Catalytic and electrocatalytic reactions in solid oxide fuel cells

Solid State Ionics 28-30 ( 1988 ) 1521-1539 North-Holland, Amsterdam

CATALYTIC AND ELECq'ROCATALYTIC REACTIONS IN SOLID OXIDE FUEL CELLS Costas G. VAYENAS Department of Chemical Engineering, institute of Chemical Engineering and nigh Temperature Chemical Processes, University of Patras, Patras 26110, Greece Received 29 August 1987

Solid oxide electrolyte cells operating on H2 and CO as the fuel have been studied extensively in recent years and may soon provide a viable option for power generation. It has been shown recently that the same type of cells, with appropriate catalytic electrodes, can be used to carry out a number of industrially important complete or partial oxidation reactions. In this "chemical cogeneration" mode of operation valuable chemicals and electrical l~ower are produced simultaneously. Using similar solid oxide cells operating in the oxygen pump mode, it has been found recently that the rate and product selectivity of several catalytic reactions can be altered significantly. This reversible electrochemical modification of the activity and selectivity of metal catalysts is of significant theoretical and practical interest. Progress in the areas of the use of solid electrolyte cells for chemical cogeneration and for catalytic activity and selectivity enhancement is reviewed, including some very recent unpublished results. The use of solid oxide cells for mechanistic studies of catalytic phenomena is also briefly surveyed.

1. Introduction

Since the pioneering work of Kiukkola and Wagner in 1957 [ 1 ] solid-state electrochemical cells involving oxygen-ion-conducting solid electrolytes have been investigated extensively, mainly for fuel cell applications [2-141, but also for .~te~m elecuoly~is [ 5,15-19 ]. The thermodynamic efficiency of solid oxide fuel cells (SOFC) compares favorably with thermal power generation which is limited by Carnot-type constraints. Their high operating temperature offers some important advantages over other types of fuel cells, including the possibility of internal reforming. Consequently solid oxide fuel cells appear to be a promising candidate for power generation in the future. State of the art SOFC units operating on HE and CO as the fuel can maintain surface power densities of order 0.4 W / c m 2 of electrolyte ever se~,erai ~housands of hours [ 7,8 ]. This combined with some new monolithic fuel cell designs, which offer unparalleled volume power densities [913 ], could make commercialization of SOFC feasible by the early nineties [14]. So far the only commercial use of solid electrolyte cells has been as oxygen sensors, particularly for the monitoring and control of combustion processes [ 20-22 ].

In recent years some new potential applications of SOFC have emerged, i.e. their use for simultaneous coproduction of chemicals and power. This mode of operation combines the concepts of a fuel cell and of a chemical reactor. Its feasibility was first demonstrated in t 980 when ~t was shown that SOFCs with Pt based electr<~des can qua.atitativcl~ . .,..,,,~,.~"*, ..... NH3 tc NO with simultaneous generation of elecmcal power [23-25]. Subsequent work has shown that with the use of appropriate catalytic anodes several other exotherrnic reacUens of industrial importance can be carried out successfully in SOFC reactors ~ith simultaneous generation of electrical power rather than heat. These include the methanol oxidation to formaldehyde [26 ], the Andrussov process for the produc::ion of HCN [271, the oxidative dehydrogenation of ethylbenzene to styrene [28,29 ] and 1butene :e butadiene [ 30 ] and the direct oxidation of H2b tO SC~a [ 3 i ]. in severa~ of these reactions the value of the electrical energy byprodact is comparable te the value of the chemical product 'self. Consequently this type of fuel cell reactor could become very attractive as soon as commercial SOFC units become available. Parallel work in the area of aqueous electrochemistry has been reviewed recently [321.

0 167-2738/88/$ 03.50 © Elsevier Science Publishers B.V. ( North-Holland Physics Publishing Division )

1522

C. G. Vayenas /Solide oxide fuel cells

An equally promising recent development in this area has been the discovery that SOFC operating in the oxygen "pump" mode can be used to enhance the rate and the selectivity of several important catalytic reactions. This was first demonstrated with ~ e reaction of NO decomposition [ 33,34 ] and more recently with the cases of ethylene and propylene epoxidation [35-37], methane eovversion to C2 hydrocarbons [38,39], propylene oxidation to acrolein [40], methanol conversion to formaldehyde [ 26 ] and CO hydrogena,:on [ 41,42 ] and oxidation [431. A very interesting and somehow suprising aspect of many of these studies is that the observed phenomena are non-Faraaaie in the sense that the observed enhancement in the rates of the partial or complete oxidation on the catalyst-electrode exposed to fuel,-oxygen mixtures can exceed the rate of O 2- traasport to or from the catalyst electrode by as much as a factor of 500. In these eases O 2- pumping causes dramatic and reversible changes to the catalyric properties of the catalyst-electrode. This in situ reversible modification of the catalytic properties of metal and, possibly, metal oxide catalysts is potentially of great practical importance, h results from the unique possibility offered by solid oxide electrolytes to apply electrochemical techniques at the operating temi~crature of man}' important catalytic processes. Because of their abi!i~y to operate at temperatures of catalytic interest, solid electrolyte cells offer several possibilities for the study of the mechanism of heterogeneous catalysis by metals. These include passive potentiometric measurements of 1the activity of oxygen on working metal catalysts [44-51 ] and determination of coverages and chemisorption equilibrium constants from exchange cun'ent density measurements [ 52,30 ]. In the present review an attempt is made to summarize previous work related to the small but rapidly expanding field of electrocatalytic and catalytic phenomena in solid oxide fuel ceUs~ Some very recent unpublis/~ed results on ~he methanol partial oxidation, H2S oxidation and the CO oxidation systems are also included. The wealth of published informarion on solid oxide electrolytes and their thel cell applications with H2 and CO as the fuel been summarized in several ireviews [2-4,6].

2. The use of solid electrolyte cells for the study of heterogeneous cat~dysis The interesting role that solid electrolyte cells can play in the study of heterogeneous catalysis was first realized by C. Wagner [44 ] who proposed the use of such cells for the measurement of the activity of oxygen on metal and metal oxide catalysts. This technique, which is frequently called Solid Electrolyte Potentiometry (SEP), has been used in conjunction with kinetic measurements to study the mechanism of several catalytic oxidations in metals [45-51 ]. The technique is particularly suitable tbr the study of oscillatory catalytic reactions. Fig. I shows a typical experimental arrangement for the use of SEP. The bottom porous metal electrode of the solid electrolyte cell is exposed to ambient air and serves as a reference electrode. The top electrode serves both as an electrode and as a catalyst for the catalytic reaction to be studied. The open circuit EMF E of the solid electrolyte cell is given by

(1)

E= ( R T /4F) [#02-P02., ] ,

where #o., and go2.r are the chemical potential of oxygen at the catalyst and reference electrode, respectively. Eq. ( 1 ) is derived on the assumption that the electrolyte is a pure O 2- conductor and also that the u#ra~orr

°°"°'

./P~ lead wire

. llHN-

........

cooling coils ~ I

zirconiu yttria-- .

t lead wire

tube

l

Pt reference electrode

I

mac~r inc,ulator

Fig 1. Schematic of ~olid electrolyte cell used for simul:aneous kinetic and SEP meas ~remcnts.

