Electrochemical behavior of iron-hexacyanoruthenate(II) thin films in aqueous electrolytes: potential analytical and catalytic applications

Materials Science and Engineering B83 (2001) 97 – 105 www.elsevier.com/locate/mseb

Electrochemical behavior of iron-hexacyanoruthenate(II) thin films in aqueous electrolytes: potential analytical and catalytic applications Kasem K. Kasem * Department of Natural, Mathematical and Information Sciences, Indiana Uni6ersity Kokomo, Kokomo, IN 46904 -9003, USA Received 25 May 2000; received in revised form 11 December 2000; accepted 15 December 2000

Abstract Electrochemically prepared thin films of iron-hexacyanoruthenate(II) KFex [Ru{CN}6]y or Ruthenium purple (RP) were used as a surface modifier for the glassy carbon electrode (GCE). The redox behavior of the counter/central ions of RP has been investigated in aqueous electrolytes using CV, CC and electrochemical impedance spectroscopy (EIS). The effect of supporting electrolyte components on the electrochemical behavior was also investigated. In aqueous KClO4, RP showed three redox waves with E of :0.2, 0.9 and 1.2 V versus Ag/AgCl. RP thin films were found to be either ionic size discriminatory or very sensitive to the level of doping with some organic solvents. Acetonitrile as a solvent and/or Li+ have a great distorting effect on these redox waves. The potential analytical application of these films was explored. The EC behavior was compared with that of related HCM (hexacyanometalate) compounds, such as KRux [Ru{CN}6]y, KRux [Fe{CN}6]y and KFex [Fe{CN}6]y. Kinetic studies corresponding to the iron and ruthenium centers in RP have been achieved. Obtained results indicated evidence for catalytic behavior of the RP towards water decomposition. © 2001 Elsevier Science B.V. All rights reserved. Keywords: Electrochemical; Compounds; Decomposition

1. Introduction Insoluble mixed-valence hexacyanometalates (HCM) are important poly-nuclear compounds that attract the attention of many investigators [1 – 7]. The general formula of these cyanometalates is M A [M B(CN)6].xH2O, where M A and M B are transition metal ions with different formal oxidation numbers and are referred to as counter and central ions, respectively. Two major characteristics of this class of inorganic redox films promote its usefulness in applied capacity. Firstly, the oxidation or reduction of these solid films can proceed without dissolution of the solid compound, the film maintains its neutrality by an ion diffusion process. Secondly, formation of a bi- or multi-layered structure is possible using an insertion-substitution mechanism. Furthermore, their well-defined redox transition makes HCM applicable to the fabrication of * Tel.: +1-765-4559245; fax: +1-765-4559566. E-mail address: [email protected] (K.K. Kasem).

sensors [8–10], displays [11], solid state batteries [12,13] and as ion selective electrodes [14]. The electrochemical behavior of the central atom in some hexacyanometalates was investigated using In3 + as a counter cation [15]. Although Prussian blue (PB) has been studied exclusively, studies on Ruthenium purple (RP) as electrochromic films have been carried out [6] and further investigations related to the preparation and characterization of these RP films have been reported [16 –19]. Most of the studies were carried out on the counter ion redox centers. The magnetic and electrochromic properties of these films can be controlled and enhanced by manipulating the structure of the film, the type and degree of doping and the interface composition. Very limited studies have been performed on the central atom redox center or on correlating its behavior to the other redox center. The electrochemical behavior of the iron and ruthenium centers in the KFex [Ru(CN)6]y cluster in aqueous and organic solvents has not yet been fully studied.

0921-5107/01/$ - see front matter © 2001 Elsevier Science B.V. All rights reserved. PII: S 0 9 2 1 - 5 1 0 7 ( 0 1 ) 0 0 5 0 3 - 7

98

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

This paper reports the electrochemical behavior of the multiple redox centers of RP. The counter cation and the central atom redox behavior in different supporting electrolytes were studied. The charge transfer kinetics of iron and Ru centers of this compound were also investigated. Discrimination character of RP towards the ion’s size and charge were investigated for possible analytical application of RP as a selective electrode. Studies were also extended to involve mixed solvent electrolyte. Evidence for electrocatalytic activities of RP towards the oxidation of water is also presented.

