Journal ofElectroanalytical
Chemistry
Journal of Electroanalytical Chemistry 567 (2004) 139–149
www.elsevier.com/locate/jelechem
Temperature dependence of the COads oxidation process onPt(1 1 1), Pt(1 0 0), and Pt(1 1 0) electrodes
Enrique Herrero a, Bernab�e �Alvarez a, Juan M. Feliu a,*, Sonia Blais b,Zorana Radovic-Hrapovic b, Gregory Jerkiewicz c,*
a Departamento de Qu�ımica F�ısica, Universidad de Alicante, Apdo. 99, Alicante E-03080, Spainb D�epartement de Chimie, Universit�e de Sherbrooke, 2500 boul. Universit�e, Sherbrooke, QC., Canada J1K 2R1
c Department of Chemistry, Queen�s University, 90 Queens Crescent, Kingston, Ont., Canada K7L 3N6
Received 15 July 2003; received in revised form 4 December 2003; accepted 15 December 2003
Available online 28 February 2004
Abstract
The influence of temperature variation on the COads oxidative desorption at Pt(1 1 1), Pt(1 0 0), and Pt(1 1 0) electrodes in 0.5 M
H2SO4 and 0.1 M HClO4 solutions is examined. A distinct shift of the COads stripping peak towards less-positive potentials is
observed as the temperature is increased from 273 to 333 K. Despite the displacement of the desorption peak towards lower po-
tentials, its current density and shape remain unaffected for the Pt(1 1 1) and Pt(1 0 0) electrodes, therefore indicating that the ox-
idation mechanism for the COads is not influenced by the temperature variation in this particular range. On the other hand, the
morphology of the peak on the Pt(1 1 0) electrode changes significantly with the temperature, as a result of the interaction with the
oxide formation. The activation energies for the oxidation processes in the two media have been evaluated through the dependence
of the peak potential and the voltammetric current of the peak with the temperature. The activation energy obtained in H2SO4 is ca.
15–25 kJmol�1 higher than that in HClO4, highlighting the role of the anions in the oxidation process.
� 2003 Elsevier B.V. All rights reserved.
Keywords: Pt(1 0 0); Pt(1 1 1); CO overlayer; Electro-oxidation; Oxidative desorption; Temperature dependence; Surface kinetics; Activated complex;
Activation energy
1. Introduction
Carbon monoxide is a small and simple molecule that
chemisorbs strongly on platinum and other noble met-
als, whose adsorption behavior is structure-dependent
[1]. From the point of view of applied research, CO is animportant surface species because it poisons the noble
metal-based electrocatalysts used in low-temperature
fuel cells. The CO is formed as an intermediate in the
oxidation of methanol or is present in low concentra-
tions in the reformed gas used in fuel cells and therefore
its behavior on platinum electrodes is a subject of in-
tense research.
* Corresponding authors. Tel.: +1-613-533-6413; fax: +1-613-533-
6669 (Gregory Jerkiewicz), +34-965-909301 (Juan M. Feliu).
E-mail addresses: [email protected] (J.M. Feliu), grego-
[email protected] (G. Jerkiewicz).
0022-0728/$ - see front matter � 2003 Elsevier B.V. All rights reserved.
doi:10.1016/j.jelechem.2003.12.019
It is well established that the oxidation of adsorbed
CO (COads) on platinum single-crystal electrodes occurs
at potentials above 0.75 V [2,3], potentials that are much
higher than the standard potential of the process (ca. 0.0
V). The COads electro-oxidative desorption requires the
presence of oxygen-containing species such as H2O orOH� depending on the pH of the aqueous electrolyte
and the reaction takes place according to the general
reaction scheme:
PtþH2O¢Pt–OHþHþ þ e�; ð1Þ
Pt–OHþ Pt–CO ! CO2Hþ þ e� þ 2Pt: ð2Þ
In this scheme, the exact nature of the species involved in
reaction (1) is still not clear, although research on ad-atom-modified Pt(1 1 1) electrode surfaces points to the
adsorbed OH as the species involved in the process
[4,5]. Reaction (2) occurs through a Langmuir–Hinshel-
wood (L–H)-type mechanism. The kinetics of the COads
140 E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149
oxidation depend on several factors, namely, the COads
coverage and surface structure, the mobility of the
COads on the electrode surface, the presence of defects on
the surface, and the anions present in the solution.
Of all these factors that affect the kinetics of COads
oxidation, the surface structure of the COads layers
on platinum single crystals has been extensively char-
acterized by electrochemical techniques [2,3,6–8], IR
spectroscopy [9–16], in situ STM [17–19], coupled UHV-
electrochemistry techniques [20,21], and in situ grazing
incident X-ray diffraction (GIXD) [22–24]. On the
Pt(1 1 1) electrode, three distinct ordered structures were
observed using STM, namely (2� 2)-HCO ¼ 0:75[17,18], (
p19�p
19) R23.4�-HCO ¼ 0:68 [17,18], and
(p7�p
7) R19.1�-HCO ¼ 0:56 [17] and confirmed by
GIDX [22–24]. The COads surface structure is also
known to be potential-dependent and at E > 0:5 V,
STM studies show that a structural transformation
takes place leading to formation of the (p19�p
19)
R23.4�-HCO ¼ 0:68 adlayer [17,18]. The transformation
from the (2� 2)-HCO ¼ 0:75 structure to the(p19�p
19) R23.4�-HCO ¼ 0:68 structure appears to
be associated with a partial oxidation of the more
compact layer and is related to the presence of surface
defects [16].