C.G. Vayenas/Solide 9xide fuel cells

dominant exchange current reaction at both electrodes is 0 (ad) + 2 e - m 0 ~-

1523

Pco(feed), m-3bor

%c02

9z,

.03

E, mV

(2a)

or

O= (ad) + 4e- ~ 202- ,

(2b)

where O (ad) and O2 (ad) denote oxygen which is dissociatively or molecularly adsorbed on the electrode at the three phase boundary between the electrode, *.heelectrolyte and the gas phase [ 45-53 ]. The latter assumption is certainly valid for the reference electrode, but may not be always valid for the catalyst electrode under fuel rich conditions if the fuel molecules adsorb strongly on the catalyst surface. In this case, e.g. in the CO oxidation on Pt under fuel rich conditions, it is possible to obtain mixed potentials, in which case the measured EMF E provides only a qualitative measure ,,f *~c " " ~ " ' . ' ' * .,a..,~,~.,,~ surface activities [50,54,55]. When both assumpt;.ons are valid, then the EMF E provides an in situ quantitative measure of the activity of oxygen adsorbed on the catalyst. The chemical potential of oxygen adsorbed on the reference electrode is given by Po~, ~=/~2,,, +RTln(0.21 ) ,

2.76 ~

.09

.........

-.5.25

(3)

one obtains E=(RT/2F)

ln[ao/(0.21 ) t / e ] .

(4)

Consequently by measuring E and T one can compute ao which expresses the square root of Che partial pressure of gaseous oxygen that would be in thermodynamic equilibrium with oxygen adsorbed an the catalyst surface, if such an equilibrium were eso tabl~shed. By comparing ~he po~entiometnca~ty o:brained surface oxygen activity a~ with the independently measured gas phase oxygen activity ~ . it is possible to extrac~ useful information about the rate limiting step of the catalytic reaction. Thus if a~,=Po: then the oxygen chemisorption step is in

250

-700

III11 llll[![lll!ln

6.81

~

~

26

-300

.17

-Z,50

.07

750 800

(2)

where Po_,,,, o is the standard cherrfical potential of gaseous oxygen at the temperature T, and ('.21 (arm) corresponds to the gaseous oxygen activity in the reference gas, i.e. air. Defining the activity of atomic oxygen adsorbed on the catalyst ao from

-350 -650

13. 2

o Po2 =/~o2~,, + RTlna 2

.08

d

2m~

Fig. 2. Effect of inlet CO partial pressure on the reaction rate and EMF oscillationa observed daring CO oxida'don on a Pt ca~ai~'s~ electrode; Po, =0.054 bar, T= 610 K, total Volumetric flowrate 2.7X 10-4mole/s. (ref. [55,69]).

equilibrium and cannot be rate limiting. If however a 2 < P o , , as is often found to be the case [4548,51,55], then oxygen adsorption is rate limiting. If ao2 =Po2, then eq.(4) reduces to the usual Nemst equation E=(RT/4F)

ln(Po:/0.2] ),

(5)

which is valid when gaseous and adsorbed oxygen are ~ equilibrium, tt should be noticed that ao obtained from eq. (4) is measured on the catalyst surface a'~ the three phase boundary between the electrol)~e, the electrode and the gas phase. Consequently ao can give useful information about the reaction mechanism provided it is uniform over the entire catalyst surface, i.e. provided there are no sig-

1524

C G. Vaye,~as/Solide oxide fuet ceils

nificant surface concentration gradients on the catalyst surface. To this end it is important to use thin (e.g. 5-10 gm) ~orous ,~atalyst films in order to avoid the creation of gaseous and consequently surface concentration gradients inside the porous catalyst film [46,48 ]. Despite its limitations, SEP is one of the very few techniques that can be used to extract in situ information about adsorbed species on catalyst surfaces at practically important pressure and temperature conditions.. It is particularly useful for the study of oscillato~ catalytic reactions. Fig. 2 shows typical rate and EMF oscillations obtained during CO oxidation on a Pt catalyst electrode [ 55 ]. It can be seen that E, and consequently no, oscillate in phase with the rate of the reaction, which is monitored by an IR CO2 analyzer, and that increasing rate corresponds to increasing ao during the oscillations. Fig. 3 shows the dependence of the frequency and amplitude of the rate and EMF oscillations on no. It can be seen that the high frequency transition (bifurcation) between oscillatory and non-oscillatory states occurs near the decomposition pressure of surface PtOz which is given by lna~ = 12.5- 12000/T.

(6)

This dissociation pressure expression has been oh-

tained from similar experiments during ethylene oxidation on Pt [48,56,57 ], which is also an oscillatory reaction. The expression is in good agreement with independent high precision resistance measurements [58 ] and more recent XPS measurements, which have positively confirmed the existence of surface PtO2 [ 59,60]. The information offered by SEP, i.e. that at the transition between oscillatory and non-oscillatory states during C2H4 and CO oxidation ao is very near the dissociation pressure of surface PtO2 led to the conclusion that surface PtO2 plays a key role to the origin of the observed oscillations under atmospheric pressure conditions [48,56,61 ]. This has led to mathematical models which describe the observed oscillatory phenomena in a semiquantitative manner [ 56,61 ]. By using SEP another interesting problem related to catalytic oscillations has been recently resolved, i.e. the problem of communication and synchroniration of catalyst crystallites during oscillations [ 55 ]. The question, which has been addressed by numerous researchers in the field of oscillatory heterogeneous reactions is the following [62]: When a supported catalyst exhibits self-sustained oscillations of the type shown here, how do the trillions of catalyst crystaUites dispersed on the inert support

I

I

I

~i02 decompos#~on. essure, ref.(54J

A

0

_,-,=~ I -30

n

..~El

-20

- 10

I 0

in 0o Fig. 3. Eff::ct of ao on the frequency, and amplitude of ao oscillations during CO oxidation on P'~. Open and dark symbols cerr.espon4 m oscillations obtained on a preoxidized and prereduced Pt catalyst electrode respectively; T= 610 K. (refs. [ 55,69] ).

