J. Phys. Chem. 1995, 99, 6635-6639
6635
Silver Zeolite-Modified Electrodes: Electron-Transport Mechanism and Ion-Exchange Kinetics for Partially Silver Exchanged Zeolite Y Mark D. Baker,* Jiwu Zhang, and Michael McBrien Department of Chemistry and Biochemistry, Guelph-Waterloo Centre for Graduate Work in Chemistry, University of Guelph, Guelph, Ontario, Canada N1 G 2WI Received: October 4, 1994; In Final Form: February 7, 1995@
The pseudo-fust-order rate constants for ion exchange of partially silver ion exchanged Y zeolites (AgnNa56-,Y) with sodium ions are determined from chronoamperometric studies of zeolite-modified electrodes in aqueous solution. In the case of Ag,Na56-,Y when n I5.9 the electrochemical reduction of intrazeolite Ag+ cations is controlled by an intracrystalline ion-exchange process rather than by counterdiffusion of exchanging cations through the interconnected supercage system of the zeolite. Herein the methodology for analysis of this situation is developed where the rate of the electrochemical reaction is equated with the velocity of the ratelimiting ion-exchange reaction. Good agreement between theory and experiment was observed. The pseudofirst-order rate constant for Ag+ exchange with Na+ is relatively insensitive to the concentration of sodium in the electrolyte solution, but decreases as the silver loading in the zeolite is increased up to 5.9 silver ions per unit cell. The data are consistent with electron transfer to silver ions occurring subsequent to the ionexchange step, confirming once again that electron transfer occurs outside of the zeolite framework.
Introduction The ion-exchange properties of zeolites are important in many of the application areas associated with these materials. These include, for example, uses in catalysis, wastewater cleanup, ion chromatography, synthesis of quantum dots and wires, sequestration of radioactive cations and detergent technology. In the past few decades, a considerable research effort has been devoted to the understanding of zeolite ion exchange. The majority of these studies have focused on the equilibrium properties of binary and temary exchange systems, and this area has been extensively The link between the zeolite ion-exchange process and the electrochemical reduction of intrazeolite cations is a very strong since it is often the case that the electroactive intrazeolite cation is completely or partially charge balancing. Consequently, the electrochemical reduction of zeolite extraframework cations demands the ingress of an electrolyte co-ion into the zeolite pore system. In the case of AgY zeolite-modified electrodes, the ion-exchange step clearly occurs before the electron transfer p r o ~ e s s .When ~ the concentration of co-ion in solution is small, the very early stages of the ion-exchange reaction are probed at the zeolite-modified electrode,’ since the quantity of Ag+ reduced per electrochemical scan is small compared to the total Ag+ present in the zeolite. Since the zeolite-modified electrode is clearly far from equilibrium, in an ion-exchange sense, this necessitates a careful consideration of the kinetic parameters that control such a situation. In contrast to equilibrium studies, there is a paucity of data available pertaining to cation-exchange kinetics, although there has been some progress in this area. Rees and his co-workers have successfully followed divalent cation self-exchange for several zeolites using radiotracer showing that in these cases the kinetics were diffusion controlled. Kinetic studies of ion exchange in zeolites A12 and ZSM-513,’4using pressure-jump techniques were also reported by Ikeda et al. in @Abstractpublished in Advance ACSAbstracts, April 1, 1995.
the early 1980’s. The assembly of kinetic data conceming monovalent ion exchange in large-pore zeolites, however, presents a considerable challenge. In the case of zeolites X and Y, for example, the oxygen 12-ring apertures that connect the supercages are large enough for small cations to move through rapidly. In this case temperature-jump methods15 can be employed. However, ion exchange involving monovalent cations in zeolites X and Y is in general very difficult to follow.16 In this paper we present an electrochemical method for following fast ion exchange between highly mobile monovalent cations for the special case in which the intrazeolite cation is electroactive. When the intrazeolite counterdiffusion of exchanging cations is rate limiting, the rate of ion exchange is linear with t-l’* and the counterdiffusion coefficient can be readily determined, as shown, for example, for aqueous Na+/Ca2+ exchange in zeolite A,17 and for nonaqueous divalent cation exchange in zeolites X and Y.18,19 Chronoamperometric data for Ags6Y display Cottrellian behavior, indicating a diffusion-controlled electrochemical r e a ~ t i o n .In ~ this case the rate-limiting step involves cation counterdiffusion through interconnecting supercages and is not described by diffusion-controlled electron hopping.20s21 Since the reduction of silver cations is limited by the cation dynamics in the zeolite, the electrochemical response becomes a useful probe of the ion-exchange kinetics. In the remainder of this paper, we will describe the modeling of the electrochemical response of AgnNas6-,Y-modified electrodes ( n I5.9) in terms of a rate-limiting pseudo-first-order intracrystalline ion-exchange reaction. In so doing we will choose, for the sake of clarity, to describe zeolite Y as being composed of two interconnected channel systems. The larger channel is composed of interconnected supercages joined via oxygen 12-rings, and the smaller channel system is composed of interconnected sodalite cages and hexagonal prisms. Note that diffusion of extraframework cations is much faster in the large channel due to steric constraints of the small channel. In low-loading AgnNa56-,,Y, silver cations preferentially occupy
0022-365419512099-6635$09.0010 0 1995 American Chemical Society
Baker et al.