2. Experimental

2.1. Reagents Reagent grade K4[Ru(CN)6]2H2O, FeCl3, RuCl3, KCl, KClO4, LiClO4 and analytical grade acetonitrile were obtained from Aldrich Chemical Co. All other reagents were of at least reagent grade and were used without further purification.

2.2. Instrumentation and methods Electrochemical experiments used the electrochemical cell of the conventional three-electrode type, carried out in a small vial container. The reference electrode was Ag/AgCl (saturated KCl) half-cell whose potential was − 45 mV versus SCE, the counter (auxiliary) electrode was a platinum (Pt) wire and the working electrode was glassy-carbon (GC (0.071 cm2 surface area, Bioanalytical Systems)). The working electrode was cleaned by polishing with 1 mm a-alumina paste and rinsed with water and acetone prior to use. BAS 100 B electrochemical analyzer (Bioanalytical Co.) was used to perform the electrochemical studies. EQCN 501 (Elchema, NY) was used to follow up with the mass change during the course of the electrochemical process.

2.3. Impedance measurement Electrochemical impedance spectroscopy (EIS) studies were carried out using a BAS impedance module. Faradic impedance measurements were carried out within a frequency range between 0.1 mHz to 1 kHz. The current response was monitored by a current transducer, whose sensitivity range was automatically adjusted for each frequency range examined. The data were analyzed by a Fourier transform algorithm.

2.4. Chronocoulometry Experiments were performed with 50 m s − 1 pulse width. All electrochemical experiments were carried out

by de-oxygenation using a nitrogen (N2 99.99%) atmosphere. All experiments were performed at room temperature (25.1°C).

2.5. Electrode modification Films of RP were electrodeposited on the GCE surfaces by repetitive potential cycling of a glassy carbon (GC) disk electrode between − 0.20 and + 1.0 V versus Ag/AgCl in freshly prepared aqueous solutions containing equi 0.5 mM of K4[Ru(CN)6]2H2O and FeCl3 and 40 mM KCl (pH 2). Scans were carried out at a sweep rate of 100 mV s − 1. The electrodes were then rinsed thoroughly and transferred to reactant-free electrolyte, where their electrochemical response was examined. The electrode surface coverage (G) of the RP was determined by integrating the area under voltammetric I–E curves.

3. Results and discussion

3.1. Electrodeposition of KFex[Ru(CN)6]y Cyclic voltammograms of a glassy carbon (GC) disk electrode between −0.20 and + 1.0 V versus Ag/AgCl, in aqueous solutions containing equi 0.5 mM of K4[Ru(CN)6].2H2O and FeCl3 and 40 mM KCl (pH 2) is shown in Fig. 1(A). The growth of the redox wave as successive scanning took place is evidence for the build up of the RP film. Fig. 1(B) is a plot of ip versus number of cycles as a function of time. It can be noticed that the cathodic current is always greater than the anodic current. The ratio of ipc/ipa in the last 15 cycles is steady and : 1.3 (Fig. 1B). In the reduction scan, the following EC take place: II + II II FeIII x [Ru (CN)6]y + 4K + 4e=K4Fex [Ru (CN)6]y

Excess Fe(III) that occupies the octahedral interstitial sites will engage in the reduction process, which causes the increase in the reduction current, over the oxidation current leading to the greater ratio of ipc/ipa. Studies indicated that after ten cycles steady DEp of 160 mV is reached; before that time, the formed film permeability for the ions’ movements causes less resistance. The film II permeability decreases as more K4FeII x [Ru (CN)6]y is deposited and the resistance of the film increases causing the IR drop that add to DEp up to 160 mV. II 3.2. Rate of K+Fe II x [Ru (CN)6]y deposition

The slope of the plot of ipc versus number of cycles as a function of time, (Fig. 1C) is 5×10 − 7 mC per cycle. Considering the effective time of deposition per cycle is 2 s, the rate of the film build up is 2.5 mC s − 1. This rate is equivalent to surface coverage of 1.85× 10 − 11 mol

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105 Fig. 1. (A) Repetitive CV at 200 mV s − 1 of GCE in equi 0.5 mM of K4[Ru(CN)6] and FeCl3 in 50 mM KCl/HCl (pH=2). Scan started at − 0.20 V versus Ag/AgCl. (B) ipC/ipA ratio versus number of cycles. (C) ip versus number of cycles.