C.G. Vayenas/Sofide oxide fuel cells

1525

.24 %C02

( ~

..... 100 Et, mV .... -200 --" -100 E2.mV .200

elecfrodd

~ •.~:;v.::.:.-.,'....v~ .,-:~.--:.-.'~.~:i:~ •

electrode (a)

i lminl

electrodeelectrode (b)

Fig. 4. Electrode configuration for monitoring the EMF on two catalyst films exposed to the same reactor environment (a) and for io measurements (b). (refs. [55. 69] ). communicate and synchronize with each other to jointly I;roduce the isothermal rate spikes and oscillations of the type shown in fig. 2? Is it thermal communication through the support or is it communication by fast mass transfer via the gas phase? T . e answer appears to be the ~atter according to the experiment shown in fig. 5. In this experiment two porous Pt catalyst films are placed in the same reactor and SEP is used to simultaneously measure the oxygen activity on each of them. It can be seen that the EMF oscillations are synchronous among themselves and also synchronous with the independently monitored CO oxidation reaction rate. Since the characteristic time for heat conduction through the zirconia between the two films is at least a factor of 10 longer than the observed rate and EMF relaxation times [ 55 ], it can be safely concluded that catalyst communication and synchronisation occurs via mass transfer through the gas phase. Since this is the case with two separated Pt films, it is reasonable to expect that the same conclusions hold when triHion~ of separated Pt crystallites are involved in a commercial suppo~,ed catalyst. Another interesting technique tbr the study of chemisorption on metals is the one recently proposed by Wang and Nowick r52] and used by Manton [ 30 ] to study oxygen chemisorption on Pt, In this technique the catalyst electrode is exposed to the chemisorbing gas and the three-electrode system shown in fig. 4b is used to measure the exchange cur-

.~5

~ ElI

vst 'cotal elecfrode ],so],rid .,)

r'eference electrode Fig. 5. Demonstration of catalyst crystallitecommunication and synchrony using SEP on two Pt catalyst films used for CO oxidation in the same reactor. The EMF oscillations are also syncl-ono"~ to the rate oscillations; Pco=4.9Xl0 -~ bar, Po:=4.2X 10-2 bar, T=556 K. (refs. [55,69] ). rent density io at the membelectrolyte-gas interface. Assuming Langmuir ~ype chemisorp~ion i~ cap, be sho,~n [ 52 ] that io = L . 0 " ( 1 - 0 )

'~ ,

(7)

where 0 is the catalyst electrode surface coverage, Kr depends on temperature only and a = 1 or 2 for molecular and dissociative chemisorption respectively. By varying the partial pressure P of the chemisorbing gas at constant temperature one obtains the value Pma.~ for which io is maximized. According to eq. (7), Pm.x corresponds to @= !/2. From the definition of the Langmuir ismhern~, i.e.

,.
(S )

it follows that K = P -x when O= 1/2. By repeating the measurements at various temperatures one can determine K ( T ) and thereby extract the neat anO standard entropy of adsorption of the chemisorbing gas. h appears feasible to generalize the approach so that the Langmuir adsorption assumption can be

1526

C.G. Vayenas/Solide oxide fuel cells

avoided. The approach has been used already for the case of oxygen chemisorption on Pt [ 52,30 ], Au and Ag [ 52 ]. Its application for the study of oxygen chemisorption on other metals appears quite promising.

The first demonstration of the feasibility of coproducing useful chemicals and electrical energy in a SOFC was for the case o f NH3 oxidation to N O in

the cell NHs, NO, No, Pt IZrO2 (8 mo!% Y203 )lPt, air

(8) 3. Chemical cogeneration

Solid electrolyte cells with appropriate catalytic electrodes can be used to simultaneously generate chemicals and electrical power. In this mode of operation, frequently called chemical cogeneration, the solid electrolyte cell serves both as a fuel cell and as a chemical reactor. There are several reactions of industrial importance with AG values comparable to that of H2 oxidation which are currently carded out commercialy using catalytic reactors which generate heat. It would be highly desirable from an environmental and from an energy conservation point of view to obtain that energy as electrical rather than thermal energy [ 63]. In some of these reactions, such as the NHa oxidation to NO and the H2S conversion to SO2, the value of the obtainable electrical energy is comparable to the value of the chemical itself [32,64]. The rather unique ability of solid electrolyte cells to coproduce chemicals and electrical power in an efficient manner stems from their high operating temperature which: (i) Decreases drastically the anodic and cathodic activation overpotendal, tha~ would very severely reduce the efficiency in a low temperature (e.g. aqueous solution) environment. (ii) >crmits the use of anodes made of catalyst materials kaown to possess high activiD and selectivity for the corresponding catalytic reactions in the same ~emperature range. As an example, it is known from the catalytic literature [65 ] that Pt or Pt-Rh catalysts catalyze effectively the industrially important NH3 oxidation to NO at temperatures between 750 K and 1150 K. Above 1150 K significant amounts of N2 byproduct are formed, while below 750 K the prising then !hat aqueous electrolyte ammcnia fuel cells w{th P~ electrodes produce only N 2 [6<~i while solid electrolyte cells with Pt based anodes produce the desired product NO with selectivity of order 97% {23-25 ], i.e. comparable to that of the comesponding industrial calalytic process.

operating at temperatures between 800 and 1100 K [ 23-25 ]. It was found that selectivity (i.e. moles NO produced/moles NHa reacted) of order 97o/o could be achieved when the ratio M of the fluxes of 0 2 (i.e. i/2F) and of NH3 to the anode is above roughly 0.5 as shown in fig. 6. The experimental observations have shown that the overall anodic reaction is 2NHa + 502--~2NO+ 3 H 2 0 + 10e-.

(9)

The elementary steps of (8) have not yet been fully elucidated. The Pt electrode also catalyzes the undesirable reaction 2NHa + 3NO-~ SN2 + 3 H 2 0 ,

(10)

which leads to N2 formation. The importance of the latter reaction can be minimized by proper desihn of the anode and proper choice of the operating c~uditions [23-25]. A mathematical model based on the above reactions (9) and (10) has been R,,md to describe the experimental results in a quantitative manner [24,25]. As shown in fig. 6 the power outI .0

~

i

i

le.3 CC STP/MIN O. 8

1131 0. S2%~'NH3

m#

i

lO

d¢

/i

o -8~ rq

:13

>-

~O.B

-6 m ° z co m

W--

-H

U e.j

-'O.H

.4-<

u~J Crl

\

-2C~

T

" o

io

~o

ao

~o

5o

CURRENT (mAHP)

Fig. 6. Cogeneration of power and NO. Effect of current on selecuvity tbr NO and power produced. E]ectrol~x~surface area 2 cm2; T-- 1131 K, Ptanode. (ref. [25]).

