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ferric ions at zeolite-modified electrodes
63
J. Electroanal. Chem, 246 (1988) 63-72 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands
EFFECTS OF SOLUTION pH ON THE VOLTAMMETIUC OF FERROUS/FERRIC IONS AT ZEOLITE-MODIFIED
CHIAICI IWAICURA,
SEIJI MIYAZAIU
and HIROSHI
Department of Applied Chemistry, Faculty of Engineering, Osaka 565 (Japan)
BEHAVIOUB ELECIBODES
YONEYAMA Osaka University,
(Received 4th August 1987; in revised form 2nd December
Yamadaoka 2-1, Sdta,
1987)
ABSTRACT The redox reaction of Fe2+/Fe3+ tons at zeolite-modified electrodes is influenced by the solution pH in such a way that the redox peak currents in the cyclic voltammograms increase almost linearly with the H+ ion concentration in the pH region of ca. 2.0-1.0 but at pHs above ca. 2.0 the redox peaks disappear. The observed pH dependence is discussed in terms of the participation of an equilibrium reaction involving H+ ions in the redox process in the zeolite cavities.
INTRODUCTION
Recently, several groups have paid attention to zeolite-modified electrodes because zeolites offer the prospect of exhibiting size and shape selectivities for molecules in addition to molecular absorbability and ion exchangeability [l-6]. Although transition metal-exchanged zeolites are known to be catalytically active in many chemical reactions, there seem to be no publications on their electrochemical characteristics except for some on silver-exchanged zeolites [5,6]. We have prepared and characterized electrochemically iron-exchanged zeolite-modified electrodes for the purpose of using them eventually as catalytic electrodes for organic electrosynthesis. The present paper describes the effects of solution pH on the voltammetric behaviour of ferrous/ferric ions at zeolite-modified electrodes. EXPERIMENTAL
The zeolites used in this work were Na-X (Zeolum F-9), Na, Ca-A (Zeolum A-5), Na, K-A (Zeolum A-3) and Na mordenite, supplied by Toyo Soda Manufacturing Co., Ltd. Some physicochemical data for these zeolites are given in Table 1. Iron-exchanged zeolites were prepared first by stirring a suspension of 300 mg of the zeolite powder (63-125 pm 0) in 50 ml of 0.25 M FeSO, aqueous solution 0022-0728/88/$03.50
6 1988 Elsevier Sequoia S.A.
(initially adjusted to pH 1.5 with H,SO, solution) for 2 days; they were then centrifuged and washed with doubly distilled water. The centrifuging and washing cycle was repeated five times. The degree of ion exchange was determined by calorimetry using o-phenanthroline. For this purpose the zeolite framework was disintegrated with 0.5 M citric acid in an ultrasonic bath to dissolve iron in the solution [2], except in the case of Na mordenite, where the ion-exchanged Fe’+ was determined by comparing the absorbance of the aqueous solutions at h = 510 nm before and after ion exchange [l]. Zeolite-modified electrodes were prepared in the following way. Iron-exchanged zeolite prepared by the above-described procedure (12.7 mg) was suspended in 0.20 ml of tetrahydrofuran containing 3 mg of dissolved polystyrene as a binder. The suspension was applied onto a platinum substrate (2.54 cm2) and allowed to dry in air. Cyclic voltammetry was carried out at 30 o C using a Pyrex glass H-type cell and a platinum counter-electrode. The potentials are referred to a saturated calomel electrode (SCE). The electrolyte used mainly was a mixed aqueous solution of NaCl and HCl whose pH was adjusted to 1.5 using a 0.1 M solution of each. In this paper, such a solution is denoted as 0.1 M NaCl/O.l M HCl (pH 1.5). The point of zero charge (pzc) was determined by a potentiometric titration method [7,8]; 0.01 M HCl was added dropwise to 50 ml of lo-100 mg zeolite-suspended 0.1 M NaCl solution whose pH was originally adjusted to pH 10. The pH of the solution was recorded and plotted as a function of the volume of titrant. RESULTS AND DISCUSSION
As shown in Table 1, the amount of iron in the zeolite decreased with decreasing effective pore size of the zeolite under the present experimental conditions, i.e. Na-X > Na mordenite > Na, Ca-A > Na, K-A. The crystal ionic diameters and effective diameters of the hydrated ions are known to be about 0.01 and 0.9 nm for H+ with coordination number n = 2,0.23 and 0.35 nm for Na+ with n = 4 or 6, and 0.15 and 0.6 nm for Fe2+ with n = 4 or 6, respectively. The cyclic voltammograms of iron-exchanged zeolite-modified electrodes were measured in 0.1 M NaCl/O.l M HCl @H 1.5). Some data obtained with the iron-exchanged Na-X-modified electrode are shown in Fig. 1. The pair of current
TABLE 1 Some physicochemical data for the zeolites used in this work Zeolite SiO, /AlTO
ratio
EffeCtivepOre size/nm Point of zero charge (pzc) Amount of iron in zeoIite/~mol Degree of ion exchange/%
10 mg-’
Na-X
Na, Ca-A
Na, K-A
Na mordenite
2.5 0.8 6.7 18.5 64
2.0 0.5 6.9 9.05 28
2.0 0.3 7.1 2.20 7.2
19.2 0.7 8.0 16.7 78
65
1.0
d -1.5 -0.2
I 0
I 0.2
I 0.4
I 0.6
I 0.8
I 1.0
E/V vs.(SCE)
Fig. 1. Cyclic voltammogramsof an iron-exchangedNa-X-modified electrodein 0.1 M NaCl/O.l M HCl @H 1.5). Sweeprate: 100 mV s-l. (a) 10th cycle; (b) 50th cycle; (c) 100th cycle; (d) 150th cycle.