The kinetics of CO oxidation on Pt single-crystal
electrodes have been studied extensively in recent years
[6,25–32]. These studies have revealed that the CO oxi-
dation on Pt(1 1 1) and vicinal electrodes takes place ac-cording to the mean field L–H-type mechanism [29–31].
Since the oxidation process takes place at very localized
sites, where OHads and COads species can interact, the
mean field L–H-type mechanism is fulfilled only if CO
diffusion on the surface is fast [33]. It has also been ob-
served that defects play a critical role in the oxidation
process. An extrapolation of the rate constants for the
COads oxidation obtained with the Pt(1 1 1) vicinal elec-trodes to the ‘‘ideal’’ (defect-free) Pt(1 1 1) electrode in-
dicates that the oxidation on real Pt(1 1 1) electrodes
takes place mainly at the defect sites [31].
In order to provide a further mechanistic insight
into the COads oxidative desorption, we engaged in
temperature-dependent research on Pt(1 1 1) in 0.5 M
aqueous H2SO4. The only previous study of tempera-
ture effects on CO oxidation was performed for twoother systems, namely for Pt(1 1 1) in aqueous NaOH
with CO present in solution [27] and for Pt(1 1 1) in
aqueous H2SO4 again with CO in solution [28]. These
measurements allowed us to examine the T-dependence
of the COads surface coverage, the relation between the
peak potential and T , and subsequently to elucidate the
apparent activation energy of the process. The latter
value sheds light on the nature of the activated com-plex, and thus on the mechanism of the process. This
paper is a continuation of our previous research on the
Pt–COchem system.
We present new experimental results for CO ad-
sorption and oxidative desorption at the Pt(1 1 1),
Pt(1 0 0), and Pt(1 1 0) electrodes in aqueous H2SO4 and
HClO4. Comparative analysis of these results leads to an
assessment of the influence of the surface geometry,electrolyte nature, and concentration on the CO oxida-
tive stripping through the determination of kinetic pa-
rameters of the process.
2. Experimental
The Pt(1 1 1), Pt(1 1 0) and Pt(1 0 0) electrodes wereobtained from small (2 mm in diameter) single-crystal
beads according to the well-established procedure [34].
The crystals were oriented, cut, and polished with di-
amond paste down to 0.25 lm; then they were an-
nealed in a gas-oxygen flame for 30 min. Prior to each
experiment, the electrodes were flame-annealed and
cooled in an Ar+H2 atmosphere followed by attach-
ment of a water droplet [35]. Subsequently, a cyclicvoltammetry (CV) profile of the electrode was recorded
in 0.5 M H2SO4 or 0.1 M HClO4 at each temperature
in order to verify the surface order and the solution
cleanliness.
The chemisorption of CO was carried out at 0.1 V
(RHE) for 1 min in 0.5 M aqueous H2SO4 and 0.1 M
HClO4 solutions, respectively, containing CO. After-
wards, the excess CO present in solution (COsol) wasremoved by purging with Ar for ca. 10 min. The strip-
ping of COads was accomplished by the application of a
single, positive-going CV scan from 0.05 to 1.00 V. A
CV profile following the COads oxidative desorption was
also recorded and was compared to that preceding the
CO chemisorption. No difference was observed between
the initial profile prior to CO adsorption and that ob-
tained after CO oxidation, thus confirming that (i) allCOads had been completely oxidatively desorbed and (ii)
dissolved CO (COsol) had been effectively removed from
the electrolyte.
The temperature-dependent measurements were per-
formed by immersing the cell in a water bath (Haake),
with the temperature controlled to within 0.5 K. The
water level in the bath was maintained above that of the
electrolyte in the cell to ensure uniformity of the tem-perature distribution. The experiments were conducted
between 273 and 333 K with a 5� interval.All potentials were measured versus a reversible hy-
drogen electrode maintained at the same temperature as
the working electrode (RHE, T). In order to compare
the potentials for the COads oxidation at different tem-
peratures and to obtain the apparent rate constant, all
potentials were converted to a common standard scale,the standard hydrogen electrode at 298 K (SHE, 298 K).
This was achieved by using the following equations
[36,37]:
-100
-50
0
50
100
0.0 0.2 0.4 0.6 0.8 1.0
0.0 0.2 0.4 0.6 0.8 1.0
-200
-150
-100
-50
0
50
100
150
200
E/V (RHE)
j/µA
cm
-2
Pt(111)
E/V (RHE)
278 K 293 K 308 K
j/µA
cm
-2
Pt(100)
Fig. 1. Evolution of the initial voltammetric profile with temperature
for Pt(1 1 1) and Pt(1 0 0) electrodes in 0.5 M H2SO4. Scan rate 50
mV s�1.
E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149 141
ESHE;T ¼ ERHE;T � 0:021 V; ð3Þ
ESHE;298 ¼ ESHE;T þ ðT � 298ÞoE=oT ; ð4Þwhere oE=oT is the variation of the potential of the SHE
measured versus a SHE at 298 K with the temperature
and has a value of 8.4� 10�4 VK�1 [36].