C G. Vayenas/Soiide oxide fuel cells

put of the first NH3 SOFCs was of order a few m W / cm 2 of electrolyte and was limited by the rather thick ( ~ l mm) electrolyte components used. Ohmic polarization was found to be the major source of overpotential. A problem of the early ammoma solxd electrolyte fuel cells was the short lifetime of the Pt

anode, i.e. a few days [25 ]. This problem has been overcome by using Pt-Rh alloy anodes [ 25,67 ], with compositions similar to those used in the industrial catalytic process. Similar to the NO producing SOFC is the HCN producing fuel cell:

1527

T=C~C XME T =5.6 .tO"2 r~ H CHO CO.CO)

¢~.

T-62S*¢ XMET..-5.7 40 -2 ,, ,.

HCHO CO.CO2

T=6S:C

/

-.

o

.

~ ,~.

HCHO

C0.C02 /

CH4, NH3, HCN, CO, N2, Pt(Rh) I ZrO2 (8 mol%

Y203 )IPt, air

/

( 11 )

/

t/o

/

which is based on the electrochemical equivalent of the Adrussow reaction, i.e. CH4 +NH3 + 3 0 2 - -*HCN+ 3 H 2 0 + 6e - .

(12)

This type of cell has been recently operated with success at temperatures 1100 to 1300 K [27]. The selectivity to HCN is moderately affected by current density and can exceed 75%, i.e. it is comparable to the selectivity obtained in the industrial catalytic Andrussow process. The cell power output was of order 8 m W / c m ' and was again limited by the electrolyte thickness. In a very recent study [26] methanol has been found to be converted almost quantitatively to fornaldehyde in the cell: CH3OH, HCHO, Agl ZrO2 (8 mol% Y 2 0 3 ) l A g , air

(13) operating at temoeratures 800 to 900 K. As shown in fig. 7 the rate of formaldehyde production matches exactly the flux of O2- reaching the anode and reacting with methanol: C H 3 O H + O 2 - - ~ H C H O + H 2 0 + 2e - .

(14)

Formation of the undesired byproducts CO and CO~ results from the catalytic CH3OH and HCHO decomposition on the silver catalyst. The rate offermarion of CO and C Q increases wilt, temperaure, current and CH3OH partial pressure trig. I~) and is very low at temperatures near 800 K, where the selectivity to formaldehyde exceeds 95%. Stabilized zirconia solid electrolyte cells can also

~. /

c

-~_--_-~-_-_T~.-- ~ - -

-

~b

~

~5

s~ mA

Fig. 7. Effect of current and temperature on the rates of HCHO and CO, CO., production. Solid line corresponds to the stoichiometric requiremem for HCHO production. E!ectre!y~c ~uffacc area 2 cm 2 (ref. [ 26 ] ).

be used to carr3 out o~idative dehydrogenation reactions. Two such systems have been already studied for the conversion of ethylbenzene to styrene [ 28,29 ] and of 1-butene to butadiene [30]. Styrene is produced industrially either by the catalytic dehydrogenation of ethylbenzene, which is endotherrnic and equilibr'um limited, or by catalytic coproduction with propylene oxide via the oxirane process [28 ]. Attempts to develop a viable catalytic oxidative dehydrogenation process have not met with success. Solid electrolyte cells offer an attractive alternative, i.e. the electro,;atalytic oxidative dehy:arogenation of ethylbenzene which eliminates she thermodynamic limitations of the catalytic dehydrogenation and ac~ua!!y leads t~ powe~ c~roduc6er~ and quite satisfactor3~ selectivity to styrene [28,29]. The reaction was studied in the cell shown in fig. 9 at temperatures near 850 K using porous Pt oleore, des. Styrene, CO and CO: were found to be the major products, with lesser ~mounts o~"benzene and

C G. Vayenas/Solide oxide fuel cells

1528

~'''

I ....

I ....

I''''=~

T=600° C T=650°C

2.5 o r,,..z-1.~s o ~" 2.04s

XMET=5.1.10"2 o HCHO

=

/

a r,,^.~-1.25s

/

/

: "-

/

CO.CO2

XMET=2.9"IO"2 0=

a

HCHO

,,

CO,CO2

1.5

XMET=a6.10-2 •, ,,

/'/CHO co. co2

/

/

"~

to

0

I

.0~

Ilia

2.0

I*

i

l * , , , l * , , ,

z..O

6.0

8.0

i (A/m z ; Fig. lO. Effect of cell current on the rates of oxidative dehydrogenation (DH) and deep oxidation (DO) of ethylbenzene. (ref.

[2s]). 0

5

rnA

10

toluene also formed. The selectivity to styrene (moles styrene procued/moles ethylbenzene reacted) was ,,1,ypically of order 75%. Increasing supply of 0 2- to the anode enhances both the rate of styrene production and the rate of CO and C02 formation (deep

15

Fig. 8. Effect of current and feed methanol partial pressure on the rates of HCHO and CO, CO2, production. Solid line corresponds to the stoichiometric requirement. (ref. [ 26 ] ).

AN~'~ULAR ELECTROCHEMICAL REACTOR RADIAL CF:OSS- SECT!Ot'~

FUSED QUARTZ ROD ~

/

~

~

0.03¢m

REACTOR VOLUME : 2.13 ¢¢

FUEL -SIDE CROSS SECT iON

¢,

R E F E R E ~ A.~.t~'.

.

~lg~h~ A~DE

¢

ELECTRODELEAE)S

AI~-SO~T. C~OSS SECtiON

REFt, REF~E C & T H ~

~NQNG CATHODE

QUARTZ ROD

Fig. 9. Schematic diagram of annular s¢,"bili:ed z-rcor, ia fuel cell reactor u~ed for ~he oxidative dehydrogenation of ethylbenzene (refs. r2~291 and 1-butene (ref. [30]).

1529

C G. V a y e n a s / S o l i d e o x i d e f u e l cells

1501.-- o i= o Atrn.2

a

"-

30AI~

/;

t- -- Calculated

,2.0 F ,o.or-

In a recent study of the anodic oxidation of light hydrocarbons and alcohols in stab ]ized zirconia cells with Pt and Au, electredes, Mason and coworkers [68] found evidence of CO and formaldehyde formation dur;ng CH4 electrooxidation. However no quantitative data have been yet reported. The direct a.~odic oxidation of HzS to SO, has been recently achieved in the cell

.