peaks at around 0.5 V (vs. SCE) are attributed to the redox reaction of Fe2+/Fe3+ ions incorporated in the zeolites. Only a small fraction (less than 1%) of the iron species in the zeolite is estimated to participate in the redox reaction, as judged from a comparison of the quantity of charge involved in the redox waves at the 150th cycle with the amount of Fe2+ incorporated in the zeolite. When the amount of iron-exchanged zeolite loaded on the platinum substrate was decreased to one tenth, the utilization fraction was found to be increased by about ten times. These findings seem to suggest that the iron species that are close to open windows of zeolites and make contact with the platinum substrate are electroactive. As can be seen from Fig. 1, the redox peak currents increased with increasing number of potential cycles up to about the 150th cycle and then reached steady values. The steady values were kept for several hundreds of cycles although their magnitude depended greatly on the solution pH. Figure 2 shows the cyclic voltammograms of an iron-exchanged Na-X-modified electrode in 0.1 M NaCl/O.l M HCl solutions of different pH. For comparison purposes, the voltammograms of a platinum electrode measured in 0.05 M FeCl,containing 0.1 M NaCl/O.l M HCl solutions of different pH are shown in Fig. 3. In the case of the platinum electrode, the magnitude of the Fe2+/Fe3+ redox peaks was almost independent of the solution pH, as expected. In contrast, the redox peaks at the iron-exchanged Na-X-modified electrode increased as the pH of the solution decreased from 2.0 to 1.0, but in solutions of pH above ca. 2.0 they disappeared. The redox peaks were recovered when the electrode was transferred from solutions of pH greater than 2.0 to solutions of pH 2.0-1.0. In solutions of pH less than 1.0, large redox peaks appeared in the initial stage of the potential sweeps,
66
3
2
-2
-3
-4
I -0.2
I 0
f
I
I
I
I
0.2
0,4
0.6
0.8
1.0
I
E/V vs.kXE)
Fig. 2. Cyclic voltammograms of an iron-exchanged Na-X-modified electrode in 0.1 N NaCl/O.l M HCI of different pH. Sweep rate: 100 mV s-‘. (a) pH 1.0; (b) pH 1.2; (c) pH 1.5; (d) pH 2.0.
6 ,
-6
1
-0.2
I
I
0
0.2
I 0,4
I
I
I
0.6
0.8
1.0
E/V vs.(SCE)
Fig. 3. Cyclic ~01~0~s of a Pt electrode in 0.05 M FeCl+ontaining 0.1 M NaCl/O.l M HCI of different pH. Sweep rate: 100 mV s-l. (a) pH 1.0; (b) pH 1.5; (c) pH 2.0.
-4
-6 -0.2
I
I
I
I
I
I
0
0.2
004
0.6
008
1.0
I
E/V vs.(SCE)
Fig. 4. Cyclic voltammograms of an Na-X-modified electrode in 0.1 M NaCl/O.l M HCl (pH 1.5) (solid line) and 0.1 M NaCl (dashed line), both containing 0.05 M FeCl,. Sweep rate: 100 mV s-l. (a) 0.1 h; (b) 3 h; (c) 18 h.