The COads coverage is defined as the number of COads
molecules per number of platinum surface atoms. TheCOads coverage was determined from the total COads
stripping charge density (measured between 0.35 and
0.93 V) corrected for the double-layer charging and for
the charge density associated with the anion re-adsorp-
tion [3]. This correction is known to provide coverage
values that are in agreement with the coverage deter-
mined by other techniques [17,18,38]. As the anion ad-
sorption and double layer charging on the Pt(1 1 1) andPt(1 0 0) electrodes are only slightly affected by T vari-
ation [39,40], the charge density correction for the entire
T range was the same and it stood at 144 lCcm�2 for
the Pt(1 1 1) and 120 lCcm�2 for the Pt(1 0 0) electrodes,
respectively. The COads coverage was calculated as
follows:
HCO ¼qCO2q1e
; ð5Þ
where qCO is the corrected charge density for COads
stripping, 2 is the number of electrons exchanged in the
COads oxidative desorption, and q1e is the charge densityassociated with transfer of 1 electron per surface atom.
The values of q1e for Pt(1 1 1), Pt(1 0 0), and Pt(1 1 0)electrodes are 241, 209, and 150 lCcm�2, respectively.
All solutions were prepared using ultra-pure water
(Nanopure, BrandsteadTM) and BDH Aristar grade
H2SO4 and Merck Suprapur HClO4. All gases were of
ultra-high purity (Ar, 99.999%; CO, 99.997%; and H,
99.999%, Air Products).
3. Results
Fig. 1 shows the initial CV profiles of the Pt(1 1 1) and
Pt(1 0 0) electrodes in 0.5 M H2SO4 at three different
temperatures that are representative of the overall ten-
dency. As can be seen, there are only minor changes in
the CV transients and they are in agreement with pre-
viously published results [39,41]. However, the totalcharge associated with the anion and hydrogen ad-
sorption on both electrodes remains the same.
In the case of the Pt(1 1 1) electrode, the CV features
associated with the anion adsorption/desorption (be-
tween 0.35 and 0.6 V) move towards more positive po-
tentials on the RHE scale. In fact, the onset for anion
adsorption remains at approximately the same potential
(ca. 0.33 V), but the final part of the adsorption process,which is characterized by the sharp spike associated with
the phase transition in the (bi)sulfate layer, moves to
more positive potentials. These results show that the
(bi)sulfate adsorption requires a wider potential window
to be completed at high temperatures, an indication that
the adsorption process is less favored at high tempera-
tures. On the other hand, the spike becomes sharperwith the temperature increase, thus indicating that the
kinetics of the phase transition from the disordered
(bi)sulfate adlayer to the ordered one become faster. The
effect on the hydrogen adsorption process on this elec-
trode has been discussed elsewhere [42,43].
In the case of the Pt(1 0 0) electrode, the peak at ca.
0.37 V, associated with the concurrently occurring hy-
drogen desorption and (bi)sulfate adsorption (or viceversa) processes [44], moves towards less-positive nega-
tive potentials upon T increase. This means that the
strength of the (bi)sulfate adsorption relative the to the
hydrogen adsorption process is higher as the tempera-
ture is raised. However, nothing can be said with respect
to the absolute values, since the respective signals
overlap due to the competitive nature of the processes. It
142 E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149
is conceivable that both processes become less favored
with the temperature increase, with H adsorption being
affected more.
In order to examine the kinetics of the CO oxidation
process in relation to the geometry of the Pt electrode, theCO adsorption and oxidation experiments were carried
out over a wide T range and at well-defined surfaces. In
Fig. 2, we showCOads oxidative desorption CV transients
in 0.5 M aqueous H2SO4 for the Pt(1 0 0) and Pt(1 1 1)
electrodes. In each case, the process is manifested by a
well-defined anodic peak. We observe that the COads
oxidation peak moves towards less positive potentials on
the RHE scale as the temperature is raised. In addition,the morphology of the oxidation peak remains un-
changed for temperatures up to 323 K, i.e., the peaks
have (within the reproducibility of the experiment) the
same maximum current density and full width at half
height. This fact indicates that there are no changes in the
oxidation mechanism within this temperature range.
It is also important to verify whether the initial con-
ditions for the oxidation experiments are the same forthe entire set of experiments. This is especially important
in the case of the Pt(1 1 1) electrode since three different
0.7 0.8 0.9 1.0
0
200
400
600
800
1000
1200
1400
0
200
400
600
800
1000
1200
1400
0.7 0.8 0.9 1.0
j/µA
cm
-2
E/V (RHE)
E/V (RHE)
j/µA
cm
-2
278 K 293 K 308 K
Pt(111)
Pt(100)
Fig. 2. Evolution of the voltammetric profile with temperature for the
COadsoxidation for Pt(1 1 1) and Pt(1 0 0) electrodes in 0.5 M H2SO4.
Scan rate 50 mV s�1.
ordered adlayers have been reported for COads at room
temperature (T ¼ 293–298 K), namely (2� 2) at
HCO ¼ 0:75 [17,18], (p19�p
19) R23.4� at HCO ¼ 0:68[17,18], and (
p7�p
7) R19.1� at HCO ¼ 0:56 [17]. The
stability of these different structures is likely to be af-fected by T . Therefore, if kinetic information is to be
extracted from the COads oxidative desorption tran-
sients, then they have to correspond to the same COads
overlayer having the same coverage. The CO coverage
was determined as a function of T using the CO oxi-
dation charge density and the methodology described in
Experimental (Fig. 3). For the Pt(1 1 1) and Pt(1 0 0)
electrodes, the total charge density is 460 lCcm�2. Thischarge has to be corrected for the anion re-adsorption as
CO becomes desorbed. In the case of the Pt(1 1 1) elec-
trode, the correction charge density amounts to 140
lCcm�2, which gives a values of qCO ¼ 320 lCcm�2.
