Electron transfer between immiscible solutions

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J. Electroanal. Chem., 244 (1988) 21-37 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

ELECTRON TRANSFER BETWEEN IMMISCIBLE SOLUTIONS THE HEXACYANOFERRATE-LUTETIUM BIPHT’HALOCYANINE SYSTEM

G. GEBLEWICZ Chemistry Department,

The Universi&

Southampton

SO9 5NH (Great Britain)

D.J. SCHIFFRIN Wolfson Centre for Electrochemical Science, Chemistry Department, Southampton SO9 5NH (Great Britain)

The University,

(Received 31st August 1987; in revised form 9th December 1987)

ABSTRACT Electron transfer between aqueous and immiscible non-aqueous solutions of redox species has been demonstrated. A macrocyclic Lu complex has been used in the organic phase. Both the complex and its oxidation product are completely insoluble in water. It is shown that it is possible to predict the transfer potentials from simple voltammettic experiments. The kinetic results are analysed using Marcus’ theory and it is concluded that the redox centres are further apart than is the case in studies at the metal/solution interface.

INTRODUCTION

There has been a great deal of interest recently in the study of heterogeneous electron transfers occurring between r&lox centres present in two different media [l-5]. Examples of systems where this type of transfer is important are the electron transfer chain in mitochondria, where NADH is oxidised by dioxygen through a series of intermediate electron transfers between membrane constituents and redox enzymes. Immiscible electrolyte solutions represent simple models for these studies, since the interfacial electrical potential can be easily controlled externally and therefore, charge transfers between redox centres present in different media can be studied conveniently. The only system that has been studied so far with this approach is ferrocene in nitrobenzene in contact with the Fe(CN)z-13- redox couple in water [4]. However, the use of ferrocene as the organic redox system presents some problems due to the possible interfacial transfer of the ferricenium ion in the experimental conditions 0022-0728/88/$03.50

Q 1988 Elsevier Sequoia S.A.

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used, and hence coupling of electron and ion transfer may occur in this case. Recent work appears to indicate that the transfer potential of ferricenium occurs at potentials less positive than electron transfer, implying that these processes occur independently of each other [5]. In order to study the kinetics of true heterogeneous electron transfers, it is desirable therefore to find organic redox couples that are sufficiently hydrophobic so that electron transfer is the only process observed. In the present study, lutetium biphthalocyanine (Lu(PC),) has been chosen as the organic redox system for several reasons: (a) the very hydrophobic nature of this high molar mass compound excludes the possibility of its partition to the water phase from 1,Zdichloroethane (DCE); (b) the extensive macrocyclic system present in this type of molecule results in the existence of many redox states involving electron transfers to the rings, thus allowing for a more flexible choice of the aqueous redox couple and, (c) N, macrocyclic metal chelates are of importance in biological systems as they occur in the cytochromes associated with the respiratory electron transfer chain, as energy transfer intermediates in photosynthesis, in hydroxylating enzymes, etc.

EXPERIMENTAL

A four electrode potentiostat with ohmic compensation was used. The potentiostat and cell were placed inside a Faraday cage. This was found important for achieving a good degree of ohmic compensation. The currents measured were always very low, of less than 1 PA, and therefore the results are free from uncompensated resistance errors. The cell was the same as used before [6], with an interfacial area of 0.385 cm’. The ac impedance experiments to measure the rate constant of the hexacyanoferrate couple were carried out with a Solatron 1250 Frequency Response Analyser controlled with a BBC B + microcomputer. The electrolytes were chosen to obtain a very wide polarisation window. The aqueous electrolyte was 1.5 M Li,SO, (AnalaR); as shown previously, this high concentration of Li,SO, results in the salting out of the organic electrolyte with the corresponding increase in the available limits of polarisation. The organic electrolyte was lo- 3 M tetraphenylarsonium tetrakis(4-chlorotetraphenyl)borate (TPAs TPBCl). The latter salt was prepared by precipitation in water of TPAsCl and NaTPBCl (Fluka) followed by recrystallization from acetone. The preparation and purification of the Lu(PC), have been described elsewhere [7]. AnalaR grade K,[Fe(CN),], K,[Fe(CN),] and Aldrich Gold Label DCE were used without further purification. The concentration of the aqueous redox couple was chosen much higher than that of the organic phase in order to simplify the diffusional problem. In consequence, only the diffusion of Lu(PC), need be considered, and the analysis of the voltammetric results can be carried out using the classical theory for metal electrodes. The overall reaction should become, in principle, first order with respect to the organic redox couple.