but when the potential cycling was continued, they soon became small, probably due to gradual dissolution of iron from the zeolite cavities. The characteristic features of the iron-exchanged Na-X-modified electrodes described here were observed at all the iron-exchanged zeolite-modified electrodes tested. In order to confirm that the above-mentioned pH dependence of the voltammettic peak currents observed at the iron-exchanged zeolite-modified electrodes is characteristic of zeolite-modified electrodes, the voltammograms of Na-X-modified electrodes without iron exchange were measured in 0.1 M NaCl/O.l M HCl (pH 1.5) and 0.1 M NaCl, both containing 0.05 it4 FeCl,. With continuation of the potential sweep cycles, the redox waves gradually developed, as Fig. 4 shows, indicating that the incorporation of Fe’+ ions into the zeolite cavities progressed slowly in the acidic solution. However, in the neutral solution, the redox waves disappeared. It is then concluded that the voltammetric behaviour of Fe’+ shown in Figs. 1 and 2 is due to the Fe2+ held in the zeolite. On the basis of Fig. 2 and related data, the Fe2+/Fe3+ redox peak currents at the iron-exchanged zeolite-modified electrodes are plotted in Fig. 5 as a function of the H+ ion concentration of the electrolyte solution. A nearly first-order dependence on the H+ ion concentration can be seen at the Na-X- and Na, Ca-A-modified electrodes, but for the Na mordenite-modified electrodes the dependence seems rather complex. In the case of the Na, K-A-modified electrode, the peak currents were very small because of the small amount of incorporated iron species (see Table l), but the dependence is judged to be virtually the same as those of the Na-X- and Na, Ca-A-modified electrodes.
68
[a+1/niol
m-3
Fig. 5. Plots of the anodic (open symbols, i,,) and cathodic (closed symbols, ipc) peak currents against the H+ ion concentration. Electrolyte and sweep rate as in Fig. 2. (a) Iron-exchanged Na-X-modified electrode; (b) iron-exchanged Na, Ca-A-modified electrode; (c) iron-exchanged Na, K-A-modified electrode; (d) iron-exchanged Na mordenite-modified electrode.
By analogy with the electrochemical reduction of a silver-exchanged the Fe2+/Fe3+ redox reaction may be written as (Z0)3Fe3+
+ M+ + e- = (ZO)3Fe2+,
M+
zeolite [5,6],
(I)
where (ZO) represents the zeolite lattice oxygen ion and M+ represents the electrolyte cation (Na+ or H+ in this case) which enters the zeolite cavities for charge compensation. Reaction (1) must take place during the redox reaction, but it cannot explain the observed pH dependence of the peak currents, because in this case both H+ and Na+ ions can participate equally in the redox reaction of Fe2+/Fe3+ so that voltammetric current peaks should appear in neutral NaCl solutions as well as in acidic NaCl/HCl solutions, but this is contrary to the observed results. It is believed that the pH dependence of the peak current is not brought about by destruction of the zeolite framework to dissolve Fe’+ in solution, because the pH dependence is
69
also observed for electrodes prepared from Na mordenite, which has a high acidic stability. One plausible explanation is to assume in the zeolite cavities the participation of equilibrium reactions involving H+ ions [9-111 such as (ZO),Fe*+(H,O).
+ (ZO) = (ZO),Fe*+(H,O)“_,(OH-)
+ (ZO)H+
(2)
where (H,O), represents hydrated water molecules and n = 3 holds in the case of totally hydrated zeolite [lo]. A similar equation can be written for Fe3+ ions, which are known to form more stable hydroxides. The dissociation of the hydrated water in the zeolite cavities is caused by proton uptake by the lattice oxygen of the zeolite framework (ZO), and its degree must depend on the pH of the electrolyte solutions. The formation of (ZO),Fe2’(H20),_,(OH-) by reaction (2) must not exclude the equivalent reaction with the solution species, i.e. (ZO),Fe’+(H,O).
+ Na:
= (ZO)sFe*+(H,O),,_,(OH-)Na++
(s = solution)
Hz (3)
the pH of which is dependent on the acidity of zeolite. Hereafter, (ZO),Fe2’(H20), is abbreviated to Fe*+(H,O) and (ZO),Fe2’(H20),_,(OH-) to Fe”OH-. If it is assumed that Fe’+(H,O) is electroactive while Fe*+OHis electroinactive, the amount of electroactive species Fe*+(H,O) increases and thereby the redox current i increases linearly with an increase in the H+ ion concentration as expressed by i = I;ilk[Fe2+(H20)]
= I;AkK;‘[Fe*+OH-][H+]
= FAkK;‘[Fe*+OH-(Na+)]
[H:]/[Naz]
where K, and K, are the equilibrium constants of reactions (2) and (3), respectively. Equation (4) can explain fairly well the experimental data, except for those obtained at iron-exchanged Na mordenite-modified electrodes. In the case of Na mordenite, the ratio of SiO,/Al,O, as its constituents is quite different from those of the other three kinds of zeolites Na-X, Na, Ca-A and Na, K-A as shown in Table 1, and the proton uptake behaviour of Na mordenite would then be different from that of the other zeolites, requiring some modification of eqn. (2). In order to clarify this point, the following pzc data were obtained. The values of the pzc obtained from titration curves for zeolites without iron exchange are given in Table 1. The adsorption density of potential-determining ions, (I’,+ - IoH-), was determined from the titration curves for the zeolites with and without iron exchange and is shown in Fig. 6 as a function of pH. Unfortunately, the adsorption density at pHs lower than those given in this figure was uncertain because of large errors in the readings of pH on the titration curves and is not shown in the figure. It is seen from the dependence of adsorption density on the pH that H+ uptake or OH- release progresses with decreasing pH for both zeolites with and without the incorporated Fe 2+. However, the rate of H+ uptake or that of OHrelease with decreasing pH is quite low at Na mordenite, compared with those at the other three kinds of zeolites. The presence of incorporated Fe’+ does not alter basically the H+ uptake behaviour with the pH, but it has a tendency to make the
70 -1.5
-1.0
7OI-0.5 g '3: Lo I :=
0
0.5
1.0 3
I 4
b
,d
,
I
I
I
I
5
6
7
8
9
10
PH
Fig. 6. Adsorption density of potential-determining ions, (r,+ - ro,- ), at zeolites (open symbols) and iron-exchanged zeolites (closed symbols) as a function of pH. (a) Na-X; (b) Na, Ca-A; (c) Na, K-A; (d) Na mordenite.