Using Eq. (5), this value is found to correspond to
a CO coverage of 0.67, a value that agrees well with
the coverage obtained for the (p19�p
19) R23.4� ad-
layer observed by STM after the pre-oxidation wave orunder mild dosing conditions [17,18]. However, as T is
raised further (here 323 6 T 6 333 K), a steep dimi-
nution in the coverage is observed, indicating a change
in the stability of the COads adlayer, and the new
value of the COads coverage is 0.53� 0.02. For the
Pt(1 0 0), the correction amounts to 120 lCcm�2, which
results in qCO ¼ 340 lCcm�2 and a coverage value of
HCO ¼ 0:80, thus similar to the values reported in theliterature [21,45].
Different behavior is observed for the CO oxidation
process at the Pt(1 1 0) electrode, as can be seen in Fig. 4.
270 280 290 300 310 320 330 3400
100
200
300
400
500
q/µC
cm
-2
T/K
Fig. 3. Oxidation charge for the COads oxidation process versus the
temperature for the Pt(1 1 1) (closed symbols) and Pt(1 0 0) (open
symbols) electrodes prior to (squares) and after correction (circles) in
0.5 M H2SO4.
0.0 0.2 0.4 0.6 0.8 1.0
-400
-200
0
200
400
600
800
-400
-200
0
200
400
600
800
0.0 0.2 0.4 0.6 0.8 1.0
j/µA
cm
-2
E/V (RHE)
T=273 K
E/V (RHE)j/µ
A c
m-2
T=333 K
Fig. 4. Initial voltammetric profile for a Pt(1 1 0) electrode prior to CO
adsorption and COads stripping voltammograms after CO adsorption
at two different temperatures in 0.5 M H2SO4. Scan rate 50 mVs�1.
E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149 143
In this case, the oxidation of CO coincides with theonset of the oxidation of the Pt(1 1 0) surface. Since the
surface oxidation is temperature dependent, the co-ex-
istence of these two processes complicates the interpre-
tation of the COads oxidation transients. Specifically, the
morphology of the peak changes significantly with the
temperature, as shown in Fig. 4, indicating a very
complicated dependence on temperature. Since the effect
of the surface oxidation on the CO oxidative desorptionprocess is difficult to account for quantitatively, no ki-
netic studies will be performed for the Pt(1 1 0) electrode.
4. Discussion
Kinetic information regarding the COads oxidative
desorption may be obtained by examining the change ofthe Ep vs. T relation [28]. The overall reaction taking
place at the working electrode is as follows:
Ptð111Þ–COþH2Oþ xA�
! Ptð111Þ–Ax þ CO2 þ 2Hþ þ ð2þ xÞe�; ð6Þ
where A is an anion present in solution that is ad-
sorbed on the electrode surface after CO oxidation.
One of the main objectives of this research is the de-
termination of the activation energy of this reaction. In
order to accomplish this objective, an oxidationmechanism for COads has to be proposed. In this
mechanism, there is an initiation process (precursor),
which may account for 2–3% of the overall charge-
density for process on Pt(1 1 1) electrodes [29], and the
main oxidation process, which can be modeled using a
mean-field Langmuir–Hinshelwood equation [29,30]. It
should be mentioned that the transients have also been
modeled with a modification of the nucleation andgrowth model proposed by Love and Lipkowski [25],
which includes an induction time for the oxidation [32].
However, since the mean field equation is simpler, we
have used this latter equation to model the kinetic
behavior. It is also known that the Tafel slope for the
oxidation desorption of CO whose surface coverage is
below 0.68 is ca. 70 mV/decade [6,25,29,30]. Using
these data, the following reaction mechanism can beproposed:
Pt–H2O¢k1
k�1
Pt–OHþHþ þ e� ðfastÞ; ð7Þ
Pt–COþ Pt–OH!k2 Pt–COOH–Pt ðrdsÞ; ð8Þ
Pt–COOH–Pt¢ 2Ptþ CO2 þHþ þ e� ðfastÞ; ð9Þ
Ptþ xA�¢Pt–xAx þ xe� ðfastÞ: ð10Þ
This mechanism satisfies both conditions found exper-
imentally: (i) the rate determining step (rds) is a
chemical reaction after the first electron transfer has
occurred, which should give a Tafel slope close to 60
mV and (ii) it implies the reaction between adsorbed
CO and OH. In the first step, the adsorbed watermolecule dissociates to produce the adsorbed OH
species. The presence of adsorbed water in the sup-
porting electrolyte, here 0.5 M sulfuric acid, at the
potentials at which CO is oxidized can be argued. As
aforementioned, (bi)sulfate anions are adsorbed at any
adsorption site that is not occupied by COads. How-
ever, the adsorbed (bi)sulfate does not preclude the
presence of water in the double layer. There is FTIRevidence pointing to water molecules being included in
the adsorbed layer [46]. Moreover, the structures found
by STM for the adsorbed (bi)sulfate in this potential
range has been interpreted as a water-(bi)sulfate layer
[47], an explanation that has been corroborated by
radiotracer [48] and chronocoulometry experiments
[49]. Therefore, the proposed first step is reasonable
under the present experimental conditions: water mol-ecules are adsorbed at the electrode together with the
(bi)sulfate anions. According to the mean field L–H-
type mechanism, in the second step the two adsorbed
144 E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149
species, COads and OHads, react to form an adsorbed
intermediate, COOHads. The final steps of the process
include the desorption of the COOHads intermediate
species to yield CO2 and the resulting re-adsorption of
anions.For this mechanism, the reaction rate kinetics are
given by the following equation [28]:
m ¼ � dhCOdt
¼ k1k2k�1
expF E � E0ð Þ
RT
� �hCOhH2O
aHþ; ð11Þ
where E is the applied potential, E0 is the standard po-
tential at 298 K for the COads oxidation reaction, and k1,k�1, and k2 are the standard rate constants for the steps
1, )1, and 2, respectively. In this case, hCO is defined as
hCO ¼ HCO
HCO;max
; ð12Þ
where HCO is the CO coverage defined as the number of
COads molecules per surface platinum atom and HCO;max
is the maximum CO coverage value for which all plati-
num adsorption sites have been completely covered. For
the present treatment and in agreement with previously
published data [21,45], we assume that the values of
HCO;max are 0.68 and 0.81 for the Pt(1 1 1) and Pt(1 0 0)
electrodes, respectively, under the experimental condi-
tions described above.