.

.

.

.

.

/

/

'

/:

H2S, Sx, 8 0 2 , 8 0 3 ,

}'t.I ZrO2 ( 8 % Y 2 0 3 ) I P t , air

(16)

X~B{%) Fig. 11. Effect of inlet ethylbenzene partial pressure and cell current on the rate of styrene production. (ref. [28] ).

oxidation) as shown in figs. 10 and 11. It was found that the dehydrogenation rate on the Pt catalystelectrode can be enhanced by as much as 600% relative to the open-circuit value by moderate current densities. An interesting and practically useful finding of this study was that addition of gas phase hydrogen suppresses the deep oxidation rate more than the rate of styrene production. A two-site LangmuirHinshelwood type reaction mechanism was found to quantitatively describe the catalytic and electrocatMytic results [ 28,29 ]. Similar in scope is a recent investigation of the oxidative dehydrogenation of butene to butadiene in the solid electrolyte cell

at temperatures near 1000 K [69]. At lower temperatures and low ( < 5 mA/cm 2) current densities elemental S production was observed. This causes reversible poisoning of the Pt catalyst electrode. At high current densities ( > 10 m A / c m 2 ) this problem disappears and H2S is converted quantitatively to SO2 (fig. 12). Selectivity exceeding 95% at 95% conversion has been obtained. The direct electrochemical conversion of H2S to SO2 is of significant I

|

20 !

/

2

0

\

0

O//a , &&

//

0 0 0~/& - 00

o/.

I

C4Hs, C4H6, H2, P~(Rh)IZrC) 2 (8 tool% Y203 )l Pt, air

/

( 15 )

at temperatures near 800 K [ 30 ]. In this study the Pt-Rh alloy anode exhibited very little catalytic dehydrogenation activity under open-circuit conditions and gaseous oxygen was added to the fuel leading to selectJvi:y of order 80%. Increasing 0 2 flux to the anode caused a linear increase in the rate of deep oxidation and had very little effect on the rate of butadiene production. Consequently the effect of increasing current on selectivity to butadiene is in this case detrimental.

•/ O~ 0

A&&

,,,i 25 currem

a ~0 density, mA.'cm2

75

Fig. 12. Effect of current c~ensityon the rates of H2S consumption ( o ) and SO2 production ( A ). T= 1003 K, electrolyte surface 2 cm 2, porous Pt electrodes, total anodic molar flowrate 3.80 X 10- s mol/S, feed H2S mole fraction 5.1 × 10-3. The straight line corresponds to the stoichiometric O 2- requirement of eq. ( 16 ). The uppermost point corresponds to complete ( > 99%) H2S conversion. (refs. [ 31,691 ).

1530

C.G. Vayenas/Solide oxide fuel cells

practical interest, because of the very negative AG of the reaction and because of the technological importance of SOz for the production of H : S O , The feasibility of large scale coproduction of chemicals and power in solid electrolyte fuel cell reactors is necessarily linked with the future of classical, i.e. H2and CO consuming, solid electrolyte fuel cells. If the latter were to become commercially available, then they could, with relatively minor modifications of the anode, become attractive to the chemical industry. Although the feasibility of chemical cogeneration has been demonstrated in the laboratory for a number of important exolhermic industrial reactions there are still numerous technical problems to be investigated. In most of the chemical cogeneration studies reviewed here the surface power densities achieved so far are of the order of a few roW/era 2 of electrolyte at most. State-of-the-art SOFCs operating on H2 or CO as the fuel can maintain surface power densities of order 490 m W / c m 2 of electrolyte for many thousands of hours ~5-8 ]. The difference is not due to the AG of the oxidation reactions involved, which are, in many of the cases discussed, comparable to that of H2 oxidation, but are due to differences in electrolyte thickness, electrode morphology and conductivity, fuel partial pressure and operating conditions. It will be necessary to combine state-of-the art technology for producing thin film (40 gin) zircenia films with the anode materials discussed here te make chemical cogeneration in solid electrolyte fuel cells a commerciaUy attractive option.

4. Catalyst activity and selectivity enhancement in soli¢l electrolyte cells A common feature of the systems discussed in the previous seclion is that all the observed current induced effects on reaction rates and selectivities can be also obtaiv.ed by supplying the equivalent (i/4F) amount of gaseous oxygen to the anode. The observed changes in selectivity with celt cu~ent and potential are reproducible with gaseous supplied oxygen at the anode and can be explained by changes in surface coverages of adsorbed species by o,lassical catalytic models [24,26,27]. There is no detectable change in the intrinsic catalytic properties of the an-

o~tdic catalyst electrode as the cell potential varies. It is true that ordy in few of the above studies was the a aodic overpotential ~ measured sepmately. In these cases [ 26,29,30 ], ~ was generally found to be rather small and this could be related, as discussed below, to the constancy of the atalytie properties of the anode. In recent years a number of systems has been found where the rate and selectivity of a catalytic reaction can be altered dramatically in a solid electrolyte cell, operating in the oxygen "pump" mode [33-43,26]. Furthermore in some of these systems where both fuel and oxygen are fed to the anode [ 35-37,43 ] the observed increase or decrease Ar in the partial or complete oxidation reaction can be as much as a factor of 500 higher that the rate of 02- pumping (i/ 2F) to or from the anode. This phenomenon, which indicates a significant change in the catalytic properties of the catalyst electrode, is frequently called non-Faradaic electrocatalysis in the sense that [Arl >> i / 2 F ,

(17)

where Ar refers to the change in the rate of the catalytic reaction occuring on the catalyst electrode and not, of course, to the rate of the electrocatalytic reaction which must always obey Faraday's law. It is useful to define an enhancement factor A [43 ] from A=Ar/(i/2F)

.