adsorption density negative. This is understandable when it is considered that the potential-determining ions on the oxide surface are H+ and OH- and that the positive charges of the incorporated iron ions bind the negative charge (OH-) in excess of those of zeolites without the incorporated iron ions. Anyway, the difference in the proton uptake behaviour between Na mordenite samples and the other samples could give a qualitative basis for the appearance of the different dependences of these voltammetric peak currents on the pH as given in Fig. 5. The reaction given by eqn. (1) does not control the redox reaction in 0.1 M NaCl/O.l M HCl solutions as discussed above, but electrolyte cations must be involved in the redox process in such a manner as that given by eqn. (1) to maintain charge neutrality in the zeolite cavities. If the electrolyte cations were too large to
71
a
1,O
-0.5
-1,o
-1.5 1 ' 0
I
I
I
I
0,2
0.4
0.6
0.8
E/V vs.(SCE)
Fig. 7. Cyclic voltammograms of an iron-exchanged Na-X-modified electrode in different electrolyte solutions. Sweep rate: 100 mV s-l. (a) 0.1 M NaCl/O.l M HCl (pH 1.5); (b) 0.1 M (C,H,),NCl/O.l M HCl (pH 1.5); (c) 0.1 M (C,H,),NCl/O.l M HCl @H 1.5).
enter the zeolite cavities, the redox reaction would be suppressed. In order to confirm this hypothesis, the effects of electrolyte cations on the voltammetric behaviour of the iron-exchanged zeolite-modified electrodes were examined for Na-X whose pore size is 0.8 nm. The results are shown in Fig. 7. The diameters of the N(C,H,): and N(C.,H,): ions used as the electrolyte cations are known to be about 0.7 and 1 nm, respectively [12]. It can be seen that the redox peak currents were markedly suppressed when N(C,H,): was used, as expected. Since the solution pH was 1.5, a contribution of H+ to the charge compensation in the redox reaction may be expected for all the electrolyte solutions used. Nevertheless, the effect of cation size appeared clearly.
ACKNOWLEDGEMENT
The present work was partially supported by a Grant-in-Aid for Scientific Research (No. 61470080) from the Ministry of Education, Science and Culture.
72
REFERENCES 1 2 3 4 S 6 7 8 9 10 11 12
Z. Li and T.E. Mallouk, J. Phys. Chem., 91 (1987) 643. H.A. Gemborys and B.R. Shaw, J. Electroanal. C&em., 208 (1986) 95. B. De Viimes, F. Bedioui, J. Devynck and C. Bied-Charreton, J. Electroanal. Chem., 187 (1985) 197. C.G. Murray, R.J. Nowak and D.R. Rolison, J. ElectroanaJ. Chem., 164 (1984) 205. J. Sarradin, J.-M. Louvet, R Messina and J. Perichon, Zeolites, 4 (1984) 157. J.-P. PereiraRamos, R Messina and J. Perichon, J. Electroanal. Chem., 146 (1983) 157. G.A. Parks and P.L. de Bruyn, J. Phys. Chem., 66 (1962) 967. S.M. Ahmed, J. Phys. Chem., 73 (1969) 3546. B.L. Dickson and L.V.C. Rees, J. Chem. Sot., Faraday Trans. 1, 70 (1974) 2038, 2060. Z. Gao and L.V.C. Rees, Zeolites, 2 (1982) 221. W.N. Delgass, R.L. Garten and M. Boudart, J. Phys. Chem., 73 (1969) 2970. B.M. Lok, T.R. Cannan and C.A. Messina, Zeolites, 3 (1983) 282.