The standard rate constants can be written as afunction of a pre-exponential factor A and an activation
energy DH 6¼
m ¼ � dhCOdt
¼ A1A2
A�1
exp
� DH 6¼
1 þ DH 6¼2 � DH 6¼
�1
RT
!
� expF E � E0ð Þ
RT
� �hCOhH2O
aHþ: ð13Þ
Since the rate constants are referred to the standard
potential of the process, the activation energy will bealso referred to this potential. The current density is
then given by
j ¼ � qCOð þ qanionsÞdhCOdt
¼ qCO þ qanionsð ÞA1A2
A�1
exp
� DH 6¼
1 þ DH 6¼2 � DH 6¼
�1
RT
!
� expF E � E0ð Þ
RT
� �hCOhH2O
aHþ; ð14Þ
where qCO is the charge density required to oxidize the
CO layer and qanions is the charge density transferred in
the anion re-adsorption. The value (qCO þ qanions) is thecharge density determined through the integration of the
CO oxidation peak (without corrections). The peak
current density, jp can be described by the followingequation:
jp ¼qCO þ qanionsð ÞA1A2
A�1
exp
� DH 6¼
1 þ DH 6¼2 � DH 6¼
�1
RT
!
� expF Ep � E0� �
RT
� �hH2O;phCO;p
aHþ; ð15Þ
where the subscript p indicates the values at the peak
potential, Ep. Solving the equation for Ep leads to
Ep � E0 ¼ RTF
lnjpA�1aHþ
qanions þ qCOð ÞA1A2hCO;p; hH2O;p
þ DH 6¼1 þ DH 6¼
2 � DH 6¼�1
F: ð16Þ
The first term in this equation depends on T and other
variables and the second one is a function of the acti-vation energy. All the variables in the first term can be
considered constant in the temperature range between
273 and 318 K because (qCO þ qanions) and jp remain
constant for both electrodes, as is evident from the data
shown in Fig. 2. Moreover, the integrated charge be-
tween the onset of COads oxidation and Ep is constant,
indicating that the CO coverage has a constant value at
the peak potential. Assuming that the water coverage isconstant at that peak potential, the relation of Ep � E0
vs. T should give a straight line and its intercept leads to
the determination of the activation energy. The standard
potential for the bulk CO/CO2 process in aqueous so-
lutions is )0.106 V [50]. In order to calculate the stan-
dard potential for the COads/CO2 process from this
value, the Gibbs energy of adsorption for CO on Pt is
required. Since this value is not known, only an estimateof its value can be made from the voltammetric behavior
of the CO/CO2 pair on platinum. COads is oxidized at a
very low rate at potentials higher than 0.3 V, and the
reduction of CO2 to COads is achieved only at potentials
below 0.25 V [51]. Therefore, the standard potential for
COads has to remain close to 0 V. Since the exact value of
E0 is not known, a value of 0 V has been used for
practical reasons, i.e., all the activation energies deter-mined in this work are referred to 0 V vs. SHE.
In Figs. 5 and 6, we show Ep vs. T for the Pt(1 0 0)
and Pt(1 1 1) electrodes in aqueous H2SO4 and HClO4.
We find that these relationships are linear, thus, con-
firming the validity of the hypothesis made initially;
namely jp, HCO, and HH2O were assumed to be constant.
We also observe a change in the Ep vs. T behavior at
323 K for Pt(1 1 1) that can be related to the change ofthe initial CO coverage. On the basis of the linear rela-
tionships presented in Figs. 5 and 6, we determined the
activation energy for the COads oxidation and found
that its value was 131� 2 kJmol�1 for Pt(1 1 1) and
139� 5 for Pt(1 0 0) (Table 1).
