- Email: [email protected]
Binary electrocatalysts for organic oxidations
J. Electroanal. Chem., 70 (1976) 73--86
73
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
BINARY ELECTROCATALYSTS FOR ORGANIC OXIDATIONS
D.F.A. KOCH, D.A.J. RAND and R. WOODS CSIRO Division o f Mineral Chemistry, P.O. Box 124, Port Melbourne, Victoria 3207 (Australia) *
(Received 30th September 1975)
ABSTRACT The electro-oxidation of formaldehyde and methanol has been studied on a number of binary platinum electrocatalysts. These comprised mixed electro-deposits of Pt with Sb, As, Bi, Hg, Re, Te or Sn and a range of homogeneous Pt--Rh alloys of different, known, surface composition. These systems were found to exhibit an enhanced activity over that of platinum alone, and this behaviour was correlated with the eas e of adsorption of oxygen on the added metal. The activities for organic oxidation were compared with predictions of a model involving reaction between adsorbed oxygen and organic species on the metal surface. The proposed mechanism accounts for the behaviour of both homogeneous and heterogeneous alloy systems.
INTRODUCTION H y d r o c a r b o n s , c a r b o n m o n o x i d e a n d water-soluble organic c o m p o u n d s such as m e t h a n o l , f o r m a l d e h y d e a n d f o r m i c acid have b e e n c o n s i d e r e d as rea c t a n t s in fuel cell systems. P l a t i n u m has b e e n f o u n d t o be t h e m o s t active single m e t a l f o r t h e c o m p l e t e e l e c t r o - o x i d a t i o n o f these c o m p o u n d s t o c a r b o n d i o x i d e in b o t h acid a n d alkaline solutions. It is well r e c o g n i z e d t h a t these fuels (e.g. e t h a n e [1], e t h y l e n e [2], c a r b o n m o n o x i d e [3], m e t h a n o l [4,5], f o r m a l d e h y d e [6] a n d f o r m i c acid [7] ) f o r m s t r o n g l y a d s o r b e d i n t e r m e d i a t e s as given b y A -~ Bads + n l H + + n l e
(1)
T h e i n t e r m e d i a t e can t h e n react with an o x y g e n species derived f r o m water, H 2 0 ~ OHad s + H + + e
(2)
Bads + OHads -~ CO2 + n2H + + n2e
(3)
The a d s o r p t i o n r e a c t i o n (1) involves d e h y d r o g e n a t i o n o f t h e o r g a n i c species; B~d~ f r o m m e t h a n o l h a v i n g been identified as - - C O H [4,8,9] or - - C O [5, * All correspondence concerning this paper to be addressed to The Chief, CSIRO Division of Mineral Chemistry, at address given.
74 10,11]. Strong evidence has been presented [10--12] to support the conclusion that the intermediate has the same constitution in the oxidation of methanol, formaldehyde, formic acid and carbon monoxide. Reaction (1) is irreversible; an electrode held at low potentials can be removed from a cell containing a solution of the organic species, washed and returned to a cell containing acid alone without loss of adsorbate from the electrode surface [11,13]. In the potential region where (3) becomes significant, the surface coverage of Bad s is determined by a balance between its formation by (1) and its further oxidation to CO2 [5]. The existence of a parallel path, more rapid than (1)--(3), has been established for the oxidation of formic acid [12]. The additional route proceeds through a weakly bonded intermediate and is inhibited by the formation of the strongly bonded species. There is continuing controversy regarding whether such a mechanism also applies to the oxidation of methanol in acid solution. The formation of formaldehyde and formic acid in solution during methanol oxidation has been considered [14] to be evidence in support of a parallel path mechanism. However, reaction (1) is n o t a single step since its kinetics [4,5] obey the rate equation for a single electron transfer process, whereas nl is 3 or 4 depending on whether Bad s is --COH or --CO. Hence, the side products could be derived from precursor adsorbates produced by consecutive removal of hydrogen atoms from methanol [15]. Supporting evidence in favour of this suggestion is obtained from electrochemical [16] and radiochemical investigations [17] which have failed to detect any intermediates other than those which result in the formation of the strongly bonded species. Experimental evidence [13,18] which suggests that the steady state rate of methanol oxidation is more than four times faster than the rate of oxidation of the adsorbed layer, determined after removal of the alcohol from the cell solution, lends support to a parallel path mechanism. Other workers [14,16], however, have found the t w o rates to be identical as must be the case if the predominant oxidation route is (1)--(3). Capon and Parsons [12] compared the oxidation of methanol with that of formic acid on platinum and palladium electrodes. In contrast to platinum, palladium does not form a strongly b o n d e d intermediate" from formic acid, and hence in this case oxidation does not become inhibited. The complete inactivity of palladium for methanol oxidation was considered to be strong evidence that the oxidation of this c o m p o u n d requires the participation of the strongly adsorbed intermediate, i.e. it can proceed only by (1)--(3). Investigations on the influence of the presence of other organic substances on methanol oxidation [19] also support a single reaction pathway. The mechanism of methanol oxidation on rhodium electrodes has also been shown [15] to proceed b y the mechanism (1)--(3). It has been suggested by Podlovchenko and Petukhova [20] that (1)--(3) constitutes the mechanism on activated smooth platinum under conditions which might be called "quasistatic", but that a parallel path operates in the final "static state". Biegler [ 21] has shown that the instantaneous methanol oxi-
75 dation current on an activated platinum electrode decreased rapidly as the rates of reactions (1) and (3) approach each other. This rapid decrease was followed by a slower rate of decay due to the accumulation of impurities on the electrode surface. He concluded that the current after the initial rapid decrease (i.e. the equivalent of a "quasistatic state") gave the best measure of the activity of the surface. In our experiments, this current was taken to represent the activity of Pt--Rh alloys for methanol oxidation and can be related to mechanism (1)--(3). For formaldehyde oxidation, a parallel path involving the formation of a blocking polymeric surface species has been shown [6] to occur at potentials above 0.4 V. In the work reported here, activities were compared at potentials significantly below this value, and hence activities for this system can also be related to mechanism (1)--(3). The search for improved electrocatalysts for fuel cells has shown that enhanced activity over pure platinum for organic oxidations can be obtained by the incorporation of other metals such as rhenium [22,23], tin [22,24], ruthenium [13,25,26] or osmium [27]. The improvement in performance in the presence of these metals has been attributed [2,22,26--28] to oxidation of the intermediate by oxygen-containing species which are associated with the added metal and are more reactive than O H a d s o n platinum. I t h a s been observed [29] that the adsorption of molybdate ions increases the activity of platinum, the proposed mechanism for this system involving oxidation of the fuel with adsorbed molybdate, followed by electro-oxidation of the reduced molybdate species. Recently, Watanabe and Motoo [26] proposed a mechanism to explain the enhancement of methanol and carbon monoxide oxidation by ruthenium adatoms on platinum. These authors consider the reaction to occur between the carbonaceous species adsorbed on platinum sites and oxygen adsorbed on ruthenium sites in a homogeneous alloy surface. This postulate assumes the platinum and ruthenium sites in this surface to have different adsorption properties. Although their explanation would apply to heterogeneous systems, it is not consistent with investigations of the chemisorption properties of homogeneous noble metal alloys [30--32]. The characteristics of adsorption of hydrogen and oxygen on such systems were shown to be a composite of those of the individual metals, demonstrating that the surface atoms of both metals behave in an identical manner. Hence, all sites on a homogeneous alloy surface will be expected to have the same activity for the adsorption of both oxygen and organic species. In this communication, we present the results of investigations of the activity of various binary platinum catalysts for the oxidation of methanol and formaldehyde. Two types of experiment were performed. Smooth Pt--Rh alloy electrodes gave results which were consistent with a proposed reaction model. The study was extended to the behaviour of more realistic electrodes for fuel cells consisting of co-deposits with platinum black. The Pt--Rh alloy system was chosen for investigation because the surface
76 composition can be determined in situ [30] and can be varied in a controlled manner between the pure metal limits by an electrochemical pretreatment without the occurrence of phase separation at the surface. Methanol was used here because of the extensive information already available on the oxidation mechanism: Electrodes with high surface area were prepared by co-depositing platinum with either antimony, arsenic, bismuth, mercury, rhenium, tellurium or tin. All these elements are stable in acid solutions either as the metal or as an oxide over the range of potentials where activity was measured. There is evidence that certain of these systems, e.g. Pt--Re [23] and Pt--Sn [24], are heterogeneous consisting of two solid phases. Formaldehyde was used as the organic reactant in these studies because of its relevance to a formaldehyde/air fuel cell developed in these laboratories [33,34]. The experimental results for both types of electrode are compared with the predictions of a model for both homogeneous and heterogeneous alloy surfaces in which the second component metal results in a more active adsorbed oxygen species. EXPERIMENTAL AND RESULTS Apparatus