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The redox properties of Lu(PC), in DCE on a metal electrode were studied on a 10 pm platinum microelectrode using a Hi-Tek PPRl waveform generator as a potential source and a Keithley 427 current amplifier to measure the currents. The potentials were referred to the aqueous phase by introducing a well defined organic/aqueous phase reference junction, having the TPAs+ ion present in both phases. For the liquid/liquid experiments the potentials refer to the cell: CELL I

Ag ( AgCl ( 1O-3 M TPAsCl(w)

( 1O-3 M

( u 1 1.5 M Li,SO,(w)

1 Pt

TPAsTPBCl(o)

+O.l M K,[Fe(CN),]

+x M Lu(PC),

+O.Ol M K,[Fe(CN),]

0

I E

where u represents the interface under study. For the microelectrode measured potential corresponds to:

results, the

CELL 2

Ag

AgCl

1O-3 M TPAsCl(w)

1O-3 M TPAsTPBCl(o)

Pt

+x M Lu(PC) 0

&

Thus, the convention used in the present work is to measure the potential of the aqueous with respect to the organic phase. Positive currents correspond to electron transfer from 1,ZDCE to water. The concentration of biphthalocyanine in DCE was calculated from the absorbance of the solutions, which was measured with a Pye-Unicam 8800 spectrophotometer. Since the hexacyanoferrate solutions tended to decompose with time, they were freshly prepared before each experiment. All the experiments were carried out at room temperature (22 f 2 o C). RESULTS

UV-visible

AND

DISCUSSION

spectrum

The UV-visible spectrum of Lu(PC), in 1,2-DCE is shown in Fig. 1. The concentration of the different solutions used was determined from their absorbance at 660 nm. The spectrum observed is almost the same as that reported by Corker et al. [8] from DMF. The spectrum of Lu(PC), was unaffected by dissolved TPAsTPBCl or when these solutions were shaken either with the aqueous supporting electrolyte or with the hexacyanoferrate solution. These experiments were carried out in order to ascertain that no chemical degradation occurred in the media studied, and that electron transfer could not occur between the oxidising aqueous solution and the

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A 2.0 ..

300

400

500

600

700

p.1nm

Fig. 1. UV-visible spectrum of a 2.5 x 10e5 M solutionof Lu(PC), in 1,2-DCE.

non-aqueous phase if the Galvani potential difference was not fixed by the externally applied potential (see later). Voltammetry

with a metallic electrode

Figure 2 shows the linear sweep voltammetry of Lu(PC), on the platinum microelectrode, which corresponds to a reversible one electron transfer process (see inset in Fig. 2). Two reversible oxidation and reduction waves with half wave potentials (E,,,) of 0.538 and 0.025 V respectively can be seen. The oxidation and reduction of Lu(PC), has been observed previously both in the solid phase [7] and from dichloromethane (DCM) solutions [9]. It has been proposed that the red oxidation product of the green form of Lu(PC), is a radical dication [8] and from the cyclic voltammetry results of L’Her et al. [9], this species is stable in DCM. Similar conclusions regarding the stability of the oxidation product were observed by Corker et al. [8], who studied its ESR spectra in DMF. In the solid state, the oxidised form of Lu(PC), is stable for long periods of time, even in the presence of water, as shown by the long term stability of thin film electrochromic devices [7]. From this point of view, lutetium biphthalocyanine appears to be a suitable hydrophobic electron donor for studies of electron transfer across interfaces of immiscible electrolytes. The present results are very close to those in dichloromethane [9], since the differences in E1,2 between oxidation and reduction are 0.51 and 0.48 V, respectively. The small differences can be due to the different electrostatic contributions to the solvation energy and to ion pairing effects. For a microelectrode, the limiting current, 1, is given by [lo]: I, = 4nFDcr where n is the number of electrons,

(1) F is the Faraday constant,

D the diffusion

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lOCJ+) 11)

a0

-1.0

\

I I 0.6

E/V

I

0.4

a5 0.6

IOpA

I

I:;; b)