J. Electroanal. Chem, 246 (1988) 63-72 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands
EFFECTS OF SOLUTION pH ON THE VOLTAMMETIUC OF FERROUS/FERRIC IONS AT ZEOLITE-MODIFIED
CHIAICI IWAICURA,
SEIJI MIYAZAIU
and HIROSHI
Department of Applied Chemistry, Faculty of Engineering, Osaka 565 (Japan)
BEHAVIOUB ELECIBODES
YONEYAMA Osaka University,
(Received 4th August 1987; in revised form 2nd December
Yamadaoka 2-1, Sdta,
1987)
ABSTRACT The redox reaction of Fe2+/Fe3+ tons at zeolite-modified electrodes is influenced by the solution pH in such a way that the redox peak currents in the cyclic voltammograms increase almost linearly with the H+ ion concentration in the pH region of ca. 2.0-1.0 but at pHs above ca. 2.0 the redox peaks disappear. The observed pH dependence is discussed in terms of the participation of an equilibrium reaction involving H+ ions in the redox process in the zeolite cavities.
INTRODUCTION
Recently, several groups have paid attention to zeolite-modified electrodes because zeolites offer the prospect of exhibiting size and shape selectivities for molecules in addition to molecular absorbability and ion exchangeability [l-6]. Although transition metal-exchanged zeolites are known to be catalytically active in many chemical reactions, there seem to be no publications on their electrochemical characteristics except for some on silver-exchanged zeolites [5,6]. We have prepared and characterized electrochemically iron-exchanged zeolite-modified electrodes for the purpose of using them eventually as catalytic electrodes for organic electrosynthesis. The present paper describes the effects of solution pH on the voltammetric behaviour of ferrous/ferric ions at zeolite-modified electrodes. EXPERIMENTAL
The zeolites used in this work were Na-X (Zeolum F-9), Na, Ca-A (Zeolum A-5), Na, K-A (Zeolum A-3) and Na mordenite, supplied by Toyo Soda Manufacturing Co., Ltd. Some physicochemical data for these zeolites are given in Table 1. Iron-exchanged zeolites were prepared first by stirring a suspension of 300 mg of the zeolite powder (63-125 pm 0) in 50 ml of 0.25 M FeSO, aqueous solution 0022-0728/88/$03.50
6 1988 Elsevier Sequoia S.A.
(initially adjusted to pH 1.5 with H,SO, solution) for 2 days; they were then centrifuged and washed with doubly distilled water. The centrifuging and washing cycle was repeated five times. The degree of ion exchange was determined by calorimetry using o-phenanthroline. For this purpose the zeolite framework was disintegrated with 0.5 M citric acid in an ultrasonic bath to dissolve iron in the solution [2], except in the case of Na mordenite, where the ion-exchanged Fe’+ was determined by comparing the absorbance of the aqueous solutions at h = 510 nm before and after ion exchange [l]. Zeolite-modified electrodes were prepared in the following way. Iron-exchanged zeolite prepared by the above-described procedure (12.7 mg) was suspended in 0.20 ml of tetrahydrofuran containing 3 mg of dissolved polystyrene as a binder. The suspension was applied onto a platinum substrate (2.54 cm2) and allowed to dry in air. Cyclic voltammetry was carried out at 30 o C using a Pyrex glass H-type cell and a platinum counter-electrode. The potentials are referred to a saturated calomel electrode (SCE). The electrolyte used mainly was a mixed aqueous solution of NaCl and HCl whose pH was adjusted to 1.5 using a 0.1 M solution of each. In this paper, such a solution is denoted as 0.1 M NaCl/O.l M HCl (pH 1.5). The point of zero charge (pzc) was determined by a potentiometric titration method [7,8]; 0.01 M HCl was added dropwise to 50 ml of lo-100 mg zeolite-suspended 0.1 M NaCl solution whose pH was originally adjusted to pH 10. The pH of the solution was recorded and plotted as a function of the volume of titrant. RESULTS AND DISCUSSION
As shown in Table 1, the amount of iron in the zeolite decreased with decreasing effective pore size of the zeolite under the present experimental conditions, i.e. Na-X > Na mordenite > Na, Ca-A > Na, K-A. The crystal ionic diameters and effective diameters of the hydrated ions are known to be about 0.01 and 0.9 nm for H+ with coordination number n = 2,0.23 and 0.35 nm for Na+ with n = 4 or 6, and 0.15 and 0.6 nm for Fe2+ with n = 4 or 6, respectively. The cyclic voltammograms of iron-exchanged zeolite-modified electrodes were measured in 0.1 M NaCl/O.l M HCl @H 1.5). Some data obtained with the iron-exchanged Na-X-modified electrode are shown in Fig. 1. The pair of current
TABLE 1 Some physicochemical data for the zeolites used in this work Zeolite SiO, /AlTO
ratio
EffeCtivepOre size/nm Point of zero charge (pzc) Amount of iron in zeoIite/~mol Degree of ion exchange/%
10 mg-’
Na-X
Na, Ca-A
Na, K-A
Na mordenite
2.5 0.8 6.7 18.5 64
2.0 0.5 6.9 9.05 28
2.0 0.3 7.1 2.20 7.2
19.2 0.7 8.0 16.7 78
65
1.0
d -1.5 -0.2
I 0
I 0.2
I 0.4
I 0.6
I 0.8
I 1.0
E/V vs.(SCE)
Fig. 1. Cyclic voltammogramsof an iron-exchangedNa-X-modified electrodein 0.1 M NaCl/O.l M HCl @H 1.5). Sweeprate: 100 mV s-l. (a) 10th cycle; (b) 50th cycle; (c) 100th cycle; (d) 150th cycle.