Experiments were carried out in three-compartment glass cells, the working compartment of each cell having a side arm and an outlet tube to allow the addition and removal of solutions. Cell electrolytes consisted of B.D.H. Aristar methanol or B.D.H. AnalaR formaldehyde in a supporting electrolyte, 0.5 or 1 M H2SO4, prepared from B.D.H. A n a l a g reagent and doubly distilled water. Solutions were deoxygenated with purified [35] nitrogen. A mercury/mercurous sulphate, 1 M H2SO 4 reference electrode was used, the potential of which was 0.68 V vs. the reversible hydrogen electrode (RHE) in 1 M H2SO 4. All potentials in this paper refer to the RHE. Measurements were carried out at 25°C unless otherwise reported. Electrode potentials were controlled with a potentiostat programmed by means of potential step and linear sweep generators constructed in these laboratories. Current--potential curves were recorded on a Hewlett-Packard 7004A XY recorder. A time base Type 17172A (Hewlett-l~ackard) was used in conjunction with the XY recorder for displaying current--time traces. Ele c tro de p repa ration
Platinum, rhodium and platinum--rhodium alloy electrodes consisted of 26 gauge wire of 99.99% purity sealed into soft glass tubing. The real surface areas of platinum and rhodium electrodes were determined by measuring the hydrogen adsorbed from 1 M H2SO 4 during a cathodic potential sweep at 40 mV s-1. The electrode surfaces were first cleaned by the application of
77 triangular potential cycles between 0.07 and 1.54 V until a voltammogram of reproducible shape was obtained (~ 20 cycles). The procedure used for integrating the curves has been reported previously [35,36]. Real electrode areas were calculated on the basis that the fractional hydrogen coverage at the lower integration limit was 0.77 and 0.59 and the corresponding monolayer charge 210 and 221 pC for platinum [36] and rhodium [35] respectively. For computing the area of alloy electrodes, it was assumed that these values changed linearly with alloy composition. The preferential dissolution of rhodium which is found [30] to occur during potential cycling of platinum--rhodium alloys provides an excellent method for preparing alloys with different surface compositions. In the work reported here, platinum--rhodium alloys were prepared by subjecting an alloy containing 74% rhodium * to repetitive potential cycling at 40 mV s-1 between 0.07 and 1.54 V in 1 M H2SO 4 until the surface composition reached the required value. This was determined using the linear relationship which has been established [30] between surface composition and the potential of the oxygen desorption peak on a voltammogram. After obtaining an alloy of the desired composition, the anodic limit of the following potentia} sweep was lowered to 0.80 V and the real area of the electrode determined from the hydrogen adsorption charge. The lower anodic limit was necessary for the areadetermining sweep, as the oxygen reduction and hydrogen adsorption regions overlap on sweeps from potentials > 1 . 0 V. High surface area electrolytic co-deposits of platinum with another element were prepared by plating for 30 min at a controlled potential. This was --0.15 V except in the case of p l a t i n u m - - a n t i m o n y co-deposits where a deposition potential of 0.0 V was used. The electrolyte usually consisted of 1% chloroplatinic acid in 1 M HC1 solution to which was added the other element in the form shown in Table 1. By varying the pH of the solution and the a m o u n t of the second element present, the deposit composition was controlled at approximately 4% of the added metal. The deposits were formed as finely-divided "blacks". Platinum, tantalum and iridium foils of 1 cm 2 geometric area were tested as substrates, the latter being f o u n d to be the best choice as it was not attacked during dissolution of the deposit in aqua regia for analysis. Mercury, platinum, rhenium and tellurium analyses were made using spectrophotometric methods while antimony, arsenic, bismuth and tin were determined polarographically. After preparation of the deposit, the electrode was washed thoroughly with distilled water and transferred to a second cell in which the real surface area was determined using the hydrogen adsorption m e t h o d described above for platinum, the electrode being cycled at 20 mV s-1 between 0 and 0.6 V in 0.5 M H2SO 4. The cathodic charge in the hydrogen region for platinum-arsenic and platinum--tin deposits was found to be less than the anodic, in-
* Alloy compositions are expressed in this paper as atomic percentages.
78
I d 4 d 4 4 4 4
II
o o
o o
o o 0
o
0 °~
0
IIIII
79
dicating a contribution from an anodic process other than hydrogen oxidation. For these deposits, areas were obtained from the mean value of the anodic and cathodic charges. After determination of the real surface area of each electrode, the contents of the cell were blown out with nitrogen, the cell washed with supporting electrolyte and the methanol or formaldehyde test solution added. Activity measurements The potential program used for activity measurements on platinum, rhodium and platinum--rhodium alloy electrodes in methanol solutions is shown in Fig. 1. Any adsorbed impurities or methanol residues are removed by oxidative desorption during the first step at 1.5 V. At this potential the electrode surface is covered with chemisorbed oxygen and any oxygen evolved is removed by bubbling nitrogen through the solution. The surface layer of adsorbed oxygen is reduced rapidly during the short step at 0.10 V, after which the final potential, ¢, is applied and the current--time curve recorded. The current transient was similar to t h a t observed previously by Biegler [ 21] and
ZOO
I'5V
~ : 0"65V ~,
150
Is
E
~too I--
=O'G3V 50
oLL___c - ~ =o-ssv ~ 0
ZO
40 GO ATOMIC ~ Rb
SO
I00
Fig. 1. Dependence of electrocatalytic activity, for methanol oxidation, on surface composition of Pt--Rh alloys. Steady-state currents in 0.1 M CH3OH + 1 M H2SO 4 at 25°C using the potential program shown.
80 consisted of an initial rapid fall in current followed by a continuous slow decay. The surface composition and real area of each alloy were monitored during experiments. After a single application of the potential program, the rhodium surface concentration was f o u n d to have decreased by ~ 4% in alloys containing ~40% Rh and by ~2% in the remainder. Real surface areas in the corresponding composition ranges were f o u n d to have increased by about 10 and 3% respectively. The averages of the values of composition and area before and after each measurement were taken as representing the surface of the electrode. As rhodium was dissolved during electrode cleaning, the cell electrolyte was replaced by a fresh sample after each activity measurement. Gradual build-up of rhodium i n s o l u t i o n must be avoided otherwise the current obtained will not originate solely from methanol oxidation but will contain a cathodic contribution from deposition of dissolved rhodium. In this investigation the rate of oxidation of methanol is assumed to have reached a state corresponding to equality between the rates of reactions (1) and (3) at the end of the initial rapid decay in the current transient, the current at 100 s (/100) being taken as the electrocatalytic activity. Specific activities (activity per unit real area) of platinum--rhodium alloys at various potentials in 0.1 M CH3OH + 1 M H2SO 4 are shown in Fig. 1 as a function of surface composition. It can be seen that, at low potentials, a m a x i m u m is reached at 34% Rh. As the experimental potential is increased, the m a x i m u m is shifted towards lower rhodium surface concentrations until at 0.68 V no synergistic effect is observed. The activities for pure platinum and rhodium electrodes are also given in Fig. 1 and were found to be in good agreement with values obtained using the more vigorous cleaning program advocated by Biegler [21]. The latter procedure involves holding the electrode at a higher potential of 1.88 V for 100 s, a treatment which is unsuitable for the alloys since the increased rate of rhodium dissolution at this potential would cause appreciable changes in surface composition. Furthermore, the corresponding increased concentration of rhodium in solution would lead to an increase in the contribution of metal deposition to the over-all current: e.g. the current at ¢ = 0.68 V from a pure rhodium electrode cleaned by Biegler's program became cathodic after 25 s. Electrocatalytic activities of mixed electro-deposits were measured in 1 M HCHO + 0.5 M H2SO 4 solution at 20°C. Application of the potential program shown in Fig. 1 to these electrodes would result in significant dissolution of the non-platinum metal. Activities were therefore obtained from slow linear potential sweeps (2 mV s- 1 ) between 0.0 and 0.6 V. Over this potential range the currents observed in the supporting electrolyte were negligibly small compared with those in the presence of the organic fuel. Hence, currents due to oxidation and dissolution of the ad-element can be neglected in activity measurements for formaldehyde oxidation in these electrode systems. The specific activities of the various electrodes at 0.2 V are given in Table 2 together with the Tafel slopes of the corresponding polarization curves.