CL2

Fig. 2. (a) Linear sweep voltammetry of 5 X lo-’ M Lu(PC), + 10e3 M TPAsTPBCl in 1,2-DCE at a Pt microelectrode. The potentials refer to Cell 2. (2) Semilogarithmic plot to test for a 1 e- reversible process.

coefficient, c the concentration of electroactive species and r is the microelectrode radius. From the results shown in Fig. 2, D = 1.1 X 10e6 cm* s-l. Electron transfer across the immiscible electrolyte interface Cyclic voltammetry Figure 3 shows the voltammetric response of the immiscible redox system studied. Although other cases of electron transfer at liquid/liquid interfaces have been reported previously [3,5], due to the high insolubility of the biphthalocyanines in water, the present results can be regarded as the first evidence for electrochemical electron transfers across inhomogeneous media in which tunnelling across redox centres present on opposite sides of the interfacial region is extremely unlikely to be coupled to ionic transfers. In contrast with ferrocene, the partition coefficient of the biphthalocyanine is extremely low due to the presence of very large macrocyclic rings in these molecules, thus precluding the possibility of coupled interfacial electron and ion transfers. Ferrocene dissolved in 1,ZDCE partitions in water, whereas the complex studied in the present work does not. The diffusion coefficient of Lu(PC), in I,2-DCE The peak currents corrected for charging currents showed a linear dependence with the square root of the sweep rate (v), as shown in Fig. 4. From eqn. (2) [ll]: I P = 2 .69 x lo5 D1’*cv”*

(4

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Fig. 3. Cyclic voltammetry for the interfacial electron transfer between Lu(PC), and hexacyanofemate. The potentials correspond to Cell 1. Sweep rates: (1) 0.005; (2) 0.01; (3) 0.02; (4) 0.03; (5) 0.05 V s-l. Concentration of Lu(PC), = 5 X 10e5 M.

the average value of the diffusion coefficient calculated assuming a reversible one-electron reaction was 0.95 x 10m6 cm2 s-l, in reasonable agreement with the value obtained with the Pt microelectrode. Ir, is the peak current density. The reaction appears to be quasi-reversible as shown by the dependence of the peak potential separation on sweep rate. Also, as expected from eqn. (2), d1,/dd’2 should be proportional to the concentration of Lu(PC),. Figure 5 shows that a linear relationship is observed, corresponding to an average diffusion coefficient of 0.98 X 10m6 cm2 s-l, in good agreement with the previous estimate. It is interesting to compare the diffusion coefficient with the hydrodynamic radius of the diffusing species. From the Einstein-Stokes relationship [12]: rp = kT/(6aqD)

(3)

a diffusion radius rp can be calculated. Here k is Boltzmann’s constant, T the absolute temperature and 7 the viscosity. Taking 9 = 8 X 10e4 Pa s for 1,ZDCE at 20 o C [13], values of rp between 2.7 and 2.4 nm are obtained. The maximum radius of a single phthalocyanine ring is - 0.7 mn. Considering that the ionic radius of

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20

40

60

80

C/FM

Fig. 4. Dependence of the peak current on the sweep rate for electron transfer at the water/l,Z-DCE interface for different concentrations of Lu(PC),. (0) 1.5 X 10m5; (0) 2.86 X lo-‘; (0) 5 X 10v5 M. The concentration of hexacyanoferrate is the same as in Cell 1. Fig. 5. Dependence of dZ,/du’j2

for electron transfer on the concentration of Lu(PC),.

Lu3+ is 0.095 nm [14], the expected hydrodynamic radius of less than 0.9 nm. The diffusion experiments give therefore, a low for the known molecular dimensions. Perhaps the origin in the use of an equation for a continuous medium, such relationship.

the molecule should be value of D which is too of this discrepancy lies as the Einstein-Stokes

The interfacial electron transfer potential The general equilibrium between two immiscible redox couples is: Ox@‘) + Red@) + Red@‘) + Ox@)

(4)

where Ox and Red are the oxidised and reduced components of the redox couples present in the aqueous (w) and organic (0) phases respectively. To simplify the treatment, a 1 e- reaction will be considered. The equilibrium condition (4) leads to [15]:

(5) where A:+ is the difference in Galvani potentials $I(“) - #“, AG& is the standard chemical Gibbs energy of reaction (4) and a, are the corresponding activities. Although eqn. (5) is a non-thermodynamic relationship involving individual ionic properties, it is very useful for the prediction of the potential range where electron

34

transfer between redox couples present in immiscible phases will occur. The possibility of making these calculations relies on the assumption of the equality of the values of the Gibbs energies of transfer for tetraphenylborate and tetraphenylarsonium (TPB- and TPAs+) from water to an organic solvent [16]. With this assumption and a modified Debye-Htickel equation, the calculated value of AI+ can easily be related to the potential of Cell 1 and hence to a practical potential scale. AG & is related to the standard potentials of the two redox couples in their respective phases measured with respect to a common aqueous reference electrode [4]. The value of E,,, measured with Cell 2 will be taken as the standard potential of the Lu(PC)g/‘+ couple (the green form is a radical cation [S]). From an analysis of the potentials of Cells 1 and 2, it can easily be shown that: Q2

(Cell 1) = E1,2 (Cell 2)

(6)

i.e., the half wave potentials for the two experiments should be identical. Within experimental error, this was found to be the case (Figs. 2 and 3). This is a very important result since the equality of the two half wave potentials can serve as the basis for predicting the potential of electron transfer between redox centres present in different media, from the study of the behaviour of the individual redox couples in their respective phases. The physical meaning of eqn. (6) is that the presence of a concentrated redox system in one of the adjoining phases containing both the reduced and oxidised components of the couple has the same properties as a metallic electrode; indeed, this is to be expected as the Fermi level of the aqueous phase will remain constant throughout the experiment and the solution will therefore behave as an extension of the metal.

The rate constant of electron transfer Since the electron transfer appears to be quasi-reversible,

the rate constant was calculated by the method of Nicholson [17] and the calculated value of k’ o ; the apparent rate constant was 0.9 X 10m3 cm s-l. The electron transfer reactions of the two redox couples studied are generally regarded as outer sphere electron transfers. The solvent reorganization energies of both hexacyanoferrate and the biphthalocyanine must be taken into account in any theoretical calculation of the rate constant. A measure of this quantity is the rate constant for electron transfer at a metallic electrode. No kinetic data are available for Lu(PC), but it is known that electron transfer rates to conjugated macrocyclic compounds are very fast, with heterogeneous rates constants greater than 0.04 cm S -’ for porphyrins [18]. These fast rates are a consequence of the low reorganization energies resulting from electron delocalisation in the macrocyclic rings, and rate constants greater than the above are expected for the biphthalocyanines. The corresponding rate constant for the hexacyanoferrate couple is also very large. Peter et al. [19] showed that k’” is strongly dependent on the cation of the base electrolyte used, and for a 3 M LiNO, solution, a value of k'O of 0.02 cm s-l was obtained on a gold electrode. In the present work, k’” for the 1.5 M Li,SO,

35

solution was measured by ac impedance on a gold electrode and a rate constant of 0.035 cm s-r was found, in reasonable agreement with Peter’s results. In spite of the difficulties of measuring a rate constant accurately from cyclic voltammetry at liquid/liquid interfaces, these results and the expected value of k’” for Lu(PC), for metallic electrodes are over one order of magnitude greater than the measured rate constant of electron exchange at the liquid/liquid interface. In order to carry out a comparison, the rate constant measured at the liquid/liquid interface must be corrected for diffuse double layer effects since these could be significant in the very dilute organic solutions employed. The use of the Frumkin correction is fraught with difficulties since at present we do not have a reliable picture of the double layer or of the potential distribution at the liquid/liquid interface. If the interface can be regarded as a mixed layer rather than two “inner layers” of pure solvent molecules [20], all the applied potential would appear across the two diffuse double layers present on both sides of the interface. In common with previous work [4,21], a primitive electrostatic model can be used for calculating diffuse layer potentials, but due to the differences in electrolyte concentrations and dielectric constants between the two phases studied, nearly all the potential drop will appear in the organic phase. The Galvani potential difference at El,* can be calculated from the individual contributions in Cell 1: El/2