peaks at around 0.5 V (vs. SCE) are attributed to the redox reaction of Fe2+/Fe3+ ions incorporated in the zeolites. Only a small fraction (less than 1%) of the iron species in the zeolite is estimated to participate in the redox reaction, as judged from a comparison of the quantity of charge involved in the redox waves at the 150th cycle with the amount of Fe2+ incorporated in the zeolite. When the amount of iron-exchanged zeolite loaded on the platinum substrate was decreased to one tenth, the utilization fraction was found to be increased by about ten times. These findings seem to suggest that the iron species that are close to open windows of zeolites and make contact with the platinum substrate are electroactive. As can be seen from Fig. 1, the redox peak currents increased with increasing number of potential cycles up to about the 150th cycle and then reached steady values. The steady values were kept for several hundreds of cycles although their magnitude depended greatly on the solution pH. Figure 2 shows the cyclic voltammograms of an iron-exchanged Na-X-modified electrode in 0.1 M NaCl/O.l M HCl solutions of different pH. For comparison purposes, the voltammograms of a platinum electrode measured in 0.05 M FeCl,containing 0.1 M NaCl/O.l M HCl solutions of different pH are shown in Fig. 3. In the case of the platinum electrode, the magnitude of the Fe2+/Fe3+ redox peaks was almost independent of the solution pH, as expected. In contrast, the redox peaks at the iron-exchanged Na-X-modified electrode increased as the pH of the solution decreased from 2.0 to 1.0, but in solutions of pH above ca. 2.0 they disappeared. The redox peaks were recovered when the electrode was transferred from solutions of pH greater than 2.0 to solutions of pH 2.0-1.0. In solutions of pH less than 1.0, large redox peaks appeared in the initial stage of the potential sweeps,
66
3
2
-2
-3
-4
I -0.2
I 0
f
I
I
I
I
0.2
0,4
0.6
0.8
1.0
I
E/V vs.kXE)
Fig. 2. Cyclic voltammograms of an iron-exchanged Na-X-modified electrode in 0.1 N NaCl/O.l M HCI of different pH. Sweep rate: 100 mV s-‘. (a) pH 1.0; (b) pH 1.2; (c) pH 1.5; (d) pH 2.0.
6 ,
-6
1
-0.2
I
I
0
0.2
I 0,4
I
I
I
0.6
0.8
1.0
E/V vs.(SCE)
Fig. 3. Cyclic ~01~0~s of a Pt electrode in 0.05 M FeCl+ontaining 0.1 M NaCl/O.l M HCI of different pH. Sweep rate: 100 mV s-l. (a) pH 1.0; (b) pH 1.5; (c) pH 2.0.
-4
-6 -0.2
I
I
I
I
I
I
0
0.2
004
0.6
008
1.0
I
E/V vs.(SCE)
Fig. 4. Cyclic voltammograms of an Na-X-modified electrode in 0.1 M NaCl/O.l M HCl (pH 1.5) (solid line) and 0.1 M NaCl (dashed line), both containing 0.05 M FeCl,. Sweep rate: 100 mV s-l. (a) 0.1 h; (b) 3 h; (c) 18 h.