81 TABLE 2 Activities of co-deposits for formaldehyde oxidation (1 M HCHO; 0.5 M H2SO4) Electrode
Current at 0.2 V/ pA cm - 2
Tafel slope/mV
Upper limit of Tafel region/V
Pt Pt--Sn Pt--Sb Pt--Re Pt--As Pt--Bi Pt--Te Pt--Hg
0.004 71 3.4 1.2 3.0 30 5.0 0.2
103 125 155 105 175 100 143 93
>0.6 0.2 >0.5 >0.5 0.3 0.2 0.35 0.3
DISCUSSION A b i n a r y m e t a l catalyst m a y consist of a single h o m o g e n e o u s phase or o i an i n t i m a t e m i x t u r e o f t w o distinct phases. In b o t h cases we will assume t h a t t h e m e c h a n i s m of organic e l e c t r o - o x i d a t i o n can be r e p r e s e n t e d by reactions (1)--(3). T h e surface a t o m s o f b o t h metals in a h o m o g e n e o u s alloy have b e e n s h o w n [30] t o behave in an identical m a n n e r for c h e m i s o r p t i o n reactions, and this will result in a c o n t i n u o u s range o f activities for t h e t h r e e reactions as the c o m p o s i t i o n is varied. On the o t h e r hand, f o r h e t e r o g e n e o u s s y s t e m s the ads o r p t i o n o f the organic c o m p o u n d m u s t be e x p e c t e d to o c c u r o n l y o n t h e p l a t i n u m phase. T h e a d s o r b e d i n t e r m e d i a t e c o u l d t h e n r e a c t with o x y g e n species associated with the s e c o n d phase. In this case, t h e r e a c t i o n will p r o c e e d at t h e b o u n d a r y b e t w e e n the t w o phases. T h e rate o f r e a c t i o n (1), on p l a t i n u m or an alloy surface, is given b y
il = ( F k T / h ) a h a M e x p ( ' A G ° ~ / R T )
(4)
where aA and a M are the activities o f t h e dissolved organic c o m p o u n d and the m e t a l surface respectively, AG °¢ is the standard e l e c t r o c h e m i c a l free e n e r g y o f activation o f r e a c t i o n (1), and t h e o t h e r t e r m s have t h e i r usual significance. AG °¢ can be separated into the t e r m s Ag °, the m e t a l - i n d e p e n d e n t free energy o f t h e activated c o m p l e x at ¢ = 0, AG °, t h e free energy o f a d s o r p t i o n o f A at ~b = 0, and a p o t e n t i a l t e r m , ~F¢. T h e r e f o r e ,
il = ( F k T / h )aA (1-- 0B) e x p ( - - A g ° / R T -- ~J A G ° / R T + fl F ¢ / R T )
(5)
where (1 - - 0 B ) , t h e f r a c t i o n o f t h e m e t a l surface u n o c c u p i e d b y a d s o r b e d species, r e p r e s e n t s the activity o f t h e m e t a l surface. The a d s o r p t i o n o f organic c o m p o u n d s such as _meth~ngl [5] on p l a t i n u m follows Elovich kinetics, f r o m which it can be d e d u c e d t h a t the free e n e r g y
82 of adsorption varies linearly with coverage [37], i.e. AG ° = A G ° 0 + fRTOB
(6)
where f is a constant. Neglecting the pre-exponential 0B term [38], i 1 = ( F k T / h ) a A e x p ( - - A g ° / R T -- fl A G ° o / R T + fi F O / R T -- fifOB)
(7)
The equilibrium coverage of OHad ~ (reaction 2) will be small for platinum in the potential range where the organic oxidation was studied, and the adsorption can be considered in terms of a Langmuir isotherm 0OH = e x p ( - - A G ° / R T + F ¢ / R T )
(8)
where AG ° is the free energy of adsorption of OHad s. The rate of reaction (3) is given by i 3 = (FkT/h)OBOoH e x p ( - - A ~ , ° / R T + fl A G ° , o / R T + fi A G ° / R T + fl F O / R T
(9)
+ J3fOB)
where A~ ° is the metal-independent free energy of the activated complex at ¢ = 0. Substituting (8) in (9) and neglecting the pre-exponential 0 B term [38] gives i 3 = ( F k T / h ) e x p ( - - A g ° / R T + fl A G ° o / R T -- (1 -- i3) A G ° / R T + (1 + fl) F ¢ / R T
+ ~fOB)
(! o)
At higher coverages of OHad~, which could be appropriate for the mixedmetal systems, the adsorption will be expected to follow a Frumkin isotherm. Neglecting pre-exponential 0 terms, exp 0 o H = exp(--AG°,0/R T + F ¢ / R T)
(11)
where A G ° 0 is the free energy of adsorption of OHads at zero coverage. In this case (neglecting pre-exponential 0 terms) i 3 = ( F h T / h ) exp(--A~ ° / R T + fl A G ° , o / R T + ~ A G ° o / R T + ~ F ¢ / R T + ~fOB
(12)
+ flfOoH )
Substituting (11) in (12) gives i 3 = ( F k T / h ) e x p ( - - A $ ° / R T + fi A G ° o / R T -- (1 -- f i ) A G ° o / R T + (1 + f i ) F ¢ / R T
(13)
+ ~fOn)
which is identical in form to (10). Hence, the rate equation is independent of whether the adsorption of OH obeys a Langmuir or a Frumkin isotherm. Since reactions (1) and (3) are irreversible, the rates must be equal in the Steady state. Equating (7) with (13) gives exp [JfOB
=
aA 1/2 e x p ( - - A $ ° / 2 R T + A $ ° / 2 R T -- [3 A G ° , o / R T
+ (1 -- fl)Aa°2,o/2RT-- F O / 2 R T )
(14)
83
Substitution in (13) gives i3 = ( F k T / h ) aA 1/2 e x p ( - - A ~ ° / 2 R T - - A ~ ° / 2 R T - - (1 - - ~) A G ° o / 2 R T
+ (1 + fi --½ ) F ¢ / R T )
(15)
If we include all the free energy terms independent of the metal in a constant term, then, assuming fi - ~z , the over-all rate will be given by i = KaA 1/2 e x p ( - - A G ° o / 4 R T
+ F¢/RT)
(16)
An additional parameter, Z, must be introduced for the reaction on heterogeneous surfaces to account for the reaction zone being the interface between the t w o phases. i = KZaA 1/2 e x p ( - - A G ° , o / 4 R T + Fdp/RT)
(17)
It can be seen that the rate of the over-all reaction will be dependent on the free energy of adsorption of oxygen. In the case of heterogeneous alloys, it is postulated that o x y g e n adsorbs on the second phase. It is to be expected that the free energy will f o l l o w the order of the standard potentials for the formation of the oxide on these elements. In Fig. 2 the current density for formaldehyde oxidation is presented as a function of the standard potentials [ 39]
Sn\,
o \\ \
\
Bt o
\ \
\
\
\
\
Te
$ob~\\ \
o
Re o
"0
\\
\ k
\ \ \ \
oH9
\
_,.i
\
\ \
\
\ \ \ \ \
-2
-O'Z
\ \
I
0
I
O'Z
I
0"4
I
O'G
E°/v (~s SHE)
I 0"8
\Pt bl I'0
Fig. 2. D e p e n d e n c e of the e l e c t r o c a t a l y t i c activity, for f o r m a l d e h y d e o x i d a t i o n o f binary c o - d e p o s i t e d electrodes c o n t a i n i n g p l a t i n u m , on the m e t a l / m e t a l - o x i d e standard p o t e n t i a l o f the s e c o n d c o m p o n e n t . A c t i v i t y is the current at 0 . 2 V in 1 M H C H O + 0 . 5 M H 2 S O 4 at 2 0 ° C. E ° is the p o t e n t i a l for the l o w e s t o x i d e species.