Another important issue that has to be considered in
the data analysis is how the adsorbed anions affect the
COads oxidation process. Although not explicitly con-sidered in Eq. (5), the presence of anions is known to
270 280 290 300 310 320 330 340
0.70
0.75
0.80
0.85 0.5 M H
2SO
4
0.1 M HClO4
E=1.26-0.00181T
Ep/V
(SH
E,2
98 K
)
T/K
E=1.44-0.00223T
Fig. 6. Ep vs. T for the Pt(1 0 0) electrode in 0.5 M H2SO4 and 0.1 M
HClO4.
260 270 280 290 300 310 320 330 340
0.68
0.70
0.72
0.74
0.76
0.78
0.80
0.82
0.84
0.86
0.88
0.90
E=1.36-0.00177T
0.5 M H2SO
4
0.1 M HClO4
E=1.15-0.00149T
Ep/V
(SH
E,2
98 K
)
T/K
Fig. 5. Ep vs. T for the Pt(1 1 1) electrode in 0.5 M H2SO4 and 0.1 M
HClO4.
E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149 145
affect this reaction. The adsorption process in the
electrochemical environment should be always re-
garded as competitive. Therefore, the presence of a
strongly adsorbed anion such as (bi)sulfate should
modify the rate constants k1 and k�1 since it modifies
the environment in which the reaction takes place. In
order to analyze this effect, the same experiments were
Table 1
Activation energies obtained for the CO oxidation process using the values
E
Pt(1 1 1) 0.1 M HClO4 1
0.5 M H2SO4 1
Pt(1 0 0) 0.1 M HClO4 1
0.5 M H2SO4 1
conducted in the absence of strongly adsorbing anions
(in 0.1 M HClO4). In the case of this electrolyte, the
CV behavior of the COads oxidation process is com-
parable to that depicted in Fig. 2 and the CO coverages
are the same as those obtained in 0.5 M H2SO4 forboth electrodes. The results obtained for this medium
are shown in Figs. 5 and 6 as open symbols. As can be
seen, in the case of HClO4, the peak potential values
are lower than those obtained in 0.5 M H2SO4. For the
Pt(1 1 1) electrode, the measurements were carried out
at a scan rate of 20 mV s�1 in order to avoid inter-
ference with the butterfly that appears in the same
potential range. For that reason, the peak potentialshave values ca. 100 mV lower with respect to the sul-
furic acid medium. If the experiments were carried out
at the same scan rate, then the peak potential would be
reduced by ca. 50 mV, a value comparable to that
obtained for the Pt(1 0 0) electrode. From the intercept,
we again determined activation energy values of
111� 5 kJmol�1 for the Pt(1 1 1) and 122� 5 kJmol�1
for the Pt(1 0 0) electrode. These values are ca. 20kJmol�1 lower than those obtained in the sulfuric acid
medium, thus giving a clear indication of the anionic
effect on the reaction kinetics.
The determination of the activation energy using the
peak potential may present several problems that could
lead to erroneous results: (1) although jp is fairly con-
stant, the real values are scattered around 10% of the
mean value (see Fig. 6); (2) jp and Ep are extremelysensitive to the initial conditions of the electrode and
even a small number of surface imperfections can result
in a slightly different value of the activation energy; and
(3) the activation energy is obtained from the intercept
at T ¼ 0 K, a temperature which is far away from the
working temperature. In order to check the validity of
previous results, a new approach will be used to calcu-
late the activation energies. In this approach, the entireCV profile corresponding to the CO oxidation process
will be used to calculate an apparent constant rate.
Then, through the analysis of changes of the constant
rate as a function of temperature, an apparent activation
energy can be obtained.
In order to obtain the apparent constant rate, the
shape of the voltammetric peak has to be simulated.
According to the model shown in Appendix A, thecurrent density of the CV peak is given by the following
formula:
EpðE 6¼;1Þ and kappðE 6¼;2Þ6¼;1 (kJmol�1) E 6¼;2 (kJmol�1)
11� 5 106� 2
31� 2 134� 2
22� 5 125� 6
39� 4 135� 4
146 E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149
j ¼ qCOð þ qanionsÞkapp exp FE
RT
� �exp � kappRT
F m exp FERT
� �þ b
� �exp � kappRT
F m exp FERT
� �þ b
� �þ 1
� �2ð17Þ
where kapp is the apparent rate constant (and is related to
k1, k�1, and k2), m the scan rate, and b is a parameter that
groups all the initial values for the process (Eini and
hCO;ini). This equation will be used to fit the CV peaks
using (qCO þ qanions), kapp, and b as parameters. As can be
seen, this method is not affected by the aforementioned
sources of error, since the initial conditions of the elec-
trode are parameters within the model and they do notrely on constant jp values.