6636 J. Phys. Chem., Vol. 99, No. 17, 1995 the small-channel system. When the zeolite cocation is sodium, and when there are fewer than 4 silver cations per unit cell, silver ions have been reported to exclusively occupy the hexagonal prisms .
paper, namely, reduction of Ag+ originally situated in the smallchannel system of zeolite Y. extrazeolite electron transfer Ag+(z)
Experimental Section Zeolite Y (LZY 52) was donated by UOP Corporation (Whistler, AL). Ion exchange was effected in 0.01 M or more dilute solutions of aqueous silver nitrate to produce the desired ion exchange level. Typically 1 g of zeolite was exchanged in 250 mL of solution. The ion-exchange reaction was performed in the dark, and following an ovemight exchange, the zeolite sample was carefully washed to remove occluded material and surface-adherent salt, air dried, ground to a fine powder, and stored over saturated ammonium chloride solution in order to maintain a constant high humidity. The level of ion exchange was quantified by using atomic absorption spectroscopy for silver in the filtrates. Determinations of the unit-cell compositions are subject to a 5% error. The unit-cell compositions of the samples used in this study were as follows: Ag56Y, Ag5.9Naso.lY, Ag3,1Na5~.9Y, and A ~ I S N ~ ~ ~ S Y . A two-compartment three-electrode electrochemical cell was used. The counter electrode was a 1-cm2 platinum flag separated from the working electrode by a glass frit. A platinum wire quasi-reference electrode was also employed. The modified electrodes were formed by coating clean blanks of indium tin oxide (ITO) on glass with a thin film of the modifying layer as previously de~cribed.~ Note that an equal mass of graphite powder was used in the coating. This produced larger Faradaic currents which can be ascribed to an effective increase in the surface area of the electrode. Chronoamperometry was performed using a PAR 273A (EG&G) potentiostat. During chronoamperometry,the zeolite-modified electrodes were biased 300 mV negative of the most cathodic reduction peak potential. Current decays were directly collected in digital form, and the data were fitted to eq 9 (vide infra). Each electrode was subjected to one pulse program immediately following immersion in the electrolyte solution. Electrochemical measurements were carried out in aqueous solutions of 0.1 M tetrabutylammonium fluoride (TBAF), obtained from Aldrich. The electrolyte solution also contained NaN03 at concentrations varying between 0.01 and 0.4 mM. Prior to measurements, the electrolyte solution was purged with oxygen-free nitrogen. Barnstead purified water (resistivity 18 MQ cm) was used in both ion exchange and electrochemistry.