84 for the formation of the first oxide species. Also included in this Figure is a line with the slope expected from (17) assuming that the adsorption of OH species occurs at a potential close to the value for the formation of the first stable oxide species and that Z is constant. The reaction zone parameter Z could differ significantly between the different electrodes. Hence, an exact correlation cannot be expected. However, it can be seen that there is a definite trend towards greater electrocatalytic activities as the E ° value decreases. The activity displayed by heterogeneous platinum--gold alloys for methanol oxidation has been shown [40] to be simply that of the platinum phase exposed at the surface. This is to be expected from the mechanism proposed here since gold only adsorbs oxygen at potential above 1.35 V [35]. Gold, therefore, cannot supply an oxygen species to oxidize methanol adsorbed on the platinum phase. Equation (16) is applicable to homogeneous alloys, and this equation predicts that the rate of oxidation will be independent of any changes in the free energy of adsorption of the organic species with change in composition of the alloy surface. An increase in activity over that of platinum will be evident, however, if the free energy of adsorption of OH, AG°0, becomes more negative with addition of the alloying metal. Rhodium adsorbs OH at a lower potential than platinum; the potentials corresponding to hydrogen and oxygen adsorption and desorption on platinum--rhodium alloys change in a linear manner with alloy composition [ 30]. Therefore, alloying rhodium with platinum will result in a more negative AG°,0 and this should lead to an enhanced activity. Figure 1 shows that this system does exhibit a synergistic effect. The decrease in activity at higher rhodium concentrations can be explained by a corresponding decrease in coverage of the organic intermediate to values below that at which Elovich kinetics apply. In this case we can assume Langmuir adsorption for the organic compound, and the over-all rate will have the same form as eqn. (5). Hence, i = g a A exp(--~ A G ° / R T + ~ F ¢ / R T )
(18)
where, again, all terms independent of the metal are included in the constant. Methanol is adsorbed to a much smaller extent on rhodium than on platirmm in the potential range investigated here [15], and hence AD o will become less negative with increase in rhodium content of the surface and the reaction rate will decrease. The mechanism therefore, accounts for the m a x i m u m in the activity--composition curve even though the free energies of adsorption of the intermediate species are assumed to change in a m o n o t o n i c manner with surface composition. The change in properties on going from Temkin to Langmuir conditions [38] can be considered as a change in the rate determining step from (3) to (1). The m a x i m u m on the curves in Fig. 1 shifts to lower r h o d i u m concentrations as the potential at which the activity is determined is increased. There is no m a x i m u m on the curve corresponding to 0.68 V. This behaviour is due to inhibition of the electrochemical reaction at the higher potentials. Inhibi-
85 tion, which occurs for a range o f o x i d a t i o n reactions including h y d r o g e n ionization [17], is a result o f a d e a c t i v a t i o n o f the surface w h e n o x y g e n is ads o r b e d t o a n y significant e x t e n t . T h e p o t e n t i a l o f c o m m e n c e m e n t o f o x y g e n a d s o r p t i o n decreases with increase in r h o d i u m c o n t e n t , and this results in a c o m p a r i s o n o f activities at 0.68 V being b e t w e e n c u r r e n t s f r o m an active plati n u m surface and alloys which have an increasing degree o f d e a c t i v a t i o n f r o m oxygen adsorption. T h e analysis o f t h e m e c h a n i s m o f organic o x i d a t i o n b y binary p l a t i n u m catalysts suggests t h a t e n h a n c e d activity over p l a t i n u m alone can be e x p e c t e d for b o t h h e t e r o g e n e o u s and h o m o g e n e o u s systems. In t h e f o r m e r case t h e activity c o u l d be limited by t h e s u p p l y o f o x y g e n t o the organic species ads o r b e d o n the p l a t i n u m phase, since this m u s t t a k e place across a phase bound. ary, b u t t h e t h e o r e t i c a l limit will be the rate o f a d s o r p t i o n of t h e organic species o n p l a t i n u m at zero coverage. The activity o f h o m o g e n e o u s alloys c o u l d n o t reach this limit because t h e rate of a d s o r p t i o n of t h e organic species will be e x p e c t e d t o decrease with the a d d i t i o n of t h e alloying metal. This will result in a l o w e r rate at l o w coverages a n d suggests t h a t a h e t e r o g e n e o u s s y s t e m w o u l d p r o v i d e t h e o p t i m u m electrocatalyst. A similar e n h a n c e m e n t o f activity of binary catalysts is t o be e x p e c t e d for the o x i d a t i o n o f organic species, like f o r m i c acid, which p r o c e e d t h r o u g h a parallel path, i n h i b i t e d b y the a d s o r b e d i n t e r m e d i a t e , Bad s. T h e increase in the rate o f r e m o v a l o f Bad~, due t o i n t e r a c t i o n with t h e m o s t active o x y g e n species associated with the a d d e d e l e m e n t , will decrease t h e coverage of t h e b l o c k i n g i n t e r m e d i a t e a n d h e n c e lead t o an e n h a n c e m e n t in the rate o f t h e over-all o x i d a t i o n process. ACKNOWLEDGEMENTS T h e a u t h o r s wish to t h a n k Mr. H. B e c k e r f o r his assistance in the experim e n t a l w o r k and Dr. K.J. C a t h r o f o r analyses o f t h e m i x e d m e t a l co-deposits. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13
S. Gilman, Trans. Faraday Soc., 61 (1965) 2546, 2561. H. Wroblowa, B.J. Piersma and J. O'M. Bockris, J. Electroanal. Chem., 6 (1963) 401. S. Gilman, J. Phys. Chem., 66 (1962) 2657. V.S. Bagotzkii and Yu.B. Vasil'ev, Electrochim. Acta, 12 (1967) 1323. T. Biegler and D.F.A. Koch, J. Electrochem. Soc., 114 (1967) 904. D.F.A. Koch in J.A. Friend and F. Gutmann (Eds.), Proc. First Australian Conference on Electrochemistry, 1963, Pergamon Press, Oxford, 1965, p. 657. S.B. Brummer and A.C. Makrides, J. Phys. Chem., 68 (1964) 1448. B.I. Podlovchenko and E.P. Gorgonova, Dokl. Akad. Nauk SSSR, 156 (1964) 673. V.S. Bagotzkii and Yu.B. Vasil'ev, Electrochim. Acta, 11 (1966) 1439. M.W. Breiter, J. Electroanal. Chem., 14 (1967) 407. T. Biegler, J. Phys. Chem., 72 (1968) 1571. A. Capon and R. Parsons, J. Electroanal. Chem., 44 (1973) 1, 239; 45 (1973) 205. O.A. Petrii, B:I. Podlovchenko, A.N. Frumkin and Hira Lal, J. Electroanal. Chem., 1(} (1965) 253.
86 14 B.I. Podlovehenko, A.N. Frumkin and V.S. Stenin, Elektrokhimiya, 4 (1968) 339. 15 V.S. Bagotzkii, Yu.B. Vasil'ev, O.A. Khazova and S.S. Sedova, Electrochim. Acta, 16 (1971) 913. 16 S.S. Beskorovainaya, Yu.B. Vasil'ev and V.S. Bagotzkii, Elektrokhimiya, 8 (1972) 376. 17 V.F. Steuin, V.E. Kazarinov and B.I. Podlovchenko, Elektrokhimiya, 5 (1969) 442. 18 M.W. Breiter, Electrochemical Processes in Fuel Cells, Springer-Verlag, New York, 1969, Ch. X, p. 166. 19 M.A. Khrizolitova, L.A. Mirkind, Yu.B. Vasil'ev, M.Ya. Fioshin and V.S. Bagotzkii, Elektrokhimiya, 8 (1972) 548. 20 B.I. Podlovchenko and R.P. Petukhova, Elektrokhimiya, 9 (1973) 273. 21 T. Biegler, Aust. J. Chem., 26 (1973) 2571, 2587. 22 D.F.A. Koch, Australian Patent No. 46,123, 1964. 23 K.J. Cathro, Electrochem. Technol., 5 (1967) 441. 24 K.J. Cathro, J. Electrochem. Soc., 116 (1969) 1608. 25 O.J. Adlhart and A.J. Hartner, Proc. 20th Power Sources Conference, PSC Publications, New Jersey, 1966, p. 11. 26 M. Watanabe and S. Motoo, J. Electroanal. Chem., 60 (1975) 259, 267, 275. 27 M. Watanabe, T. Suzuki and S. Motoo, Denki Kagaku, 40 (1972) 205, 210. 28 C.E. Heath, Proc. Int. Congr. Fuel Cells, Brussels, 1965, S.E.R.A.I.-COMASCI, 1965, p. 99. 29 J.A. Shropshire, J. Electrochem. Soc., 112 (1965) 465. 30 D.A.J. Rand and R. Woods, J. Electroanal. Chem., 36 (1972) 57. 31 D.A.J. Rand and R. Woods, J. Electroanal. Chem., 44 (1973) 83. 32 D.A.J, Rand and R. Woods in M.W. Breiter (Ed.), Proc. Symposium on Electrocatalysis, 1974, The Electrochemical Society, New Jersey, p. 140. 33 K.J. Cathro, D.F.A. Koch and H.R. Skewes, Mech. Chem. Eng. Trans., MC3 (1967) 17. 34 K.J. Cathro and C.H. Weeks, Energy Convers., 11 (1971) 133. 35 D.A.J, Rand and R. Woods, J. Electroanal. Chem., 31 (1971) 29. 36 T. Biegler, D.A.J. Rand and R. Woods, J. Electroanal. Chem., 29 (1971) 269. 37 E. Gileadi and B.E. Conway in J.O'M. Bockris and B.E. Conway (Eds.), Modern Aspects of Electrochemistry, No. 3, Butterworths, London, 1964, Ch. 5. 38 T. Biegler and R. Woods, J. Phys. Chem., 73 (1969) 3502. 39 M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solution, Pergamon Press, Oxford, 1966. 40 M.W. Breiter, J. Phys. Chem., 69 (1965) 3377.