99

100

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

cm − 2 per cycle, or 2.1× 10 − 8 g per cycle. These results are consistent with the work performed using quartz crystal microbalance measurements. II 3.3. Network structure of KFe III x [Ru (CN)6]y

The calculated unit cell dimensions suggested that the mono-layer thickness is 10.4 A, . This corresponds to a molar volume of 677.16 cm3 mol − 1. Experimentally, the consumption of 1.8 mC cm − 2 generated a film thickness of 300 A, . These experimental quantities correspond to 643.3 cm3 mol − 1. The calculated or experimental molar volume indicates that a mono-layer of : 10.0 A, thick is equivalent to 1.6×10 − 10 mol cm − 2. For multi layers of RP, the network structure, consists II mainly of KFeIII x [Ru (CN)6]y, where the high spin iron center is octahedrally surrounded by six NC ligands and some higher oxidation states ruthenium oxides. Interstitial sites of the octahedrally shaped II KFeIII x [Ru (CN)6]y are occupied by excess Fe(III) ions.

3.4. Redox beha6ior of RP films 3.4.1. Background In hexa cyano metalates (HCM) of general formula Y Y X MX A[M B (CN)6]Z, M B is the central atom, while M A is X Y the fixed counter cation/s. Both M A and M B are gener-

ally transition metals. The reduction of the fixed counter cation of these compounds can take place via electron/cation addition according to the following equation: + Y −1 [ M MX [M Y MX A[M B (CN)6]Z + 1e+M A B (CN)6]Z

(1) While the oxidation of the central atom can take place by electron loss followed by anion addition as follows: − Y Y+1 MX [ MX (CN)6X]Z A[M B (CN)6]Z − 1e+X A[M B

(2) The reduction of the counter cation according to reaction 1 produces the reduced form, while the oxidation of the central atom in the anion complex according to reaction 2 produces the oxidized form. The color of Y the reduced or oxidized forms depends on M X A and M B . In PB, the reduced form is white and is called Prussian white (PW) and the oxidized form is green and is called Prussian green (PG). Thin films of RP at 0.0V versus Ag/AgCl show no absorption peaks within the visible region, which indicates that the color of these films under this applied potential is considered to be white (RW = Ruthenium white). On the other hand, applying more positive potential, :0.50 V versus Ag/AgCl, to the same thin films gives rise to an absorption peak at : 540 nm. The complementary color for this wavelength is violet (RV= Ruthenium violet). The CV for HCM may show redox waves corresponding to the reduction of the counter ion and oxidation of the central atom and, in some cases, the redox waves (if any) of the excess counter ions occupying interstitial sites of the octahedrally shaped metal cyanide. Previous studies [2,3,20] indicated that the redox wave at less positive potential corresponds to the reduction of the counter ion, while that at higher positive potential represents the oxidation of the central atom. Some of these studies [20] indicated that the counter ion affects the formal potential of the central atom.

3.5. Electrochemical beha6ior of RP

Fig. 2. (A) CV at 200 mV s − 1 of GCE modified with RP in 50 mM KClO4. Scan started at − 0.20 V versus Ag/AgCl. (B) CV at 200 mV s − 1 of GCE modified with RP in 50 mM KClO4. Scan started at − 1 V versus Ag/AgCl.