Results and Discussion Electrochemistry of Partially Silver Exchanged Zeolite Y: Theory. We have recently shown7 that the electrochemical reduction of Ag+ at a zeolite Y-modified electrode is not diffusion controlled when the majority of the silver ions are initially located in the small cages of the zeolite. Rather, an incracrystalline ion-exchange step between Ag+ in the small cavities and mobile Na+ in the supercages is the rate-determining step, with an activation energy of 35 ( f l ) kJ mol-'. Prior to a more detailed discussion of this process it is crucial to correctly intefpret the mechanism of the electrochemical processes that occur. This is best envisaged by considering the two electron transport mechanisms originally proposed by Shaw et al.**These are reproduced below for the specific system studied in this
+ C+(s) Agf(s) + Cf(z) Ag+(s) + e- - Ago
(1)
intrazeolite electron transfer Ag+(z)
+ C+(s) + e- - Ago(z) + Cf(z)
(2)
where z and s denote the zeolite and solution phases, respectively, and C+ is the electrolyte cation involved. In (1) the electron transfer occurs subsequent to the ion-exchange step, whereas in (2) the electron transfer must stem from electron hopping or tunneling phenomena. In our work so far concerning AgY zeolite-modified electrodes, we have found no evidence for the operation of the intrazeolite electron-transport mechanism, and this is discussed in detail elsewhere.23 The modeling of the electrochemistry of AgnNa56-nY in this paper therefore is based upon the dominance of mechanism 1 above. Thus the rate of arrival of Ag+ at the electrode-solution interface is controlled by the kinetics of the ion-exchange reaction in (1). In a low-loading sample, as described in the Introduction, silver cations exclusively occupy the small-channel system of zeolite Y. The route that these small-channel silver cations take to the electrode solution interface ultimately involves diffusion through the large-channel system. However they must frst exit the small channel via intracrystalline ion exchange with mobile large-channel co-ions. This involves countennotion of silver and sodium ions through the six-rings of the sodalite cages that lead into the supercages. It is this step that is rate limiting' and can be formally written as
(3) At a modified electrode this is followed by the faster processes
and Ag',,,
+ e- - Ago
(4)
where in the above sb refers to a cation bound in the small cavities, lm refers to a mobile cation in the large channels, esi refers to the conductive electrode-solution interface, kf and kb denote the rate constants for the forward and reverse ionexchange reactions, and D is the Ag+-Na+ counterdiffusion coefficient in the large channels. These processes are depicted schematically in Figure 1. In the case of small-channel silver cations, the rate of the overall electrochemical process is controlled by (3) since the silver ions preferentially locate in the small channels for lowloading samples. The rate of the ion-exchange reaction as written above is given by
We can effectuate the analysis of this situation by making a number of simplifying assumptions. The second term in ( 5 ) concerns the back-ion-exchange reaction, which we will neglect in the remainder of this paper. At the early stages of the ionexchange reaction (vide supra) this is reasonable and also
Silver Zeolite-Modified Electrodes
. I . Phys. Chem., Vol. 99, No. 17, 1995 6631
Small Channels electrode-solution 2.0
.
1.6
-
1.2
.
0.6
.
intracrystalline ion exchange counter
Large Channel
4Uiision
-Na
+!
4:
E
5 Figure 1. A schematic representation of small- and large-channel silver cation extrazeolitic electron transport. Note that small-channel silver ions first pass into the large channel of zeolite Y whence they diffuse rapidly to the electrode-solution interface. The passage of silver ions from the small to large channels is rate determining. follows from the observation' that diffusion in the large channel is fast in comparison to intracrystalline ion exchange, that is, [Ag'lml 0. The first term in ( 5 ) contains [Na+lm],which can be taken as a constant considering the high mobility of Na+ in the largechannel system and also the electrochemical experiment samples only the early part of the ion-exchange reaction, resulting in little perturbation of the initial concentration of the cations in the zeolite. Therefore, we can reasonably approximate the ion-exchange kinetics by
0.4
-
where k = kf[Na+h,l and Ct = [Ag+sbl. Integration of (6) gives
C, = Coe-kf = Co(l - kt)
(7)
This is only true for small t (Le., at the early stages of the ion-exchange reaction). Note that Cf and COare the concentrations of Ag+ ions in the small cavities at times t and 0, respectively, and k is the rate constant of the pseudo-first-order ion-exchange reaction. The current observed at the modified electrode is determined from eq 7 in the following manner if we equate the rate of the electrochemical process to that of the rate-limiting ion-exchange step,
i = nF(-dCJdt)V
(8)
where i is the limiting current, n is the number of electrons involved in the reduction process, F is the Faraday constant, and Vis the effective volume of the small cavities of zeolite Y. The combination of eqs 6 , 7, and 8 gives i = nFkCo(l - kt)V
(9)
which can be simplified to i=a-bt
(10)
where a = nFCoM, b = nFCoVk2, and k = bta. Therefore, one would anticipate that the current observed at the Ag,Na~+,,Y-modified electrode will be linear in t at short times with an intercept a and a slope of -b. The rate constant k is equal to b/a. Note that this linear region is not expected to
t
0.0
0.2
0.4
0.6
0.6
1.0
1.2
1.4
time/sec
Figure 2. Chronoamperometry of Ag3.1Nas~gYin water at room temperature. The electrolyte solution was 0.1 M in TBAF and 0.1 mM in NaN03. The pseudo-first-orderrate constant is determined to be 0.87 S-1.