73
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
BINARY ELECTROCATALYSTS FOR ORGANIC OXIDATIONS
D.F.A. KOCH, D.A.J. RAND and R. WOODS CSIRO Division o f Mineral Chemistry, P.O. Box 124, Port Melbourne, Victoria 3207 (Australia) *
(Received 30th September 1975)
ABSTRACT The electro-oxidation of formaldehyde and methanol has been studied on a number of binary platinum electrocatalysts. These comprised mixed electro-deposits of Pt with Sb, As, Bi, Hg, Re, Te or Sn and a range of homogeneous Pt--Rh alloys of different, known, surface composition. These systems were found to exhibit an enhanced activity over that of platinum alone, and this behaviour was correlated with the eas e of adsorption of oxygen on the added metal. The activities for organic oxidation were compared with predictions of a model involving reaction between adsorbed oxygen and organic species on the metal surface. The proposed mechanism accounts for the behaviour of both homogeneous and heterogeneous alloy systems.
INTRODUCTION H y d r o c a r b o n s , c a r b o n m o n o x i d e a n d water-soluble organic c o m p o u n d s such as m e t h a n o l , f o r m a l d e h y d e a n d f o r m i c acid have b e e n c o n s i d e r e d as rea c t a n t s in fuel cell systems. P l a t i n u m has b e e n f o u n d t o be t h e m o s t active single m e t a l f o r t h e c o m p l e t e e l e c t r o - o x i d a t i o n o f these c o m p o u n d s t o c a r b o n d i o x i d e in b o t h acid a n d alkaline solutions. It is well r e c o g n i z e d t h a t these fuels (e.g. e t h a n e [1], e t h y l e n e [2], c a r b o n m o n o x i d e [3], m e t h a n o l [4,5], f o r m a l d e h y d e [6] a n d f o r m i c acid [7] ) f o r m s t r o n g l y a d s o r b e d i n t e r m e d i a t e s as given b y A -~ Bads + n l H + + n l e
(1)
T h e i n t e r m e d i a t e can t h e n react with an o x y g e n species derived f r o m water, H 2 0 ~ OHad s + H + + e
(2)
Bads + OHads -~ CO2 + n2H + + n2e
(3)
The a d s o r p t i o n r e a c t i o n (1) involves d e h y d r o g e n a t i o n o f t h e o r g a n i c species; B~d~ f r o m m e t h a n o l h a v i n g been identified as - - C O H [4,8,9] or - - C O [5, * All correspondence concerning this paper to be addressed to The Chief, CSIRO Division of Mineral Chemistry, at address given.
74 10,11]. Strong evidence has been presented [10--12] to support the conclusion that the intermediate has the same constitution in the oxidation of methanol, formaldehyde, formic acid and carbon monoxide. Reaction (1) is irreversible; an electrode held at low potentials can be removed from a cell containing a solution of the organic species, washed and returned to a cell containing acid alone without loss of adsorbate from the electrode surface [11,13]. In the potential region where (3) becomes significant, the surface coverage of Bad s is determined by a balance between its formation by (1) and its further oxidation to CO2 [5]. The existence of a parallel path, more rapid than (1)--(3), has been established for the oxidation of formic acid [12]. The additional route proceeds through a weakly bonded intermediate and is inhibited by the formation of the strongly bonded species. There is continuing controversy regarding whether such a mechanism also applies to the oxidation of methanol in acid solution. The formation of formaldehyde and formic acid in solution during methanol oxidation has been considered [14] to be evidence in support of a parallel path mechanism. However, reaction (1) is n o t a single step since its kinetics [4,5] obey the rate equation for a single electron transfer process, whereas nl is 3 or 4 depending on whether Bad s is --COH or --CO. Hence, the side products could be derived from precursor adsorbates produced by consecutive removal of hydrogen atoms from methanol [15]. Supporting evidence in favour of this suggestion is obtained from electrochemical [16] and radiochemical investigations [17] which have failed to detect any intermediates other than those which result in the formation of the strongly bonded species. Experimental evidence [13,18] which suggests that the steady state rate of methanol oxidation is more than four times faster than the rate of oxidation of the adsorbed layer, determined after removal of the alcohol from the cell solution, lends support to a parallel path mechanism. Other workers [14,16], however, have found the t w o rates to be identical as must be the case if the predominant oxidation route is (1)--(3). Capon and Parsons [12] compared the oxidation of methanol with that of formic acid on platinum and palladium electrodes. In contrast to platinum, palladium does not form a strongly b o n d e d intermediate" from formic acid, and hence in this case oxidation does not become inhibited. The complete inactivity of palladium for methanol oxidation was considered to be strong evidence that the oxidation of this c o m p o u n d requires the participation of the strongly adsorbed intermediate, i.e. it can proceed only by (1)--(3). Investigations on the influence of the presence of other organic substances on methanol oxidation [19] also support a single reaction pathway. The mechanism of methanol oxidation on rhodium electrodes has also been shown [15] to proceed b y the mechanism (1)--(3). It has been suggested by Podlovchenko and Petukhova [20] that (1)--(3) constitutes the mechanism on activated smooth platinum under conditions which might be called "quasistatic", but that a parallel path operates in the final "static state". Biegler [ 21] has shown that the instantaneous methanol oxi-
75 dation current on an activated platinum electrode decreased rapidly as the rates of reactions (1) and (3) approach each other. This rapid decrease was followed by a slower rate of decay due to the accumulation of impurities on the electrode surface. He concluded that the current after the initial rapid decrease (i.e. the equivalent of a "quasistatic state") gave the best measure of the activity of the surface. In our experiments, this current was taken to represent the activity of Pt--Rh alloys for methanol oxidation and can be related to mechanism (1)--(3). For formaldehyde oxidation, a parallel path involving the formation of a blocking polymeric surface species has been shown [6] to occur at potentials above 0.4 V. In the work reported here, activities were compared at potentials significantly below this value, and hence activities for this system can also be related to mechanism (1)--(3). The search for improved electrocatalysts for fuel cells has shown that enhanced activity over pure platinum for organic oxidations can be obtained by the incorporation of other metals such as rhenium [22,23], tin [22,24], ruthenium [13,25,26] or osmium [27]. The improvement in performance in the presence of these metals has been attributed [2,22,26--28] to oxidation of the intermediate by oxygen-containing species which are associated with the added metal and are more reactive than O H a d s o n platinum. I t h a s been observed [29] that the adsorption of molybdate ions increases the activity of platinum, the proposed mechanism for this system involving oxidation of the fuel with adsorbed molybdate, followed by electro-oxidation of the reduced molybdate species. Recently, Watanabe and Motoo [26] proposed a mechanism to explain the enhancement of methanol and carbon monoxide oxidation by ruthenium adatoms on platinum. These authors consider the reaction to occur between the carbonaceous species adsorbed on platinum sites and oxygen adsorbed on ruthenium sites in a homogeneous alloy surface. This postulate assumes the platinum and ruthenium sites in this surface to have different adsorption properties. Although their explanation would apply to heterogeneous systems, it is not consistent with investigations of the chemisorption properties of homogeneous noble metal alloys [30--32]. The characteristics of adsorption of hydrogen and oxygen on such systems were shown to be a composite of those of the individual metals, demonstrating that the surface atoms of both metals behave in an identical manner. Hence, all sites on a homogeneous alloy surface will be expected to have the same activity for the adsorption of both oxygen and organic species. In this communication, we present the results of investigations of the activity of various binary platinum catalysts for the oxidation of methanol and formaldehyde. Two types of experiment were performed. Smooth Pt--Rh alloy electrodes gave results which were consistent with a proposed reaction model. The study was extended to the behaviour of more realistic electrodes for fuel cells consisting of co-deposits with platinum black. The Pt--Rh alloy system was chosen for investigation because the surface
76 composition can be determined in situ [30] and can be varied in a controlled manner between the pure metal limits by an electrochemical pretreatment without the occurrence of phase separation at the surface. Methanol was used here because of the extensive information already available on the oxidation mechanism: Electrodes with high surface area were prepared by co-depositing platinum with either antimony, arsenic, bismuth, mercury, rhenium, tellurium or tin. All these elements are stable in acid solutions either as the metal or as an oxide over the range of potentials where activity was measured. There is evidence that certain of these systems, e.g. Pt--Re [23] and Pt--Sn [24], are heterogeneous consisting of two solid phases. Formaldehyde was used as the organic reactant in these studies because of its relevance to a formaldehyde/air fuel cell developed in these laboratories [33,34]. The experimental results for both types of electrode are compared with the predictions of a model for both homogeneous and heterogeneous alloy surfaces in which the second component metal results in a more active adsorbed oxygen species. EXPERIMENTAL AND RESULTS Apparatus