Scanning the potential of GCE modified with RP at 100 mV s − 1 between − 0.20 and 1.30 V in 0.1 M KClO4 produced the CV shown in Fig. 2(A). Three redox surface waves with formal potentials at 0.185, 0.875 and 1.230 V can be identified. The formal potentials of some related HCM, such as KRux [Ru{CN}6]y and KRux [Fe{CN}6]y, are listed in Table 1. It can be noticed that the reduction of Fe(III) as a counter cation according to reaction 1, requires less potential (more positive) than that needed for the reduction of Ru(III) as a counter cation. On the other hand, the oxidation of the central atom depends on the atoms bridging through the N–C group. The data listed in Table 1

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

101

Table 1 Formal potentials (V versus Ag/AgCl) for some immobilized hexacyanometalates in 0.1 M KClO4 at 0.2 V s−1 scan rate. Compound

1st KClO4

lst KCl

2nd KClO4

2nd KCl

3rd KClO4

3rd KCl

KFex [Ru(CN)6]y KRux [Ru(CN)6]y KRux [Fe(CN)6]y KFex [Fe{CN}6]y

0.150 0.050 0.000

0.205 0.000 0.000 0.245

0.850 0.850 0.650

0.850 0.785 0.850

1.220 1.050 1.100

1.235 1.100 1.050 1.045

suggest that the oxidation of the central atom take place at less positive potential (third wave potential) when the counter and central atoms are identical. It is clear that the oxidation potential of Ru in Ru4[Ru(CN)6] is less than that for Ru in KFex [Ru(CN)6]y. Furthermore, the oxidation potential of Fe in KFex [Fe{CN}6]y is less than that for Fe in KRux [Fe(CN)6]y. However, consideration has to be given to the higher oxidation states that only Ru ions can have (IV and VI), which cause the redox wave of formal potential :0.85 V. The fact that Ru in KFex [Ru(CN)6]y showed a redox wave at :0.85 V, indicates that the formation of higher oxidation states of ruthenium (oxoruthenium) does not depend on the position of Ru atom in the network of the polymer. Furthermore, studies showed that the redox waves are independent of each other and are not affected by oxygen reduction (Fig. 2B).

3.6. Effect of supporting electrolyte en6ironments on the redox potential of HCM 3.6.1. A-cation effect Reaction 1 indicates that the reduction process of the counter metal involves gaining electron/s and the addition of a supporting electrolyte cation. The change in the first redox wave shape and potential can be used as an indicator for the cation effect. The results indicated that in 0.1 M HClO4 in the presence of only H3O+ the redox waves steadily decayed. The decay of the redox wave can be explained on the basis of H3O+/K+ aq + exchange. As the H3O+/K+ aq ratio is very large, Kaq + concentrations are not enough to replace H3O during the oxidation process. As more reduction takes place, more H3O+ replaces K+ aq and blocks the re-interring of K+ during the next oxidation process. The fact that in aq acidic KCl the deposited film is stable suggests that reaction 1 is thermodynamically favored and took place with smaller cation size (H3O+ =900 pm, while K+ aq = 300 pm). The partial recovery of the redox wave when the modified electrode is transferred from HClO4 to aqueous KClO4 is an indication of less stability in aqueous acidic electrolytes. This is also evidence for the importance of K+ aq in the stability of the deposited film. Similar behavior was noticed in the cases of aqueous HCl and CF3COOH. The film permeability to different

cation sizes was tested by performing CV for the modified electrode in aqueous 0.1 M LiClO4. The resultant CV is displayed in Fig. 3(A). This figure illustrates that both the cathodic and anodic peak currents decreased and an ill-shaped CV wave is produced. Fig. 3 also shows the effect of the Li+/K+ ratio on both peak current (3B) and potential (3C) of the first redox wave. Fig. 3(B) also indicates that DEp increases as the concentration of Li+ increases.