From the values of kapp, the activation energy can be
calculated as (see Appendix A):
Rd ln kapp� �
dT�1ffi DH 6¼
1 þ DH 6¼2 � DH 6¼
�1: ð18Þ
Therefore, the slope of the representations of the
logðkappÞ vs. T�1 relation will be used to calculate the
activation energy.Fig. 7 shows the experimental CV transients and the
corresponding fitted curves using adjustable parameters
for COads oxidative desorption in 0.5 M H2SO4 for the
Pt(1 1 1) electrode at different temperatures. As can be
seen, the model reproduces the experimental CV tran-
sient well. A similar fitting procedure was applied to the
experimental results obtained for the Pt(1 1 1) in 0.1 M
HClO4 and for Pt(1 0 0) in both media, and a very goodagreement with the fitted data was found. It is important
to emphasize that this model is based on the fact that the
kinetics of CO oxidation at these electrodes follow the
0.76 0.78 0.80 0.82 0
0
200
400
600
800
1000
1200 293303 K
313 K
j/µA
cm
-2
E/V (S
Fig. 7. Comparison between the COads stripping peak (symbols) and the fitt
H2SO4. Scan rate 50 mVs�1.
mean field L–H-type mechanism. For the Pt(1 1 1) elec-
trode, the extrapolation of the results obtained with a
Pt(1 1 1) vicinal electrode for CO oxidation to the be-
havior of the Pt(1 1 1) electrode indicates that the oxi-
dation takes place mainly at the defect sites of thesurface [31]. In a process like this, in which the oxidation
process takes place at very localized sites, the mean field
L–H-type mechanism is fulfilled only if CO diffusion on
the surface is fast. The measurements performed in
UHV indicate a high mobility of CO on the Pt(1 1 1)
surface [52–55]. For the Pt(1 0 0) electrode, IR mea-
surements also indicate a high CO mobility [56]. The
good agreement of the fitted and experimental resultsalso supports the view of high mobility of COads on the
platinum surfaces.
The values of kapp vs T�1 obtained for the Pt(1 1 1)
and Pt(1 0 0) electrodes are shown in Figs. 8 and 9. As
can be seen, the rate constants in 0.1 H HClO4 are lower
than those obtained in 0.5 M H2SO4, as would be ex-
pected on the basis of the peak potentials of the CV
transients. It should be mentioned that the deviationsfrom linearity in Figs. 8 and 9 are smaller than those
observed in Figs. 5 and 6. Very small changes in the
initial coverage lead to significant changes in the peak
potential (>10 mV), increasing the errors and deviations
in Figs. 5 and 6. On the other hand, this second ap-
proach accounts for the influence of the initial CO
coverage in the voltammetric behavior since hCO;ini is
included in b as a fitting parameter, resulting in a moreaccurate determination of the activation energy.
The values of the activation energy obtained are,
within experimental error, the same as those obtained
with the previous approach (Table 1). As can be seen,
.84 0.86 0.88 0.90 0.92
283 K K
273 K
HE, 298 K)
ings obtained from Eq. (17) (lines) for the Pt(1 1 1) electrode in 0.5 M
0.0036 0.0034 0.0032 0.003010-16
10-15
10-14
10-13
10-12
10-11
10-10
10-9
10-8
0.5 M H2SO
4
0.1 M HClO4
k app/s
-1
T-1/K-1
Fig. 8. kapp vs. T�1 for for the Pt(1 1 1) electrode in 0.5 M H2SO4 and
0.1 M HClO4.
0.0038 0.0036 0.0034 0.003210-15
10-14
10-13
10-12
10-11
10-10 0.5 M H2SO
4
0.1 M HClO4
k app/s
-1
T-1/K-1
Fig. 9. kapp vs. T�1 for for the Pt(1 0 0) electrode in 0.5 M H2SO4 and
0.1 M HClO4.
E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149 147
there is a difference of ca. 15–20 kJmol�1 in the acti-
vation energy when bisulfate (a strongly adsorbed an-
ion) is present in the solution. In perchloric acid, the
activation energy shows a small dependence on the
electrode surface, being, for the Pt(1 1 1) electrode, ca. 10
kJmol�1 lower than that obtained for the Pt(1 0 0)electrode, indicating that the surfaces have different in-
teraction energies with the adsorbed species (COads and
COOHads). However, in the presence of sulfuric acid, the
activation energies are comparable, indicating that the
effect of the adsorbing anion has compensated these
differences in the interaction energies.
Density functional theory (DFT) has been used to
evaluate the activation energy of the COads reaction. For
step (8), the activation energy calculated is DH 6¼2 ¼ 0:44
eV [57]. Since the activation energy we have determined
in this work corresponds to DH 6¼1 þ DH 6¼
2 � DH 6¼�1 , the
value of DH 6¼1 � DH 6¼
�1 has to be estimated in order to
compare the experimental and theoretical results. From
basic kinetics, the difference between DG6¼1 � DG 6¼
�1 is
related to the equilibrium constant of the reaction. For
an electrochemical step this relation can be written as
DG6¼1 � DG 6¼
�1 ¼ F ðE01 � ErefÞ; ð19Þ
where E01 is the standard potential of this step and Eref is
the potential at which the activation energy is calcu-
lated. Since all the rate constants determined in this
manuscript are referred to the standard potential of the
overall reaction, Eref for Eq. (19) should be E0 whichhave been taken as very close to zero for practical rea-
sons. The value of E01 ¼ 0:57 V has been also estimated
by DFT, assuming that the entropy contributions to the
Gibbs energy are negligible [58] and therefore
DH 6¼1 � DH 6¼
�1 � DG6¼1 � DG6¼
�1 ¼ F ðE01 � ErefÞ
¼ 55 kJmol�1: ð20Þ
By adding the contributions, the total activation energy
calculated using the DFT data would be ca. 97 kJmol�1,
very close to the value determined experimentally.
Acknowledgements
An acknowledgment is made to the NSERC of
Canada (Discovery Grant) and Ministerio de Educacion
y Cultura, and DGES (Grant No. BQU2003-4029) for
support of this research project.