but when the potential cycling was continued, they soon became small, probably due to gradual dissolution of iron from the zeolite cavities. The characteristic features of the iron-exchanged Na-X-modified electrodes described here were observed at all the iron-exchanged zeolite-modified electrodes tested. In order to confirm that the above-mentioned pH dependence of the voltammettic peak currents observed at the iron-exchanged zeolite-modified electrodes is characteristic of zeolite-modified electrodes, the voltammograms of Na-X-modified electrodes without iron exchange were measured in 0.1 M NaCl/O.l M HCl (pH 1.5) and 0.1 M NaCl, both containing 0.05 it4 FeCl,. With continuation of the potential sweep cycles, the redox waves gradually developed, as Fig. 4 shows, indicating that the incorporation of Fe’+ ions into the zeolite cavities progressed slowly in the acidic solution. However, in the neutral solution, the redox waves disappeared. It is then concluded that the voltammetric behaviour of Fe’+ shown in Figs. 1 and 2 is due to the Fe2+ held in the zeolite. On the basis of Fig. 2 and related data, the Fe2+/Fe3+ redox peak currents at the iron-exchanged zeolite-modified electrodes are plotted in Fig. 5 as a function of the H+ ion concentration of the electrolyte solution. A nearly first-order dependence on the H+ ion concentration can be seen at the Na-X- and Na, Ca-A-modified electrodes, but for the Na mordenite-modified electrodes the dependence seems rather complex. In the case of the Na, K-A-modified electrode, the peak currents were very small because of the small amount of incorporated iron species (see Table l), but the dependence is judged to be virtually the same as those of the Na-X- and Na, Ca-A-modified electrodes.
68
[a+1/niol
m-3
Fig. 5. Plots of the anodic (open symbols, i,,) and cathodic (closed symbols, ipc) peak currents against the H+ ion concentration. Electrolyte and sweep rate as in Fig. 2. (a) Iron-exchanged Na-X-modified electrode; (b) iron-exchanged Na, Ca-A-modified electrode; (c) iron-exchanged Na, K-A-modified electrode; (d) iron-exchanged Na mordenite-modified electrode.
By analogy with the electrochemical reduction of a silver-exchanged the Fe2+/Fe3+ redox reaction may be written as (Z0)3Fe3+
+ M+ + e- = (ZO)3Fe2+,
M+
zeolite [5,6],
(I)
where (ZO) represents the zeolite lattice oxygen ion and M+ represents the electrolyte cation (Na+ or H+ in this case) which enters the zeolite cavities for charge compensation. Reaction (1) must take place during the redox reaction, but it cannot explain the observed pH dependence of the peak currents, because in this case both H+ and Na+ ions can participate equally in the redox reaction of Fe2+/Fe3+ so that voltammetric current peaks should appear in neutral NaCl solutions as well as in acidic NaCl/HCl solutions, but this is contrary to the observed results. It is believed that the pH dependence of the peak current is not brought about by destruction of the zeolite framework to dissolve Fe’+ in solution, because the pH dependence is
69
also observed for electrodes prepared from Na mordenite, which has a high acidic stability. One plausible explanation is to assume in the zeolite cavities the participation of equilibrium reactions involving H+ ions [9-111 such as (ZO),Fe*+(H,O).
+ (ZO) = (ZO),Fe*+(H,O)“_,(OH-)
+ (ZO)H+
(2)
where (H,O), represents hydrated water molecules and n = 3 holds in the case of totally hydrated zeolite [lo]. A similar equation can be written for Fe3+ ions, which are known to form more stable hydroxides. The dissociation of the hydrated water in the zeolite cavities is caused by proton uptake by the lattice oxygen of the zeolite framework (ZO), and its degree must depend on the pH of the electrolyte solutions. The formation of (ZO),Fe2’(H20),_,(OH-) by reaction (2) must not exclude the equivalent reaction with the solution species, i.e. (ZO),Fe’+(H,O).
+ Na:
= (ZO)sFe*+(H,O),,_,(OH-)Na++
(s = solution)
Hz (3)
the pH of which is dependent on the acidity of zeolite. Hereafter, (ZO),Fe2’(H20), is abbreviated to Fe*+(H,O) and (ZO),Fe2’(H20),_,(OH-) to Fe”OH-. If it is assumed that Fe’+(H,O) is electroactive while Fe*+OHis electroinactive, the amount of electroactive species Fe*+(H,O) increases and thereby the redox current i increases linearly with an increase in the H+ ion concentration as expressed by i = I;ilk[Fe2+(H20)]
= I;AkK;‘[Fe*+OH-][H+]
= FAkK;‘[Fe*+OH-(Na+)]
[H:]/[Naz]
where K, and K, are the equilibrium constants of reactions (2) and (3), respectively. Equation (4) can explain fairly well the experimental data, except for those obtained at iron-exchanged Na mordenite-modified electrodes. In the case of Na mordenite, the ratio of SiO,/Al,O, as its constituents is quite different from those of the other three kinds of zeolites Na-X, Na, Ca-A and Na, K-A as shown in Table 1, and the proton uptake behaviour of Na mordenite would then be different from that of the other zeolites, requiring some modification of eqn. (2). In order to clarify this point, the following pzc data were obtained. The values of the pzc obtained from titration curves for zeolites without iron exchange are given in Table 1. The adsorption density of potential-determining ions, (I’,+ - IoH-), was determined from the titration curves for the zeolites with and without iron exchange and is shown in Fig. 6 as a function of pH. Unfortunately, the adsorption density at pHs lower than those given in this figure was uncertain because of large errors in the readings of pH on the titration curves and is not shown in the figure. It is seen from the dependence of adsorption density on the pH that H+ uptake or OH- release progresses with decreasing pH for both zeolites with and without the incorporated Fe 2+. However, the rate of H+ uptake or that of OHrelease with decreasing pH is quite low at Na mordenite, compared with those at the other three kinds of zeolites. The presence of incorporated Fe’+ does not alter basically the H+ uptake behaviour with the pH, but it has a tendency to make the
70 -1.5
-1.0
7OI-0.5 g '3: Lo I :=
0
0.5
1.0 3
I 4
b
,d
,
I
I
I
I
5
6
7
8
9
10
PH
Fig. 6. Adsorption density of potential-determining ions, (r,+ - ro,- ), at zeolites (open symbols) and iron-exchanged zeolites (closed symbols) as a function of pH. (a) Na-X; (b) Na, Ca-A; (c) Na, K-A; (d) Na mordenite.