= (EFwIj,Fe(IIIj

-

EAg,Agd

+ K&,2

-

&hAs+

(7)

are the potentials of the silver/silver chloride and EAp/AgCl and EFe(II)/Fe(III) hexacyanoferrate couples with respect to the standard hydrogen electrode, A’&rAs+ is the Galvani potential difference of the junction between the reference electrode and the solution under study, and A’&,2 is the Galvani potential drop across the immiscible electrolyte system under study at the standard potential of the Lu(Pc):‘2+ couple. The activity coefficient of Cl- ion in the TPAsCl solution was calculated from the extended Debye-Htickel theory taking a value of the ion size parameter of 0.6 run [12], from which ycl- = 0.966. + was calculated from: 4hPAS

where

(8) has been given as -0.364 V [22] and -0.338 V [23]. The activity FOOTPAD+ coefficient of TPAs+ in the organic solution was calculated from the extended form of the Debye-Htickel theory using a value of the ion size parameter of 0.66 run, based on the values given by Abraham and de Namor [23]. In the absence of thermodynamic information, the association constant of TPAsTPBCl was taken as equal to that of TPAsTPB (600 M-l [23]). Since this term represents a small contribution to A:+rrAs+ due to the very dilute solutions employed in the present work, the small difference in ion pair contact distance will not affect the calculated potential. A value of yTpAs+= 0.535 and a degree of dissociation of 0.870 was calculated for the 10m3 M solution employed. The potential of the Fe(CN)d-/3couple in 1.5 M L&SO, was measured against a saturated calomel electrode using a

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giving

Pt el=trode, &e(IIj,Fe(IIIj- 0.565V vs. SHE. From the above, the Galvani potential difference at ~!?i,~ is A;&,2 = - 9 or + 16 mV according to which value of + is chosen. It must be concluded that the Frumkin correction to the A%oTpAs measured values of k’ o will be small. The lower value for k’ ’ observed at the liquid/liquid interface than that measured on metal electrodes can be due mainly to two reasons: (a) a change in the solvent reorganization energy due to the replacement of the metal in half the space by a solvent and (b) a greater distance between redox centres at the immiscible electrolyte interface as compared with the sum of distances of closest approach to metal electrodes. The kinetic equation for heterogeneous electron transfer is [24]: k’=XZexp(-g) where x is the transmission coefficient taken as 1 for adiabatic reactions, the pre-exponential term 2 has a value of lo4 cm s-l [24] and AG” # is the standard activation energy of the reaction. The only component of AG o # that concerns us here is the solvent dependent term, AG,I #, which according to Marcus’s theory is given by [25]: AG;*

=A,/4

00)

with the “outer” reorganization energy h, given by:

(11) where a is the ionic radius, R is twice the distance from the ion to the electrode, Do and D, are the optical and static dielectric constants, respectively. The equivalent relationship for an immiscible electrolyte interface when the two reactants are located at an equal distance I from a hypothetical interfacial plane is [26]:

(12) where b is the radius of the redox species in the organic phase. As can be seen by comparing eqns. (9) and (lo), the solvent reorganisation energies for the liquid/liquid case are the same as those for the individual couples at metal electrodes, and the activation energy is the sum of the individual activation energies provided R and I are the same. The rate constant determined by cyclic voltammetry corresponds to a condition of equimolar surface concentration of Lu(PC): and Lu(PC);‘. Since the concentrations of the fast aqueous redox couple is more than three orders of magnitude greater than that of the organic couple, the value of the rate constant calculated from cyclic voltammetry must be related only to the activation energy of the Lu(PC), couple. In this respect, the highly concentrated aqueous redox couple

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behaves as a metallic electrode, as the number of available energy levels for electron exchange is very large compared with those in the organic phase. However, k’” obtained from these experiments is very low compared with the expected electron transfer rate at metallic electrodes. It is proposed, therefore, that the observed irreversibility is related to the greater separation between redox centres in the liquid/liquid case compared with the metal/solution interphase. From eqns. (9) and (10) it can be seen that this would result in higher values of the reorganization energy and hence, in lower rate constants. More accurate determinations of rate constants are required to test the proposed comparison between metal and redox electrodes. ACKNOWLEDGEMENTS

The support by SERC (U.K.) is gratefully acknowledged. Part of the equipment was purchased with funds from the University Grants Committee Special Award to the Electrochemistry Group at Southampton, 1985. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26

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