Experiments were carried out in three-compartment glass cells, the working compartment of each cell having a side arm and an outlet tube to allow the addition and removal of solutions. Cell electrolytes consisted of B.D.H. Aristar methanol or B.D.H. AnalaR formaldehyde in a supporting electrolyte, 0.5 or 1 M H2SO4, prepared from B.D.H. A n a l a g reagent and doubly distilled water. Solutions were deoxygenated with purified [35] nitrogen. A mercury/mercurous sulphate, 1 M H2SO 4 reference electrode was used, the potential of which was 0.68 V vs. the reversible hydrogen electrode (RHE) in 1 M H2SO 4. All potentials in this paper refer to the RHE. Measurements were carried out at 25°C unless otherwise reported. Electrode potentials were controlled with a potentiostat programmed by means of potential step and linear sweep generators constructed in these laboratories. Current--potential curves were recorded on a Hewlett-Packard 7004A XY recorder. A time base Type 17172A (Hewlett-l~ackard) was used in conjunction with the XY recorder for displaying current--time traces. Ele c tro de p repa ration
Platinum, rhodium and platinum--rhodium alloy electrodes consisted of 26 gauge wire of 99.99% purity sealed into soft glass tubing. The real surface areas of platinum and rhodium electrodes were determined by measuring the hydrogen adsorbed from 1 M H2SO 4 during a cathodic potential sweep at 40 mV s-1. The electrode surfaces were first cleaned by the application of
77 triangular potential cycles between 0.07 and 1.54 V until a voltammogram of reproducible shape was obtained (~ 20 cycles). The procedure used for integrating the curves has been reported previously [35,36]. Real electrode areas were calculated on the basis that the fractional hydrogen coverage at the lower integration limit was 0.77 and 0.59 and the corresponding monolayer charge 210 and 221 pC for platinum [36] and rhodium [35] respectively. For computing the area of alloy electrodes, it was assumed that these values changed linearly with alloy composition. The preferential dissolution of rhodium which is found [30] to occur during potential cycling of platinum--rhodium alloys provides an excellent method for preparing alloys with different surface compositions. In the work reported here, platinum--rhodium alloys were prepared by subjecting an alloy containing 74% rhodium * to repetitive potential cycling at 40 mV s-1 between 0.07 and 1.54 V in 1 M H2SO 4 until the surface composition reached the required value. This was determined using the linear relationship which has been established [30] between surface composition and the potential of the oxygen desorption peak on a voltammogram. After obtaining an alloy of the desired composition, the anodic limit of the following potentia} sweep was lowered to 0.80 V and the real area of the electrode determined from the hydrogen adsorption charge. The lower anodic limit was necessary for the areadetermining sweep, as the oxygen reduction and hydrogen adsorption regions overlap on sweeps from potentials > 1 . 0 V. High surface area electrolytic co-deposits of platinum with another element were prepared by plating for 30 min at a controlled potential. This was --0.15 V except in the case of p l a t i n u m - - a n t i m o n y co-deposits where a deposition potential of 0.0 V was used. The electrolyte usually consisted of 1% chloroplatinic acid in 1 M HC1 solution to which was added the other element in the form shown in Table 1. By varying the pH of the solution and the a m o u n t of the second element present, the deposit composition was controlled at approximately 4% of the added metal. The deposits were formed as finely-divided "blacks". Platinum, tantalum and iridium foils of 1 cm 2 geometric area were tested as substrates, the latter being f o u n d to be the best choice as it was not attacked during dissolution of the deposit in aqua regia for analysis. Mercury, platinum, rhenium and tellurium analyses were made using spectrophotometric methods while antimony, arsenic, bismuth and tin were determined polarographically. After preparation of the deposit, the electrode was washed thoroughly with distilled water and transferred to a second cell in which the real surface area was determined using the hydrogen adsorption m e t h o d described above for platinum, the electrode being cycled at 20 mV s-1 between 0 and 0.6 V in 0.5 M H2SO 4. The cathodic charge in the hydrogen region for platinum-arsenic and platinum--tin deposits was found to be less than the anodic, in-
* Alloy compositions are expressed in this paper as atomic percentages.
78
I d 4 d 4 4 4 4
II
o o
o o
o o 0
o
0 °~
0
IIIII
79
dicating a contribution from an anodic process other than hydrogen oxidation. For these deposits, areas were obtained from the mean value of the anodic and cathodic charges. After determination of the real surface area of each electrode, the contents of the cell were blown out with nitrogen, the cell washed with supporting electrolyte and the methanol or formaldehyde test solution added. Activity measurements The potential program used for activity measurements on platinum, rhodium and platinum--rhodium alloy electrodes in methanol solutions is shown in Fig. 1. Any adsorbed impurities or methanol residues are removed by oxidative desorption during the first step at 1.5 V. At this potential the electrode surface is covered with chemisorbed oxygen and any oxygen evolved is removed by bubbling nitrogen through the solution. The surface layer of adsorbed oxygen is reduced rapidly during the short step at 0.10 V, after which the final potential, ¢, is applied and the current--time curve recorded. The current transient was similar to t h a t observed previously by Biegler [ 21] and
ZOO
I'5V
~ : 0"65V ~,
150
Is
E
~too I--
=O'G3V 50
oLL___c - ~ =o-ssv ~ 0
ZO
40 GO ATOMIC ~ Rb
SO
I00
Fig. 1. Dependence of electrocatalytic activity, for methanol oxidation, on surface composition of Pt--Rh alloys. Steady-state currents in 0.1 M CH3OH + 1 M H2SO 4 at 25°C using the potential program shown.
80 consisted of an initial rapid fall in current followed by a continuous slow decay. The surface composition and real area of each alloy were monitored during experiments. After a single application of the potential program, the rhodium surface concentration was f o u n d to have decreased by ~ 4% in alloys containing ~40% Rh and by ~2% in the remainder. Real surface areas in the corresponding composition ranges were f o u n d to have increased by about 10 and 3% respectively. The averages of the values of composition and area before and after each measurement were taken as representing the surface of the electrode. As rhodium was dissolved during electrode cleaning, the cell electrolyte was replaced by a fresh sample after each activity measurement. Gradual build-up of rhodium i n s o l u t i o n must be avoided otherwise the current obtained will not originate solely from methanol oxidation but will contain a cathodic contribution from deposition of dissolved rhodium. In this investigation the rate of oxidation of methanol is assumed to have reached a state corresponding to equality between the rates of reactions (1) and (3) at the end of the initial rapid decay in the current transient, the current at 100 s (/100) being taken as the electrocatalytic activity. Specific activities (activity per unit real area) of platinum--rhodium alloys at various potentials in 0.1 M CH3OH + 1 M H2SO 4 are shown in Fig. 1 as a function of surface composition. It can be seen that, at low potentials, a m a x i m u m is reached at 34% Rh. As the experimental potential is increased, the m a x i m u m is shifted towards lower rhodium surface concentrations until at 0.68 V no synergistic effect is observed. The activities for pure platinum and rhodium electrodes are also given in Fig. 1 and were found to be in good agreement with values obtained using the more vigorous cleaning program advocated by Biegler [21]. The latter procedure involves holding the electrode at a higher potential of 1.88 V for 100 s, a treatment which is unsuitable for the alloys since the increased rate of rhodium dissolution at this potential would cause appreciable changes in surface composition. Furthermore, the corresponding increased concentration of rhodium in solution would lead to an increase in the contribution of metal deposition to the over-all current: e.g. the current at ¢ = 0.68 V from a pure rhodium electrode cleaned by Biegler's program became cathodic after 25 s. Electrocatalytic activities of mixed electro-deposits were measured in 1 M HCHO + 0.5 M H2SO 4 solution at 20°C. Application of the potential program shown in Fig. 1 to these electrodes would result in significant dissolution of the non-platinum metal. Activities were therefore obtained from slow linear potential sweeps (2 mV s- 1 ) between 0.0 and 0.6 V. Over this potential range the currents observed in the supporting electrolyte were negligibly small compared with those in the presence of the organic fuel. Hence, currents due to oxidation and dissolution of the ad-element can be neglected in activity measurements for formaldehyde oxidation in these electrode systems. The specific activities of the various electrodes at 0.2 V are given in Table 2 together with the Tafel slopes of the corresponding polarization curves.