occur when the ion-exchange kinetics are diffusion controlled, where the current is proportional to t-1'2 as we have recently shown.7 Chronoamperometry of AgNas6-,Y-Modified Electrodes. Typical chronoamperometric data exhibited by Ag3 1Na52 9Ymodified electrodes are shown in Figure 2. The electrolyte was 0.1 M in TBA and 0.1 mM in NaN03. The data shown are directly reproduced from five separate electrodes, although this experiment was performed many times, giving the same results. Following application of the cathodic potential to the electrode, a linear region of i vs t can be identified (at short times) as shown on the figure. The current decays shown in the figure differ slightly in their overall magnitude, which is due to slight differences in the amount of zeolite per electrode or different effective surface areas of the electrodes. This could be accounted for by multiplying eq 9 by a constant factor, which would vary from electrode to electrode, however, as we now see that this factor does not affect the determination of the rate constant for the ion-exchange reaction. From eq 10 we recall that a = nFCoVk and b = nFCoVk2. The rate constant can thus be determined from the ratio of the intercept on the current axis and the slope of the line. This leads to a rate constant of 0.87 s-l. Note that as stated above, a constant factor used in eq 9 to account for differences in electrode composition would (and does) cancel using this method. The rate constant for ion exchange between Ag3 1Na529Y and Na+ was also determined as a function of the Na+ ion concentration in solution. According to the approximations used above (see eq 6 ) , the numerical value of k should not vary as the sodium ion concentration in the external solution is changed. This was indeed the case as shown in Table 1. Chronoamperometric data for Agl5Na54 5Y and Ag5 9Nas0 1Y were also collected. For these samples a linear region in the i vs t plot was also evident at short times. The rate constants from these data were determined to be 1.28 s-l for Agl5Na54 sY and 0.80 s-l for Ags 9Na50 1Y, respectively. Note that as the sodium ion concentration in the zeolite increases, the pseudofirst-order rate constant increases as expected (see eq 6 ) . In
Baker et al.
6638 J. Phys. Chem., Vol. 99, No. 17, 1995
4
E
\ Y
E
g k
a u
I
8 V
e8
v v
. * 4
E C
2 I.
z
1 0 . 7 mA
T 0 . 7 mA
0.0
0.4
0.8
1.2
1.6
2.0
1 dtimels-" Figure 3. A comparison between the room temperature chronoamperometry of Ag56Yand Ag3.lNa5z.gY.Top: Partially exchanged sample. Bottom: Fully exchanged sample. The data are plotted according to the Cottrell equation (diffusion control).TABLE 1: Effects of Na+ Concentration in Solution on the Pseudo-First-Order Rate Constant for Silver-Sodium Exchange in Zeolite Y at Room Temperature for Ag3.1Nas2.9Y C (Na+ in solution)/& 0.01 0.04 0.1 0.4 US-' 0.87 0.88 0.87 0.87 this case, the concentration of sodium ions in the large cages is approximated to a constant, which depends on the unit-cell composition, and will thus affect the measured pseudo-firstorder rate constant. It is of interest to note that the i vs t plots depart from linearity at both long and short ('100 ms) times. The former is adequately explained by the depletion of Ag+ from the small cavities and the breakdown of some of the assumptions used to simplify eq 5 . For example, at longer times the back-ionexchange reaction may not be negligible. The nonlinearity at short times can perhaps be explained by an "induction time" required to populate the large channels following immersion of the modified electrode in the electrolyte solution. Such an induction time has been observed p r e v i o u ~ l y albeit , ~ ~ for the much slower exchange between small-channel silver ions and large-channel divalent cations. We again emphasize that the analysis presented above is appropriate only when the electrochemical reaction is controlled by intracrystalline ion-exchange kinetics. This is only true at low loadings of silver when the small-channel system holds the
0.0
0.6 0.8 1.0 1.2 1.4 time/s Figure 4. A comparison between the room temperature chronoamperometry of Ag56Y and Ag31Na529Y. Top: Chronoamperometry of partially Ag+ ion exchanged zeolite Y. Bottom: Chronoamperometry of fully exchanged samples. Data are plotted according to intracrystalline ion-exchange kinetics (see text and eq 9). 0.2
0.4
majority of the extraframework Ag+. At higher silver loadings, the large-channel system contains a large proportion of the silver cations and the ion-exchange kinetics (and the electrochemistry) becomes diffusion controlled. This is exemplified in Figures 3 and 4 where the chronoamperometry of Ag56Y and Ag3.1Na52.9Y is compared. In Figure 3, chronoamperometric data are displayed for the partially and fully exchanged samples plotted in Cottrell form. It is evident that diffusion control is only operative for the fully exchanged sample. In Figure 4 the data are plotted according to intracrystalline ion-exchange kinetics. In this case it is evident that these kinetics hold only for the partially ion exchanged sample.