(18)

There exist reactions such as the ethylene epoxidarien on Ag [ 35,36 ], '.he low temperature CO oxidation on Pt [43] and the methanol dehydrogenation to formaldehyde [ 26 ] where A values as high as 500 or as low as - 500 have been measured. Obviously to obtain ]A I values exceeding unity for an oxidation reaction, gaseous oxygen must be supplied at the anode together with the fuel. In the subsequent discussion the current i is defined to be positive when 0 2. are pumped to the catalyst electrode and negative when O 2. are pumped from the catalyst electrode. 4. i. Rale enhancemem studies

The first reaction for which a dramatic rate increase was observed in a solid electrolyte cell was the decomposition of NO [33,34]. Mason, Hu,~ins and coworkers found that with Ft and Au electrodes the

C G. Vayenas/Solide oxide fuel cells

rate of NO decomposition can be increased by as much as three orders of magnitude relative to the open-circuit value when large negative currents are applied to a stabilized zirconia cell and the catalyst electrode potential relative to the reference electrode is more negative than - 1 . 5 V [ 33,34 ]. Similar observations have been made in water vapor reduction studies [ 16-19 ]. In the case of NO decomposition, since chemisorbed oxygen resulting from the decomposition of NO is known to severely poison the catalytic reaction rate on Pt, it would appear that oxygen removal in the form of O 2- from the catalyst electrode cleans the surface from chemisorbed oxygen and consequently enhances the rate of NO decomposition. However, the fact that similar results were obtained with a Au electrode, which was inactive for NO decomposition under open circuit conditions, led to the hypothesis that the rate enhancement is due to F-center formation on the stabilized zirconia surface itself [ 33 ]. Other explanations for this phenomenon have also been proposed, including the formation of electrocatalytieally active intcrmetallic alloys ZrPh and ZrAu3 [3 ]. Further experimental work is needed to elucidate tbe exact mechP ~;sm of this interesting electrocatalyfic system. It should be noticed that all the experimental results presented for the NO decomposition system indicate (e.g. fig. 3 in ref. [33]) that IAI ~
1531

~.mv

% c~

I l l ! ! I !

t2minl

a=b

-1500~ .-~.*--i=*9001JA.,,=

i.O

. ,,

i=-~O

IJA

. ,,

i=O

Fig. 13. Effect of O 2- pumping on the rate of CO oxidation on Pt. Solid and broken lines show the cell-reactor effluent mole fraction of CO_,, and the catalyst potential relative to the reference electrode r2,pectively. Inlet conditions: Pco = 2.44 × ! 0- 2 bar, P o : = 3 . 8 4 x l0 -2 bar, T ~ ~,2~,K. (refs. [43,69] ).

Under similar catalyst-electrode overvoltage conditions Yentekakis and vayenas [43 ] have recently lbund a dramatic enhancement in the rate of CO oxidation on Pt. Using the three-electrode ~ystem shown in fig. 4b and feeding CO-O2 mixtures at the anode it was found that up to a sixfold increase in the CO oxidation rate over the open circuit value is obtained when the iR-free potential of the porous Pt catalyst electrode relative to the reference electrode becomes mere negative than - 1.5 V. This is shown iv. fig. 13 where the corresponding enhancement factor A for curve a is of order - 500. Curve a in fig. 13 was obtained when a current of - 6 0 0 gA was first applied to the cell, after repetitive positive currents. After several successive applications of - 6 0 0 gA to the cell the rate increase stabilized to the one shown by curve b. This dramatic effect was attributed to an enhancement of the CO disproportionation rate followed by fast combustion of carbon formed on the catalyst-electrode by gaseous oxygen. This route of CO2 formation is not me,r m u u y u a m ~ n y ~',mam~ under open circuit condiliens where C©a formation is known ~e proceed en Pt by reaction k egween chemisarbed CO and oxygen [55,62]. This interpretation is corroborated by following the COn formation rate during gatvanostafic transients [43]. Interesting non-Faradaic behavior is also observed upon application of positive currents [43]. When

15 3 2

C G. Vayenas/Solide oxiae fuel cells

l ~ / . - - - - - ~ /"~"~ , /

<>T=31.7°C

e T=Id2°C

/I" o ~ ~ _ ~ - . . .

o r--~s~oc

soc v r.-s T:585°C

'~~~6¢//~v

,,

, ~

/, /.

-t

/ -02

/

/

/

/

A-IO_.

/

/ I 0

.02

( i / 2 F ) / r o , Dim/less

Fig. 14+ Steady state effect of 02- pumping -n *+he~ ? el-CO oxidation for small positive and negativeapplied currents. Broken lines are constant enhancement factor lines. The rates of CO oxidation reaction and 02- pumping are nondimensionalized with respect to the open-circuit reaction rate to; Pco=4.7 × 10bar, Po.=0.16 bar. (refs. [43,69] ). 0 2- are pumped to the catalyst electrode the rate generally increases and enhancement factors A exceeding 100 are observed. Small negative currents have generally a smaller decreasing effect on reaction rate (fig. 14). The CO oxidation on P~ is well known to exhibit self-sustained rate oscillations [49,55,62]. It has been found [431 that rate and EMF oscillations can be induced or stopped at will by appropriate application of O ~- currents (fig. 15 ). The waveform of the oscillations is not altered by 0 2 - pumping, but the oscillation frequency has been found to depend linearly on the applied current, (fig. 15). These observations can be interpreted [43 ] within the framework of existing mathematical models which describe successfully the observed o scillatory phenomena during ethylene [ 56] and CO [61 ] oxidation on Pt. In the underlying physical models the oscillations are caused by formation and consumptien of surface PtO2 [ 56,61 ]. Some similar effects have been observed recently with the cell CH3OH, HCHO, Agl ZrO~ (8 mol% Y203 )iAg, a~r, (19) which has beee s+~cl;~ fer the coproduction of ~brmaldehyde and electrical power. As previoas!y discussed when the cell operates in the fuel cell mode and produces curre~+~ which is poshive according to

the convention adopted here, then the effects observed are Faradaic (e.g. fig. 7). However when negative currents are applied to ~hc cell and O ~- are pumped from the catalyst electrode, then a significant non-Faradiac enhancement is observed in the rate of HCHO production as well as on the rate of CO and H2 byproduct formation [26 ]. This rate enhancement, which appears to have a small effect on product selectivity and corresponds to A values as negative as - 10, (fig. 16) is again very pronounced when the IR-free catalyst electrode potential with respect to the reference electrode is more negative than - 1 . 5 V. At the present stage, despite some apparent similarities in the above systems, the available experimental information does not permit any unambiguous conclusions to be drawn yet on whether the observed increases m the rates of NO and H20 decomposition [ 17-19,33,34] CO hydrogenation [41,42 ], CO oxidation [ 43 ] and CH3OH dehydrogenation [26 ] upon negative current application near the zirconia blackening conditions are of common origin, related to the creation of catalytic sites on the zirconia itself, or if the observed effects are electrode-specific and result from changes in the catalyric properties of the catalyst-electrode due to ~he very low surface oxygen activity conditions created upon 0 2- removal from the catalyst electrode. It is conceivable [ 19] that the catalytic effects observed under large negative current application, near the zirconia blackening conditions, are the sum of the effect caused by changes in the catalyst electrode beha~50r and of the effect caused by the zirconia itself. It is clear that further experimental work is needed to clarify the origin of the observed phenomena which are of significant theoretical but also practical impertance. Far from the zirconia blackening conditions, as for example in the case of the O 2- pumping effect on the Pt catalyzed oscillatory. CO oxidation the observed effects are definitely electrode specific. [40] have found that application of positive currents to a stabilized zirconia cell with Au electrodes leads to partial oxidation of propylene to acrolein at temperatures near 720 K. The selectivity to acrolein is 60% at 1% propylene conversion. The ma'n byproduc~s are CO and CO2. These results are quite interesting, since the Au catalyst electrode was found

C G. Vayenas/Solide oxide fuel cells

,I- ~ = - - - - . - - . . ~ - - - t = - 2 O O p A

.....