3.6.2. B-anion effect Reaction 2 involves incorporation of supporting electrolyte anions. Studies show that changing the supporting electrolyte from KCl to KClO4 did not cause a tangible difference in the formal redox potential (Table 1). Such behavior is anticipated since the hydrodynamic radii of Cl− and ClO− ions are 300 and 350 pm, 4 respectively; however, using aqueous CF3COOK (pH 5) no tangible change in the third wave redox potential has been obtained. The greater hydrodynamic radius (600 pm) of CF3COO− did not alter the central atom redox behavior. Such results indicate that the anion size has little or no effect on reaction 2. 3.6.3. C-sol6ent effect The solvent effect has been investigated using the following scheme: first depositing an RP film from 95% acetonitrile/water solution of bath salts, then running a CV for the modified electrode with these thin film in pure and mixed solvent electrolytes. The resultant CVs showed two redox waves at : 0.20 and 1.25 V versus Ag/AgCl. No evidence for the redox wave corresponding to the oxoruthenium has been obtained. Such results indicate the rule of water in the formation of Ru high oxidation state oxides. The CVs of the modified electrode with RP films (prepared in pure aqueous bath) carried out in mixed acetonitrile/water electrolyte are shown in Fig. 4. It can be noticed that with an increase in the acetonitrile/water ratio, both cathodic and anodic peak current decreased and DEp for the first wave (at :0.2 V) increases as well. The results displayed in Fig. 4(B,C) indicates that the modified electrodes with RP thin films have potential analytical applications. Fig. 4(B) shows a linear decrease in both cathodic and anodic peak current. The maximum drop in the peak current took place

102

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

after the initial addition of 10% acetonitrile. The peak current continues to drop with less extent at greater percentages of the acetonitrile. The drop in the peak current is a function of the degree of the organic solvent doping in the RP film. Fig. 4(C) shows that the change in both anodic and cathodic peak potential (DEp). The change is not tangible for small percentage (1 –10%) of acetonitrile, however it increases as the percentage of acetonitrile increases. This greater change in anodic peak potential than that of the cathodic peak potential recommend that the change of the anodic peak potential can be used for potentiometric application. No change is noticed for the wave at : 1.20 V. Comparing Fig. 4 with Fig. 2, one can determine the obvious change for the oxoruthenuim wave. Such differences can be explained as follows. Aqueous solutions of the active components of the RP deposition bath have the following characteristics 1. Aqueous Ru(CN)6− 4 can undergo the following equilibrium:Ru(CN)6− 4 ? Ru2 + +6CN−Ru2 +

+ + H2O ? [Ru(H2O)6]2 + or Ru2(aq) [Ru(H2O)6]2 + ? Oxoruthenium + 3+ 2. FeCl3(aq) ? Fe3(aq) + 3Cl− (aq)Fe(aq) can be written as 3+ [Fe(H2O)6] 3. Ru(CN)6− 4 + H2O ? [Ru(H2O)6]2 + + 6CN − is very slow and thermodynamically is unlikely. Addition of acetonitrile affects the first redox wave (counter ion redox center). This observation concur with the following facts: 1. The Ru(CN)6− 4 is an inert complex with respect to either CH3CN or H2O molecule exchange. 2. The reduction or oxidation of iron ions (as a counter ion redox center) requires exchange of K+. Changing the solvent will affect the solvation of both iron and potassium ions.

3.7. Impedance measurement Typical Nyquest plots for GCE modified with poly KFex [Ru(CN)6]y at 0.155 V versus Ag/AgCl in 0.1 M

Fig. 3. (A) CV at 200 mV s − 1 of GCE modified with RP in 50 mM KClO4. Scan started at − 0.20 V versus Ag/AgCl in different Li+/K+ ratios (top scan is for 0.0 Li+/K+ ratio). (B) Plot of ip versus Li+/K+ ratios (" =ipA, =ipC). (C) Plot of Ep versus Li+/K+ ratios (2= EpA,

= EpC).

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

103

Fig. 4. (A) CV at 200 mV s − 1 of GCE modified with RP in 50 mM KClO4 in the presence of acetonitrile/water mixtures (scan 1 in aqueous electrolyte only). (B) Plot of ip versus acetonitrile/water ratios (= ipA, = ipC). (C) Plot of Ep versus acetonitrile/water ratios ( = EpA,

= EpC).