Conclusions A model has been developed that explains the chronoamperometric data recorded for partially exchanged AgNaY zeolite modified electrodes and is consistent with an electrochemical reaction controlled by ion-exchange kinetics. This allows the determination of the pseudo-first-order rate constant for the intracrystalline silver-sodium ion-exchange reaction. Chronoamperometry of partially silver ion exchanged Y zeolite modified electrodes showed good agreement with this model. This is the first time, to our knowledge, that kinetic data has been determined for a fast monovalent large-pore zeolite system. As well as aiding in our understanding of the importance of ion-exchange kinetics in the context of zeolite-modified electrodes it is potentially a useful general method for kinetic studies
J. Phys. Chem.,-Vol. 99, NO.17, 1995 6639
Silver Zeolite-Modified Electrodes of other ion-exchange materials where diffusion control is inoperative. Finally, this paper once again shows that the electron-transfer event occurs subsequent to the release of silver from the zeolite molecular sieve; that is, the electron transfer occurs outside of the zeolite pore system.
Acknowledgment. We gratefully acknowledge the Natural Sciences and Engineering Research Council of Canada and the Institute for Chemical Science and Technology for financial support of this research program. References and Notes (1) Dyer, A. Modem Theories of Ion Exchange and Ion Exchange Selectivity with Particular Reference to Zeolites. In Inorganic Ion Exchungers in Chemical Analysis; Qureshi, M.; Varshney, K. G., Eds.; CRC Press: Boca Raton, FL, 1991. (2) Sherry, H. S. The Ion-Exchange Properties of Zeolites. In Ion Exchange; Marinsky, J. A., Ed.; Marcel Dekker: New York, 1969; Vol. 2. (3) Breck. D. W. Zeolife Molecular Sieves: R. E. Kriger Publishing: M&&, FL i984. (4) Baker, M. D.; Senaratne, C.; Zhang, J. J. Chem. Soc., Faraday Trans. 1992, 88, 3187. ( 5 ) Baker, M. D.; Senaratne, C. Anal. Chem. 1992, 64, 697. (6) Senaratne, C.; Baker, M. D. J. Electroanal. Chem. 1992,332, 357. (7) Baker, M. D.; Senaratne, C.; Zhang, J. J. Phys. Chem 1994, 98, 1688.
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(8) Brooke, N. M.; Rees, L. V. C. Trans. Faraday SOC.1968,64,3383. (9) Brooke, N. M.; R e s , L. V. C. Trans. Faraday SOC.1%9,65,2728. (10) Duffy, S. C.; Rees, L. V. C. J. Chem. Soc., Faraday Trans. 1 1974, 70, 777. (11) Duffy, S. C.; Rees, L. V. C. J. Chem. SOC.,Faraday Trans. 1 1975, 71, 602. (12) Ikeda, T.; Sasaki, M.; Milkami, N.; Yasunaga, T. J. Phys. Chem. 1981,85, 3896. (13) Ikeda, T.; Sasaki, M.; Yasunaga, T. J. Colloid. Infer$aceSci.1984, 98, 192. (14) Ikeda, T.; Yasunaga, T. J. Colloid Inter$ace Sci., 1984, 99, 183. (15) Hoinkis, E.; Lecy, H. W. Proc. Inf. Con$ Ion Exch. Process Ind., London, July I969 1970, 339. (16) Bngham, E. S.; Snowden, P. T.; Kim, Y.; Mallouk, T. E. J. Phys. Chem. 1993, 97, 8650. (17) Drummond, D.; DeJonge, A.; Rees, L. V. C. J. Phys. Chem. 1983, 87, 1967. (18) Dyer, A.; Fawcett, J. M. J. Inorg. Nucl. Chem. 1966, 28, 615. (19) Dyer, A.; Gettins, R. B. J. Inorg. Nucl. Chem. 1970, 32, 2401. (20) Dahms, H. J. Phys. Chem. 1968, 72, 362. (21) Ruff, I. Electrochim. Acta 1970, 15, 1059. (22) Shaw, B. R.; Creasy, K. E.; Lanczycki, C. L.; Sargeant, J. A. J. Electrochem. SOC.1988, 135, 869. (23) Senaratne, C.; Zhang, J.; Baker, M. D.; Rolison, D. R.; Bessel, C. A. Manuscript in preparation. (24) Senaratne, C.; Baker, M. D. J. Phys. Chem. 1994, 98, 13687. JP9427 33+