----tO0---

, . ~

~f •i--O

1533

~ - o300

Vw~'mV

..... II

. --=-.----i=-tO0

.---400 ~

, ,.-851

..... .~,@ @ T=332°C 0 r=297°C

"p 3

:_::~-.~ ~-'~50

I

"500

0

i, pA

"500

*~50

Fig. 15. Effect of applied current on the frequency of rate and EMF oscillations of the CO oxidation on Pt. Black circles e i;c freque~c.v versus current diagram represent the oscillator5, states of this figure. Broken lines on the same diagram are model predictions (ref. [43 ]" inlet c o m p o s i t i o n : / ' c o = 4.7 X 10- 3 bar, Pc,:= 0.16 bar, 7"= 605 K. (refs. [ 43,69 ] ).

C. G. Vayenas/Solide oxide fuel cells

1534

l

~

i

=

O: i=f O: i=2 .P 0

I ,,'0"""" I,"

/,S s s .

os / o

p"

.

6

s

•

s

0

0

O~ ~ j

00,,"

s

s

s

S

s

S

S S S

o

#/

/

01

/

% tt

0

V

,t/

I

ft

,d~"

$4"

J

s S 0

I

I

100

20O

I

-

300

AV, mV ...

t

~4

0

Fig. 17. Overpotential (AV), gaseous composition (PET/Po2) and catalyst oxygen uptake effect on the rates of ethylene epoxidation ( i = 1, o ) and deep oxidation ( i = 2 , • ): r,o are the open-circuit rate values: T = 6 7 3 K. (ref. [36] ).

-i

-b

.~mA

-b

-~

Fig. 16. Effect of applied current on the rates of production of HCHO ( o ) and CO ( o ) on a porous Ag catalyst electrode exposed to CH3OH. Solid line corresponds to A = - 1; ( E] ) 856 K, PMET=5.1×I0 -2 bar, (o) 896 K, PMET=5.0× 10 -2, ( A ) 939 K, PMET=4.6× 10 -2 bar, ( ~ ) 975 K, PMET=3.6X 10 -2 bar. (ref. [26]).

composition, electrode surface area and applied current on the transient and steady state rates of ethylene epoxidadon and deep oxidation to CO2 wei-e studied in detail [35,36]. As shown in fig. 17 it was found that the relative increase in the rates of epoxidadon and deep oxidation are proportional to the .t5

to bc total!y inactive when exposeG to propylene-oxygen mixtures under open-circuit conditions [40].

T : 400%

,

PO 2 PET ---7

4.2. Selectivity enhancement sf udies T

The first application of sol: 5 electrot, te cells operating in the o×'dgen pump mode to ai'~er d:~ product selectivity of a heterogeneous catalytic reaction was in the case of ethylene oxidation on porous Ag electrodes [ 35,36 ]. Using the cell

C2H4, C e k ] 4 U ,

(k.~U2, AglZrO2 (~z2Q} 3 ) l A g , a i r

Rc~....

.10 E

o

J f

0

.05

....~ I I ~ R c 2

_

RCIO

..~, .I/'~[.~-RC2

RC6/ "~I F ~

I~RC 2

~'RC II

I

(20) at temperatures between 580 and 700 K it was found that 0 2 - pumping to the catalyst causes an increase to the rate and selectivity of ethylene oxide production. The opposite effect is found when negative currents are applied to the cell. The effect of gaseous

L

I

0

.05

I____ .10

,15

I .20

i / 4 F Q , min -I Fig. 18. Effect of Ag catalyst-electrode surface area (measured as the reactive oxygen catalyst uptake Q) on the cell ~'elaxation time constant during galvanostatic transients. Each data point is a different catalyst-electrode. Conditions: T= 673 K, PO/PET=7. (ref. [36]).

C.G. Vayenas/Solide oxide fuel ceils ''l

I

T : 400 "~O

O'7 f

1535

t

*C

P =16.~£) 3 bor

0.6-

---

0.5-

ARt

d O

0.4-

a £,, "

/

0.3-


/

/

&R 2 Ir2 o

/

o/ /

0.~-

.

/ /

o

I-

"~ 1,7
Irto

~..O. =.= ----.--=- ..O-=-.,== =-,

(%]

0.I 0 0

I, 5O

If' III | I'II

I

~ ,~i

100

' |

'-

"

150

i ,

, ,,u

2 O0

tpA Fig. ! 9. Effect of current on the steady-state increase in the rate of ethylene epoxidation r~ and deep oxidation r2. Comparison with the rate of oxygen pumping to the catalyst. Intrinsic ( i = 0 ) selectivity ~ 0.5; PET= t.6-10-: bar, Po2 = 0.1 bar. (ref. [ 3 5 ] ).

cell overpotential and inversely proportional to the surface area of the Ag cataiyst eiec~rode. ~-he reiaxation time constant of the Ag catalyst electrode during galvanostatic transients was found to l:,e proportional to the electrode surface area and irversely prooortional to the applied current (fig. 18 ). All the observations were interpreted in terms of a simple mechanistic model which postulates the creation and destruction of an active surface oxygen species by 0 2 - pumping to and from the catalyst respectively [ 35,36 ]. Enhancement factors as high as 400 were observed during this study (fig. 19). The selectivity to ethylene oxide at 673 K could be typically controlled by the rate ofO 2- pumping between 40 and 60% and the ethylene oxide yield could be varied between 1 and 4% (fig. 20). Similar results

methane to produce ethane and ethytene is possiNe •

~

........