KClO4 are shown in Fig. 5. This figure exhibited a case of overlapping kinetic/diffusional controlled regions of impedance behavior. At high frequencies, … \ 50, an ill-defined semicircular plot was obtained, while at very low frequencies, 1\… B 0.01, the impedance plot reflected a diffusional control. When measurements were extended to a very small …(0.001 Hz) an indication of charge saturation behavior was observed; however, Fig. 5 can also fit the modified Randles-type equivalent circuit with constant phase element (CPE) which acts as a capacitor. The permeability of the deposited film to aqueous K+, which was confirmed by CV studies, indicates the porous nature of this film. This porosity is a further confirmation for the equivalent circuit type of this electrochemical system. Kinetic data of different electrochemical parameters related to diffusion, charge transfer and double layer capacitance is listed in Table 2. These data were reported for measurements performed under reduction/oxidation of

each of the counter ions (0.128/0.262 V), metal oxide inclusion (0.858/0.923 V) and central atom (1.200 V). An investigation of Table 2 data, we can note the following: 1. The parameters associated with the reduction of the counter cation (Dct, io, ks) are greater than (approximately three times) those related to the oxidation. The obtained CV does not reflect a redox wave of a thin layer (film) as the DEp = 134 mV. This great DEp can not be explained on the basis of IR drop only. It indicates that the redox process is not ideally reversible, but rather sluggish. The effect of the supporting electrolyte cation can be clearly noticed by comparing Dct, io and ks data in the presence of Li+ and K+ ions. Table 2 clearly shows that these parameters in Li+ are much less than in K+. This confirms the effect of cation size in the reduction/oxidation process. The reported Dct, io, ks calculated from either CC or/and IMP measurements

104

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

are smaller than reported [21] for Co(III) hexacyanoferrate (CoHCF). 2. The data listed for the central atom redox center indicates greater io and ks than those for the counter ion. These data reflect the effect of CN− as the major network bridging ion; however, the obtained Dct, io, ks in Cl− are smaller than those in ClO− 4 . These results indicate that it is not only CN− that can affect the magnitude of these parameters but also the supporting electrolyte anion. During the oxidation– reduction process, an exchange of supporting electrolyte anion takes place (cf. reaction 2). The calculated Dct from impedance measurements were much closer to those calculated from CC studies. This can be used as a confirmation for the correctness of the suggested model (modified Randles-type equivalent circuit).

3.8. Catalytic beha6ior of RP film Thin films of RP on the surface of GCE exhibited some electrocatalytic activities towards the oxidation of water. Evidence for such catalytic behavior can be noticed in the CVs obtained in basic solution of CF3COOK (pH 8) or 0.05 M KOH. Fig. 6(A) shows a great increase in the anodic current starting at 0.725 V and an irreversible anodic peak at 0.980 V with a great anodic current (note the scale factor of the dashed curve). The film on the electrode surface underwent cracking and lost \30% of original surface coverage. Such damage was caused by evolution of O2 gas that had been produced through the catalytic oxidation of water. This catalytic behavior cannot be attributed to the inclusion of some oxoruthenium compound within the RP network structure. Modified electrodes with Fe4[Fe(CN)6]3 show similar catalytic behavior (Fig. 6B),

Fig. 5. Nyquest Plot for GCE modified with RP, at 0.155 V versus Ag/AgCl in 0.1 M KClO4. Table 2 Electrochemical impedance measurement data for KFex [Ru(CN)6]y film (G =5.9×10−9 mol cm−2), film thickness 6.72×10−6 m Electrolyte

E,V versus Ag/AgCl

Dct

io, A cm−2

Kct

0.1M KClO4 (aq)

0.128 0.262 0.858 0.923 1.250 0.180 1.280

1.18×10−10 (0.914×10−10)* 2.98×10−11 (2.64×10−11)* 2.1×10−11 1.41×10−11 6.72×10−11 (3.66×10−11)* 7.34×10−12 2.77×10−11