81 TABLE 2 Activities of co-deposits for formaldehyde oxidation (1 M HCHO; 0.5 M H2SO4) Electrode
Current at 0.2 V/ pA cm - 2
Tafel slope/mV
Upper limit of Tafel region/V
Pt Pt--Sn Pt--Sb Pt--Re Pt--As Pt--Bi Pt--Te Pt--Hg
0.004 71 3.4 1.2 3.0 30 5.0 0.2
103 125 155 105 175 100 143 93
>0.6 0.2 >0.5 >0.5 0.3 0.2 0.35 0.3
DISCUSSION A b i n a r y m e t a l catalyst m a y consist of a single h o m o g e n e o u s phase or o i an i n t i m a t e m i x t u r e o f t w o distinct phases. In b o t h cases we will assume t h a t t h e m e c h a n i s m of organic e l e c t r o - o x i d a t i o n can be r e p r e s e n t e d by reactions (1)--(3). T h e surface a t o m s o f b o t h metals in a h o m o g e n e o u s alloy have b e e n s h o w n [30] t o behave in an identical m a n n e r for c h e m i s o r p t i o n reactions, and this will result in a c o n t i n u o u s range o f activities for t h e t h r e e reactions as the c o m p o s i t i o n is varied. On the o t h e r hand, f o r h e t e r o g e n e o u s s y s t e m s the ads o r p t i o n o f the organic c o m p o u n d m u s t be e x p e c t e d to o c c u r o n l y o n t h e p l a t i n u m phase. T h e a d s o r b e d i n t e r m e d i a t e c o u l d t h e n r e a c t with o x y g e n species associated with the s e c o n d phase. In this case, t h e r e a c t i o n will p r o c e e d at t h e b o u n d a r y b e t w e e n the t w o phases. T h e rate o f r e a c t i o n (1), on p l a t i n u m or an alloy surface, is given b y
il = ( F k T / h ) a h a M e x p ( ' A G ° ~ / R T )
(4)
where aA and a M are the activities o f t h e dissolved organic c o m p o u n d and the m e t a l surface respectively, AG °¢ is the standard e l e c t r o c h e m i c a l free e n e r g y o f activation o f r e a c t i o n (1), and t h e o t h e r t e r m s have t h e i r usual significance. AG °¢ can be separated into the t e r m s Ag °, the m e t a l - i n d e p e n d e n t free energy o f t h e activated c o m p l e x at ¢ = 0, AG °, t h e free energy o f a d s o r p t i o n o f A at ~b = 0, and a p o t e n t i a l t e r m , ~F¢. T h e r e f o r e ,
il = ( F k T / h )aA (1-- 0B) e x p ( - - A g ° / R T -- ~J A G ° / R T + fl F ¢ / R T )
(5)
where (1 - - 0 B ) , t h e f r a c t i o n o f t h e m e t a l surface u n o c c u p i e d b y a d s o r b e d species, r e p r e s e n t s the activity o f t h e m e t a l surface. The a d s o r p t i o n o f organic c o m p o u n d s such as _meth~ngl [5] on p l a t i n u m follows Elovich kinetics, f r o m which it can be d e d u c e d t h a t the free e n e r g y
82 of adsorption varies linearly with coverage [37], i.e. AG ° = A G ° 0 + fRTOB
(6)
where f is a constant. Neglecting the pre-exponential 0B term [38], i 1 = ( F k T / h ) a A e x p ( - - A g ° / R T -- fl A G ° o / R T + fi F O / R T -- fifOB)
(7)
The equilibrium coverage of OHad ~ (reaction 2) will be small for platinum in the potential range where the organic oxidation was studied, and the adsorption can be considered in terms of a Langmuir isotherm 0OH = e x p ( - - A G ° / R T + F ¢ / R T )
(8)
where AG ° is the free energy of adsorption of OHad s. The rate of reaction (3) is given by i 3 = (FkT/h)OBOoH e x p ( - - A ~ , ° / R T + fl A G ° , o / R T + fi A G ° / R T + fl F O / R T
(9)
+ J3fOB)
where A~ ° is the metal-independent free energy of the activated complex at ¢ = 0. Substituting (8) in (9) and neglecting the pre-exponential 0 B term [38] gives i 3 = ( F k T / h ) e x p ( - - A g ° / R T + fl A G ° o / R T -- (1 -- i3) A G ° / R T + (1 + fl) F ¢ / R T
+ ~fOB)
(! o)
At higher coverages of OHad~, which could be appropriate for the mixedmetal systems, the adsorption will be expected to follow a Frumkin isotherm. Neglecting pre-exponential 0 terms, exp 0 o H = exp(--AG°,0/R T + F ¢ / R T)
(11)
where A G ° 0 is the free energy of adsorption of OHads at zero coverage. In this case (neglecting pre-exponential 0 terms) i 3 = ( F h T / h ) exp(--A~ ° / R T + fl A G ° , o / R T + ~ A G ° o / R T + ~ F ¢ / R T + ~fOB
(12)
+ flfOoH )
Substituting (11) in (12) gives i 3 = ( F k T / h ) e x p ( - - A $ ° / R T + fi A G ° o / R T -- (1 -- f i ) A G ° o / R T + (1 + f i ) F ¢ / R T
(13)
+ ~fOn)
which is identical in form to (10). Hence, the rate equation is independent of whether the adsorption of OH obeys a Langmuir or a Frumkin isotherm. Since reactions (1) and (3) are irreversible, the rates must be equal in the Steady state. Equating (7) with (13) gives exp [JfOB
=
aA 1/2 e x p ( - - A $ ° / 2 R T + A $ ° / 2 R T -- [3 A G ° , o / R T
+ (1 -- fl)Aa°2,o/2RT-- F O / 2 R T )
(14)
83
Substitution in (13) gives i3 = ( F k T / h ) aA 1/2 e x p ( - - A ~ ° / 2 R T - - A ~ ° / 2 R T - - (1 - - ~) A G ° o / 2 R T
+ (1 + fi --½ ) F ¢ / R T )
(15)
If we include all the free energy terms independent of the metal in a constant term, then, assuming fi - ~z , the over-all rate will be given by i = KaA 1/2 e x p ( - - A G ° o / 4 R T
+ F¢/RT)
(16)
An additional parameter, Z, must be introduced for the reaction on heterogeneous surfaces to account for the reaction zone being the interface between the t w o phases. i = KZaA 1/2 e x p ( - - A G ° , o / 4 R T + Fdp/RT)
(17)
It can be seen that the rate of the over-all reaction will be dependent on the free energy of adsorption of oxygen. In the case of heterogeneous alloys, it is postulated that o x y g e n adsorbs on the second phase. It is to be expected that the free energy will f o l l o w the order of the standard potentials for the formation of the oxide on these elements. In Fig. 2 the current density for formaldehyde oxidation is presented as a function of the standard potentials [ 39]
Sn\,
o \\ \
\
Bt o
\ \
\
\
\
\
Te
$ob~\\ \
o
Re o
"0
\\
\ k
\ \ \ \
oH9
\
_,.i
\
\ \
\
\ \ \ \ \
-2
-O'Z
\ \
I
0
I
O'Z
I
0"4
I
O'G
E°/v (~s SHE)
I 0"8
\Pt bl I'0
Fig. 2. D e p e n d e n c e of the e l e c t r o c a t a l y t i c activity, for f o r m a l d e h y d e o x i d a t i o n o f binary c o - d e p o s i t e d electrodes c o n t a i n i n g p l a t i n u m , on the m e t a l / m e t a l - o x i d e standard p o t e n t i a l o f the s e c o n d c o m p o n e n t . A c t i v i t y is the current at 0 . 2 V in 1 M H C H O + 0 . 5 M H 2 S O 4 at 2 0 ° C. E ° is the p o t e n t i a l for the l o w e s t o x i d e species.
