o r dielectric relaxation time ~-. The reason for this can be rather slow r e o r i e n t a t i o n of solvent molecules which gives a significant c o n t r i b u t i o n to the overall a c t i v a t i o n energy of the electrode process. T h e fact of the correlation of the kinetics with the relaxation times ~- of pure solvents suggests rather weak i n t e r a c t i o n b e t w e e n the r e a c t a n t a n d the m e d i u m . The concept presented above can also be used for i n t e r p r e t a t i o n of some literature data [11,49,50]. ACKNOWLEDGEMENT The work was carried out within Research Project 03.10. REFERENCES 1 J.O'M. Bockris and S.U. Khan, Quantum Electrochemistry, Plenum Press, New York, 1979. 2 R.W. Gurney, Proc. R. Soc. (London), A134 (1931) 137. 3 H. Gerischer, Ber. Bunsenges. Phys. Chem., 77 (1971) 771; H. Gerischer in P. Delahay (Ed.), Advances in Electrochemistry and Electrochemical Engineering, Vol 1, Wiley-Interscience, New York, 1961, p. 31. 4 R.A. Marcus, J. Chem. Phys., 24 (1956) 966. R.A. Marcus, J. Phys. Chem., 67 (1963) 853, 2883. R.A. Marcus, J. Chem. Phys., 43 (1965) 679. 5 R.R. Dogonadze in N.S. Hush (Ed.), Reactions of Molecules at Electrode, Ch. 3, Wiley-Interscience, New York, 1971; V.G. Levich in P. Delahay (Ed.), Advances in Electrochemistry and Electrochemical Engineering Vol 4, Wiley-Interscience, New York, 1966, p. 243. 6 R.A. Marcus, Annu. Rev. Phys. Chem., 15 (1964) 155. 7 A. Baraflski and W.R. Fawcett, J. Electroanal. Chem., 94 (1978) 237. 8 V. Gutmann, The Donor-Acceptor Approach to Molecular Interaction, Plenum Press, New York, 1979. 9 H. El~'anowska,Z. Borkowska and Z. Galus, J. Electroanal. Chem., 157 (1983) 251. 10 S. Sahami and M.J. Weaver, J. Electroanal. Chem., 124 (1981) 35. 11 W.R. Fawcett and J.S. Jaworski, J. Phys. Chem., 87 (1983) 2972. 12 A. Kapturkiewicz and B. Behr, lnorg. Chim. Acta, 69 (1983) 247.

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