adsorption density negative. This is understandable when it is considered that the potential-determining ions on the oxide surface are H+ and OH- and that the positive charges of the incorporated iron ions bind the negative charge (OH-) in excess of those of zeolites without the incorporated iron ions. Anyway, the difference in the proton uptake behaviour between Na mordenite samples and the other samples could give a qualitative basis for the appearance of the different dependences of these voltammetric peak currents on the pH as given in Fig. 5. The reaction given by eqn. (1) does not control the redox reaction in 0.1 M NaCl/O.l M HCl solutions as discussed above, but electrolyte cations must be involved in the redox process in such a manner as that given by eqn. (1) to maintain charge neutrality in the zeolite cavities. If the electrolyte cations were too large to
71
a
1,O
-0.5
-1,o
-1.5 1 ' 0
I
I
I
I
0,2
0.4
0.6
0.8
E/V vs.(SCE)
Fig. 7. Cyclic voltammograms of an iron-exchanged Na-X-modified electrode in different electrolyte solutions. Sweep rate: 100 mV s-l. (a) 0.1 M NaCl/O.l M HCl (pH 1.5); (b) 0.1 M (C,H,),NCl/O.l M HCl (pH 1.5); (c) 0.1 M (C,H,),NCl/O.l M HCl @H 1.5).
enter the zeolite cavities, the redox reaction would be suppressed. In order to confirm this hypothesis, the effects of electrolyte cations on the voltammetric behaviour of the iron-exchanged zeolite-modified electrodes were examined for Na-X whose pore size is 0.8 nm. The results are shown in Fig. 7. The diameters of the N(C,H,): and N(C.,H,): ions used as the electrolyte cations are known to be about 0.7 and 1 nm, respectively [12]. It can be seen that the redox peak currents were markedly suppressed when N(C,H,): was used, as expected. Since the solution pH was 1.5, a contribution of H+ to the charge compensation in the redox reaction may be expected for all the electrolyte solutions used. Nevertheless, the effect of cation size appeared clearly.
ACKNOWLEDGEMENT
The present work was partially supported by a Grant-in-Aid for Scientific Research (No. 61470080) from the Ministry of Education, Science and Culture.
72
REFERENCES 1 2 3 4 S 6 7 8 9 10 11 12
Z. Li and T.E. Mallouk, J. Phys. Chem., 91 (1987) 643. H.A. Gemborys and B.R. Shaw, J. Electroanal. C&em., 208 (1986) 95. B. De Viimes, F. Bedioui, J. Devynck and C. Bied-Charreton, J. Electroanal. Chem., 187 (1985) 197. C.G. Murray, R.J. Nowak and D.R. Rolison, J. ElectroanaJ. Chem., 164 (1984) 205. J. Sarradin, J.-M. Louvet, R Messina and J. Perichon, Zeolites, 4 (1984) 157. J.-P. PereiraRamos, R Messina and J. Perichon, J. Electroanal. Chem., 146 (1983) 157. G.A. Parks and P.L. de Bruyn, J. Phys. Chem., 66 (1962) 967. S.M. Ahmed, J. Phys. Chem., 73 (1969) 3546. B.L. Dickson and L.V.C. Rees, J. Chem. Sot., Faraday Trans. 1, 70 (1974) 2038, 2060. Z. Gao and L.V.C. Rees, Zeolites, 2 (1982) 221. W.N. Delgass, R.L. Garten and M. Boudart, J. Phys. Chem., 73 (1969) 2970. B.M. Lok, T.R. Cannan and C.A. Messina, Zeolites, 3 (1983) 282.