Appendix A
The current for the CO oxidation process can becalculated using
j ¼ � qCOð þ qanionsÞdhCOdt
: ðA:1Þ
The value dhCO=dt is defined according to the kinetics
law governing the process. In this case, the mean field
Langmuir–Hinshelwood equation replicates the experi-
mental behavior observed in potentiostatic CO oxida-tion experiments [27,30]. According to the detailed
mechanism proposed, the reaction rate can be written as
(Eq. (13))
m ¼ � dhCOdt
¼ k1k2
k�1
expF E � E0ð Þ
RT
� �hCOhH2O
aHþ: ðA:2Þ
148 E. Herrero et al. / Journal of Electroanalytical Chemistry 567 (2004) 139–149
The water coverage is not known, but it will be pro-
portional to the fraction of the electrode surface not
covered by CO
hH2O ¼ K 1ð � hCOÞ; ðA:3Þ
K being the proportionality constant. This proportion-
ality constant is very dependent on the anion present in
the solution, i.e., the stronger the anion adsorption, the
lower is the proportionality constant obtained. Addi-
tionally, in a voltammetric experiment, the applied po-
tential depends on the time according to
E ¼ Eini þ mt; ðA:4Þ
where Eini is the initial potential and m the scan rate.
Substituting (A.3) and (A.4) in (A.2) yields
m ¼ � dhCOdt
¼ k1k2k�1
K expF Eini þ mt � E0ð Þ
RT
� �hCO 1� hCOð Þ
aHþ
¼ kapp expF Eini þ mtð Þ
RT
� �hCO 1ð � hCOÞ; ðA:5Þ
where the terms (k1; k2; k�1; K; aH3Oþ and E0) in Eq.
(A.5) for a single experiment have been grouped in kappand is equal to
kapp ¼k1k2k�1
KaH3O
þexp
�� FRT
E0
�: ðA:6Þ
In Eq. (A.5), dhCO=dt is a function of t and hCO which
in turn is also function of t. In order to obtain the
current, an explicit relationship between dhCO=dt and thas to be obtained. For this reason, this equation has to
be integrated, solved for hCO and differentiated again toobtain dhCO=dt as a function only of t. The initial con-
ditions for the integration are a CO coverage value of
hCO;ini at t ¼ 0. Thus,Z hCO
hCO;ini
dhCOhCO 1� hCOð Þ ¼ �
Z t
0
kappRTF m
expFRT
Einið�
þ mtÞ�dt:
ðA:7ÞSolving for hCO yields
hCO ¼ 1� 1
exp � kappRTF m exp F
RT Eini þ mtð Þ� �
þ b� �
þ 1;
ðA:8Þwhere b has a value of
b ¼ kappRTF m
expFEini
RT
� �þ ln
hCO;ini
1� hCO;ini
ðA:9Þ
and groups all the initial values (Eini and hCO;ini) for a
given experiment. Under the present conditions of theexperiments, where Eini ¼ 0:06 V (RHE), the term
kappðRT =F mÞ expðFEini=RT Þ is negligible with respect to
lnððhCO;iniÞ=ð1� hCO;iniÞÞ and therefore, b depends only
on the initial CO coverage. Substituting Eq. (A.8) in Eq.
(A.1) gives
j ¼ qCOð þ qanionsÞ
�kapp exp F
RT Eini þ mtð Þ� �
exp � kappRTF m exp F
RT Eini þ mtð Þ� �
þ b� �
exp � kappRTF m exp F
RT Eini þ mtð Þ� �
þ b� �
þ 1� �2 :
ðA:10Þ
Since E ¼ Eini þ mt, the current as a function of the
electrode potential is
j ¼ qCOð þ qanionsÞ
�kapp exp FE
RT
� �exp � kappRT
F m exp FERT
� �þ b
� �exp � kappRT
F m exp FERT
� �þ b
� �þ 1
� �2 :
ðA:11Þ
This expression will be used to fit the voltammetric pro-
files obtained for CO oxidation, where ðqCO þ qanionsÞ,kapp, and b are the fitting parameters.
The values obtained from kapp will be used to estimate
the activation energy of the process. The activation en-
ergy is defined as
DH 6¼ ¼ Rd lnðkÞdT�1
; ðA:12Þ
where k is the rate constant of the process. In this case,
kapp contains, aside from the rate constants of the pro-
cess, other terms. The differentiation of lnðkappÞ with
respect to T�1 gives
Rd ln kapp� �
dT�1¼ R
d ln k1ð ÞdT�1
þ Rd ln k2ð ÞdT�1
� Rd ln k�1ð ÞdT�1
þ Rd ln Kð ÞdT �1
þ Rd ln aH3O
þ� �dT�1
þ FE0
¼ DH 6¼1 þ DH 6¼
2 � DH 6¼�1 þ R
d ln Kð ÞdT�1
þ Rd ln aH3O
þ� �dT�1
þ FE0: ðA:13Þ
Since E0 is close to 0 for the CO oxidation process and
the variations of lnðKÞ and lnðaH3OþÞ with respect to T�1
can be considered negligible in this temperature range,
Rd ln kapp� �
dT�1ffi DH 6¼
1 þ DH 6¼2 � DH 6¼
�1 ðA:14Þ
Therefore, the activation energy of the process can be
obtained from the slope of the representation of
logðkappÞ vs. T�1.
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