84 for the formation of the first oxide species. Also included in this Figure is a line with the slope expected from (17) assuming that the adsorption of OH species occurs at a potential close to the value for the formation of the first stable oxide species and that Z is constant. The reaction zone parameter Z could differ significantly between the different electrodes. Hence, an exact correlation cannot be expected. However, it can be seen that there is a definite trend towards greater electrocatalytic activities as the E ° value decreases. The activity displayed by heterogeneous platinum--gold alloys for methanol oxidation has been shown [40] to be simply that of the platinum phase exposed at the surface. This is to be expected from the mechanism proposed here since gold only adsorbs oxygen at potential above 1.35 V [35]. Gold, therefore, cannot supply an oxygen species to oxidize methanol adsorbed on the platinum phase. Equation (16) is applicable to homogeneous alloys, and this equation predicts that the rate of oxidation will be independent of any changes in the free energy of adsorption of the organic species with change in composition of the alloy surface. An increase in activity over that of platinum will be evident, however, if the free energy of adsorption of OH, AG°0, becomes more negative with addition of the alloying metal. Rhodium adsorbs OH at a lower potential than platinum; the potentials corresponding to hydrogen and oxygen adsorption and desorption on platinum--rhodium alloys change in a linear manner with alloy composition [ 30]. Therefore, alloying rhodium with platinum will result in a more negative AG°,0 and this should lead to an enhanced activity. Figure 1 shows that this system does exhibit a synergistic effect. The decrease in activity at higher rhodium concentrations can be explained by a corresponding decrease in coverage of the organic intermediate to values below that at which Elovich kinetics apply. In this case we can assume Langmuir adsorption for the organic compound, and the over-all rate will have the same form as eqn. (5). Hence, i = g a A exp(--~ A G ° / R T + ~ F ¢ / R T )
(18)
where, again, all terms independent of the metal are included in the constant. Methanol is adsorbed to a much smaller extent on rhodium than on platirmm in the potential range investigated here [15], and hence AD o will become less negative with increase in rhodium content of the surface and the reaction rate will decrease. The mechanism therefore, accounts for the m a x i m u m in the activity--composition curve even though the free energies of adsorption of the intermediate species are assumed to change in a m o n o t o n i c manner with surface composition. The change in properties on going from Temkin to Langmuir conditions [38] can be considered as a change in the rate determining step from (3) to (1). The m a x i m u m on the curves in Fig. 1 shifts to lower r h o d i u m concentrations as the potential at which the activity is determined is increased. There is no m a x i m u m on the curve corresponding to 0.68 V. This behaviour is due to inhibition of the electrochemical reaction at the higher potentials. Inhibi-
85 tion, which occurs for a range o f o x i d a t i o n reactions including h y d r o g e n ionization [17], is a result o f a d e a c t i v a t i o n o f the surface w h e n o x y g e n is ads o r b e d t o a n y significant e x t e n t . T h e p o t e n t i a l o f c o m m e n c e m e n t o f o x y g e n a d s o r p t i o n decreases with increase in r h o d i u m c o n t e n t , and this results in a c o m p a r i s o n o f activities at 0.68 V being b e t w e e n c u r r e n t s f r o m an active plati n u m surface and alloys which have an increasing degree o f d e a c t i v a t i o n f r o m oxygen adsorption. T h e analysis o f t h e m e c h a n i s m o f organic o x i d a t i o n b y binary p l a t i n u m catalysts suggests t h a t e n h a n c e d activity over p l a t i n u m alone can be e x p e c t e d for b o t h h e t e r o g e n e o u s and h o m o g e n e o u s systems. In t h e f o r m e r case t h e activity c o u l d be limited by t h e s u p p l y o f o x y g e n t o the organic species ads o r b e d o n the p l a t i n u m phase, since this m u s t t a k e place across a phase bound. ary, b u t t h e t h e o r e t i c a l limit will be the rate o f a d s o r p t i o n of t h e organic species o n p l a t i n u m at zero coverage. The activity o f h o m o g e n e o u s alloys c o u l d n o t reach this limit because t h e rate of a d s o r p t i o n of t h e organic species will be e x p e c t e d t o decrease with the a d d i t i o n of t h e alloying metal. This will result in a l o w e r rate at l o w coverages a n d suggests t h a t a h e t e r o g e n e o u s s y s t e m w o u l d p r o v i d e t h e o p t i m u m electrocatalyst. A similar e n h a n c e m e n t o f activity of binary catalysts is t o be e x p e c t e d for the o x i d a t i o n o f organic species, like f o r m i c acid, which p r o c e e d t h r o u g h a parallel path, i n h i b i t e d b y the a d s o r b e d i n t e r m e d i a t e , Bad s. T h e increase in the rate o f r e m o v a l o f Bad~, due t o i n t e r a c t i o n with t h e m o s t active o x y g e n species associated with the a d d e d e l e m e n t , will decrease t h e coverage of t h e b l o c k i n g i n t e r m e d i a t e a n d h e n c e lead t o an e n h a n c e m e n t in the rate o f t h e over-all o x i d a t i o n process. ACKNOWLEDGEMENTS T h e a u t h o r s wish to t h a n k Mr. H. B e c k e r f o r his assistance in the experim e n t a l w o r k and Dr. K.J. C a t h r o f o r analyses o f t h e m i x e d m e t a l co-deposits. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13
S. Gilman, Trans. Faraday Soc., 61 (1965) 2546, 2561. H. Wroblowa, B.J. Piersma and J. O'M. Bockris, J. Electroanal. Chem., 6 (1963) 401. S. Gilman, J. Phys. Chem., 66 (1962) 2657. V.S. Bagotzkii and Yu.B. Vasil'ev, Electrochim. Acta, 12 (1967) 1323. T. Biegler and D.F.A. Koch, J. Electrochem. Soc., 114 (1967) 904. D.F.A. Koch in J.A. Friend and F. Gutmann (Eds.), Proc. First Australian Conference on Electrochemistry, 1963, Pergamon Press, Oxford, 1965, p. 657. S.B. Brummer and A.C. Makrides, J. Phys. Chem., 68 (1964) 1448. B.I. Podlovchenko and E.P. Gorgonova, Dokl. Akad. Nauk SSSR, 156 (1964) 673. V.S. Bagotzkii and Yu.B. Vasil'ev, Electrochim. Acta, 11 (1966) 1439. M.W. Breiter, J. Electroanal. Chem., 14 (1967) 407. T. Biegler, J. Phys. Chem., 72 (1968) 1571. A. Capon and R. Parsons, J. Electroanal. Chem., 44 (1973) 1, 239; 45 (1973) 205. O.A. Petrii, B:I. Podlovchenko, A.N. Frumkin and Hira Lal, J. Electroanal. Chem., 1(} (1965) 253.
86 14 B.I. Podlovehenko, A.N. Frumkin and V.S. Stenin, Elektrokhimiya, 4 (1968) 339. 15 V.S. Bagotzkii, Yu.B. Vasil'ev, O.A. Khazova and S.S. Sedova, Electrochim. Acta, 16 (1971) 913. 16 S.S. Beskorovainaya, Yu.B. Vasil'ev and V.S. Bagotzkii, Elektrokhimiya, 8 (1972) 376. 17 V.F. Steuin, V.E. Kazarinov and B.I. Podlovchenko, Elektrokhimiya, 5 (1969) 442. 18 M.W. Breiter, Electrochemical Processes in Fuel Cells, Springer-Verlag, New York, 1969, Ch. X, p. 166. 19 M.A. Khrizolitova, L.A. Mirkind, Yu.B. Vasil'ev, M.Ya. Fioshin and V.S. Bagotzkii, Elektrokhimiya, 8 (1972) 548. 20 B.I. Podlovchenko and R.P. Petukhova, Elektrokhimiya, 9 (1973) 273. 21 T. Biegler, Aust. J. Chem., 26 (1973) 2571, 2587. 22 D.F.A. Koch, Australian Patent No. 46,123, 1964. 23 K.J. Cathro, Electrochem. Technol., 5 (1967) 441. 24 K.J. Cathro, J. Electrochem. Soc., 116 (1969) 1608. 25 O.J. Adlhart and A.J. Hartner, Proc. 20th Power Sources Conference, PSC Publications, New Jersey, 1966, p. 11. 26 M. Watanabe and S. Motoo, J. Electroanal. Chem., 60 (1975) 259, 267, 275. 27 M. Watanabe, T. Suzuki and S. Motoo, Denki Kagaku, 40 (1972) 205, 210. 28 C.E. Heath, Proc. Int. Congr. Fuel Cells, Brussels, 1965, S.E.R.A.I.-COMASCI, 1965, p. 99. 29 J.A. Shropshire, J. Electrochem. Soc., 112 (1965) 465. 30 D.A.J. Rand and R. Woods, J. Electroanal. Chem., 36 (1972) 57. 31 D.A.J. Rand and R. Woods, J. Electroanal. Chem., 44 (1973) 83. 32 D.A.J, Rand and R. Woods in M.W. Breiter (Ed.), Proc. Symposium on Electrocatalysis, 1974, The Electrochemical Society, New Jersey, p. 140. 33 K.J. Cathro, D.F.A. Koch and H.R. Skewes, Mech. Chem. Eng. Trans., MC3 (1967) 17. 34 K.J. Cathro and C.H. Weeks, Energy Convers., 11 (1971) 133. 35 D.A.J, Rand and R. Woods, J. Electroanal. Chem., 31 (1971) 29. 36 T. Biegler, D.A.J. Rand and R. Woods, J. Electroanal. Chem., 29 (1971) 269. 37 E. Gileadi and B.E. Conway in J.O'M. Bockris and B.E. Conway (Eds.), Modern Aspects of Electrochemistry, No. 3, Butterworths, London, 1964, Ch. 5. 38 T. Biegler and R. Woods, J. Phys. Chem., 73 (1969) 3502. 39 M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solution, Pergamon Press, Oxford, 1966. 40 M.W. Breiter, J. Phys. Chem., 69 (1965) 3377.