~

ilJl i i l C ~,gll

CH4, C2H4, C: H6, CO2, 02, Ag-Bi_,03 I ZrO2 (8 molqb Y203 ) lAg, air

(2~)

operating at 973 K. Electrochemically supplied oxygen leads to higher selectivities to C2 hydrocarbons than gaseous oxygen. The observed enhancement factors A are of order unity. The selectivity to C.,. hydrocarbons varies between 67 and 20% as current increases. Qualitatively similar results have been recently reported for the cell

CH4, C2H4, C2H6, CO2, 02, Ag-MgO [

ZrO2 (8 mol% Y203 )lag, air

(22)

lt~,a

on Ag catalyst-electrodes, in this case however ~he selectivity to propy!ene oxide was low, typically of order 5% [37 ]. In recent years there have been increasing efforts to use zirconia oxygen pumps for the transformation of methane to useful chemicals. Otsuka and coworkors ['38] have found that oxidative coupling of

',vh;ch was studied a~ 973 K both with positive and negative applied curr~nts [ 39 ]. Th,~se first exploratory studies for the conversion of methane to more vaiuable chemicals in zirconia cells ar~ quite encouraging. It appears likely that with the use of s~itable catalyst electrodes improved selectivities to Cz hydrocarbons, but also to partially

C.G. Vayenas/Solide oxide fuel cells

1536 _.._._.-...--q~

I

...D.---- '-" " - - " E " .0 4 0

L /. /

~----/

/

D

.035

/

.60

d

-/

.0 3 0

0

/

"O'"

"--" ~

" ' O "="" '"

I o/1°

.56

I/

i = 50pA

I

.025

V= 270

r

"<

mV

.54 .52

.020

I

I

15

30

I

I

45

60

.50

•t , rain .02.5 J

I

I

\

\

52

I

50

i = -- 5 0 , u A

\

V= - 2 7 0

mV

-.48

0 \\ ,0 20

-.46

\~.0

-- \

0 []

\ 13 ,,J W

- 44

N

\ k\

.015

.O;O

-

trl rq rpl cl .4 <

-.42

x \

-4 -<

\

X

i

I

15

30

45

Fig. 20. Transient .:ff ect of oxygen pumping to (a) and from (b) a Ag cata!yst electrode on selectivity and y:ield o f ethylene oxide. Current applied at t = 0 : ?":~ 673 K. (ref. [35] ).

oxygenated prodacts (CH3OH, HCHO) can be achieved. This could be of" significant practicai importance.

S. Remoter coasideratiens

The practical utility of solid electrolyte cells op-

C.G. J'?o,enas,So!ide oxide fuel cells

erating as fuel cells, with or without chemical cogeneration, or as oxygen pumps for the production of chemicals will depend strongly on the development of efficient and relatively iaexpensive reactor configurations. The Westinghouse parallel tube designs [5,7,8] have been tested for many years and it appears that all the related materials compatibility problems have been solved. It has been reported that mass production may start within the next few years [ 14]. A disadvantage of the parallel tube design is its low electrolyte sa,rfaee to volume ratio which results from the tubular geometry and is of order 1 c m - ~ in the state-of-the-art Westinghouse fuel cells. In order to overcome this limitation several novel monolithic flat plate designs have been recently proposed [ 913 ]. These designs can offer electrolyte surface to volume ratios of order at least 6 era-~, as already demonstrated by small prototype cross-flow monolithic fuel cell reactors of the type shown in fig. 21 [10]. This very high electrolyte surface to reactor volume ratio combined with the state-of-the-art power densities of 0.4 W/cm z, obtained in tubular designs [ 7,8 ], could icad ~o voiume power densities well in excess of 1 kW/1. Mathematical models have been already developed to describe the behavior o~ monolithic crosc-flow fuel cell reactors [9]. Substantial experimental progress in this direction has been recently reported by researchers in the Argonne National Laboratory [ 1 l, 12 ].

;£ .

AIR

~I . . . . . . .

. . . . .

" FUEL

] - - 7•

_._.i! o= L FUEL ; f._ ,

mmmmm a

_

..L.. _

A',R

L. F U E L

-1

t

i

b

Fig. 21. Schematicdiagram and equivalent circuit of a cross-flow monolith fuelcellreactor with seriesconnectionof unit cells. (Ref. [lO]).

1537

6. ConcRudin~ remarks In addition to their usual application as oxygen sensors and as power producing d e ices, solid oxide electrolyte cells offer several other au:'active possibilities which are related to the field of heterogeneous catalysis. This is due to the permissible temperature range of sol,.'d oxide cells which coincides with the operating temperature range of many of the important hete~'ogeneous catalytic processes. These new emerging applications of solid oxide cells include the cogeneration of power and chemicals and their use in a passive or active mode to study and to influence the rate and selectivity of catalytii? reactions on metals and, possibly, metal oxides. The existing literature has shown that several important complete or partial oxidation reactions can be carried out in solid oxide cells with high selectivity to the desired products. Some very interesting effects of cell current and overpotential on catalytic reaction rates and selectivities have been observed. There is certainly a lot more to be found in the near future. The influence and control of the rate and selectivity of heterogeneous cata]y*Acreactions by ciassical electrochemical techniques appears quite mtractive and coald lead to proeuct selectivities unavailable under open circuit conditions, i.e. unavailable by classical heterogeneous catalysts. The scope of most previous studies in this area has been macroscopic and exploratory. There is a need for more fundamental studies [70,71]) which will elucidate the mechanism of the electrochemical and catalytic steps occuring at the cawAyst electrode-electrolyte-gas phase interface, but also on the catalyst electx,)de surface and possibly on the electrolyte surface i~seV The use of the surface spectroscopic techniques employed in fundamental catalytic research studies could give verb, valuable inforrnmion. Many of the resu!!s su_n,eyed here are definiten_y ~l~o~md~-~pecific, but the role of the electrob, ~ itseV may be significant in some cases and needs Cu~her investigation. In practically all the chemical cogeneratien and v-:ygen pump studies reviewed here, yttria-stabilized zirconia was used as the solid electrolyte. In many cases (e.g. [28,29,36,37] ) Qis dictated cell operating temperatures higher man the optimai one~: for product selectivity m~ximization. Consequently the

1538

C. G. Vayenas/Solide oxide fuel cells

use of other solid electrolytes, such as the Bi:O3Fr:O~ ~n!id soh!fi.ons which exhibit higher icmic ecmductivity [ 72 ] and lead to lower electrode resistance [ 73 ] may prove quite advantageous in the future.

Acknowledgement Partial financial support by the VW Stiftung (Fed. Rep. of Germany) is gratefully acknowledged.

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