1.49×10−4 4.47×10−5 5.36×10−5 3.20×10−5 3.58×10−4 3.92×10−6 9.33×10−5

2.46×10−5 7.50×10−6 1.13×10−4 6.74×10−5 3.39×10−4 1.65×10−6 1.57×10−5

0.1 M LiClO4 (aq) 0.1 M KCl (aq) * CC data.

K.K. Kasem / Materials Science and Engineering B83 (2001) 97–105

105

of HCM were focused on the counter cation redox potential (at less positive potential). This study gives details for all redox potentials in HCM film containing two transition metals centers. The studied thin films indicated discrimination character towards the cation’s size. The leaner drop in peak current (Fig. 3B) can be invested in using HCM films as sensors for Li+. The sensitivity of the peak current to the level of organic solvent doping inside the film (Fig. 4) can also be used to develop another sensor for organic solvent contaminants. The studies demonstrated the possibility of seeding variety of redox centers in these important inorganic polymers. Manipulating the structures of HCM to perform a desired function is now possible. Potential analytical applications of these thin films as a sensor for Li+ and acetonitrile is in progress. References

Fig. 6. Electrocatalytic activities of RP. (A) CV at 100 mV s − 1 of GCE modified with RP in 50 mM CCl3COOK (pH= 8). (B) CV at 100 mV s − 1 of GCE modified with RP in 50 mM KCl/KOH (pH=8).

which suggests that the catalytic behavior of the HCM film is a characteristic of the film fixed sites not due to the inclusion in the film. Fig. 6 also indicates that the catalytic activity is promoted by the oxidation of the central atom. This can be possible when OH− is attached to the complex followed by the following reaction: 2OH − [ H2O + 12O2 +2e It can be concluded that the catalytic oxidation of water is achieved by the redox center related to the central atom in HCM.

4. Conclusion Most of the published studies of the electrochemistry

.

[1] V.D. Neff, J. Electrochem. Soc. 125 (1978) 886. [2] K. Itaya, H. Akahoshi, S. Toshimo, J. Electrochem. Soc. 129 (1982) 1498 – 1500. [3] L.M. Siperko, T. Kuwana, J. Electrochem. Soc. 130 (1983) 396 – 402. [4] Z. Gao, G. Wang, P. Li, Z. Zhao, Electrochim. Acta 36 (1991) 147 – 152. [5] K.K. Kasem, J. Appl. Electrochem. 29 (1999) 1473. [6] T.R. Catldi, C.E. de Benedetto, C. Campa, J. Electroanal. Chem. 437 (1997) 93 – 98. [7] S. Ohkoski, A. Fujishima, K. Hashimoto, J. Am. Chem. Soc. 120 (1998) 5349 – 5350. [8] A.A. Karyakin, O.V. Gitelmacher, E.E. Karyakina, Anal. Chem. 67 (1995) 2419. [9] A.A. Jaffari, A.P.F. Turner, Biosens. Bioelectron. 12 (1997) 1. [10] H. Li, E. Wang, Microchem. J. 49 (1994) 91. [11] K. Itaya, K. Shibayama, H. Akahoushi, S. Toshima, J. Appl. Phys. 53 (1982) 804. [12] V.D. Neff, J. Electrochem. Soc. 132 (1985) 1382. [13] P.J. Kulesza, J. Electroanal. Chem. 289 (1990) 103. [14] P.J. Kulesza, K. Dobbofer, J. Electroanal. Chem. 274 (1989) 95. [15] M.A. Malik, G. Horanyi, P.J. Kulesza, G. Inzelt, V. Kertesz, R. Schmidt, E. Czirok, J. Chem. 452 (1998) 57 – 62. [16] K.P. Rajan, V.D. Neff, J. Phys. Chem. 86 (1982) 4361. [17] K. Itaya, T. Ataka, S. Toshimo, J. Am. Chem. Soc. 104 (1982) 3751. [18] F.M. Crean, K. Schug, Inorg. Chem. 23 (1984) 853. [19] P.J. Kulesza, J. Electroanal. Chem. 220 (1987) 295. [20] S. Jayarama Reddy, A. Dostal, F. Scholz, J. Electroanal. Chem. 403 (1996) 209 – 212. [21] Z. Gao, J. Bobacka, A. Ivaska, Electrochim. Acta 38 (1993) 379 – 383.