Oxygen Reduction Reaction at Pt single crystals: A critical overview
Ana Mª. Gómez–Marína,b, Rubén Rizoa, Juan M. Feliu a,*
a
Instituto de Electroquímica, Universidad de Alicante, Apt. 99, Alicante, E-03080,
Spain.
b
Basic Science Department, Fundación Universitaria Luis Amigó, Transversal 51A
#67B-90, Medellín, Colombia.
Juan M. Feliu-juan.feliu@ua.es. *Corresponding author. Tel.: +34 965 909 301; fax:
+34 965 903 537.
Abstract
Oxygen reduction reaction (ORR) dynamics at platinum single crystal surfaces is
reviewed, and experimental results in acid and alkaline solutions are discussed in the
framework of theoretical studies. Special emphasis is devoted to point out the role of the
surface charge, water structure and adsorbed oxygen containing species. Additionally, a
discussion about the possible relevance of the hydrogen peroxide as intermediate
species has also been included. It is shown that ORR is a complex process, affected by
many different factors and so, neither surface charge nor oxygen-containing species
coverage alone are determining factors of electrode activity. Instead, the structure and
interactions between water and water dissociation products affect the energetics of
adsorption processes. In this way, the nature of adsorbed species, such as H2Oads, OHads,
Oads or PtO oxides may determine the surface reactivity. Finally, if H2O2 is an
intermediate product in the ORR mechanism, it would be crucial to find a proper
catalyst able to effectively reduce it at high potentials and inhibit its oxidation.
Introduction
The oxygen reduction reaction (ORR) is one of the fundamental reactions in electrocatalysis and certainly, constitutes one of the main drawbacks for the development of
fuel cells [1]. On Platinum, the most active pure metal for this reaction, oxygen reduces
to water through a four-electron process with a large overpotential, η~0.3 V [1,2,3,4,5].
Unfortunately, despite many years of research, the ORR mechanism is still unknown
[3,6,7,8], likely because currently available experimental techniques are not able to
detect the intermediate species involved in the reaction [6,7,8,9]. Hence, the best
improvement in the ORR activity, compared to the activity of pure platinum, has been
an overpotential reduction of less than 100 mV [5,10], reported for platinum binary
alloys [11,12,13]. A generally accepted reaction scheme, in which hydrogen peroxide
may be a stable reaction intermediate species, is depicted by [14]:
Scheme 1. Reaction pathways proposed for the ORR given by Wroblowa et al. [14].
Under this scheme, oxygen can directly reduce to water, without detectable intermediate
species:
O2 + 4H + + 4e− ⇄ 2H2 O
E0 = 1.229 V (1)
or follow a serial route, in which hydrogen peroxide is formed as an intermediate
species:
O2 + 2H + + 2e− ⇄ H2 O2
E0 = 0.695 V (2)
Of course, both reaction schemes may also occur concomitantly [15,16]. When H2O2 is
formed, it can be re–oxidized to oxygen (Eq. 2) or reduced to water:
H2 O2 + 2H + + 2e− ⇄ 2H2 O
E0 = 1.763 V (3)
In addition, it can also be transported to the bulk [14,15,16], or disproportionate into
water and oxygen in a non–electrochemical reaction:
1
H2 O2 ⇄ O2 + H2 O
2
ΔG0 ≈ -1.096 eV (4)
From a technological point of view, a full understanding of the fundamental factors
controlling the ORR activity on Pt surfaces is required for the rational design of new
and efficient ORR catalysts. In this perspective, a critical analysis of the current state of
art, combined with recent results from our laboratory at platinum single crystals, is
reported. The use of single crystals simplifies the study and opens the possibility of
correlating specific interfacial properties and the electrochemical processes occurring at
the interface. This information can be further extended to understand molecular
phenomena occurring at Pt nano–particles, which catalyze the ORR in polymer
electrolyte membrane fuel cells, assuming that they can be well described as a collection
of ordered facets having {111} and {100} terraces, as well as step and kink surface sites
[17,18].
Oxygen reduction at Platinum Single Crystals in O2 saturated solutions.
First works about ORR on Pt single crystals began almost 20 years ago [17,19,20,21].
In alkaline or acidic non-adsorbing electrolyte solutions, free of species that could
strongly adsorb on the electrode surface and interfere with species coming from water
adsorption, oxygen reduces to water at Pt(111) surfaces between the reaction onset at
~1.0 V (vs. RHE) and the potential at which hydrogen adsorption (Hads) occurs, Figure
1A [8,17,19,22,23]. At E < 0.3 V, two current drops, together with H2O2 detection
[17,19], indicate that Hads may block surface sites or inhibit O–O bond scission and so,
oxygen reduces to H2O2 in a two-electron process [17,19,22,23]. In agreement, a similar
decrease in current has also been reported during the H2O2 reduction reaction (HPRR)
on Pt(111) in this potential region [24].
0.0
-1/2

-1
3 6 9 12 15 18
-2.0
-2.0
-2
-4.0
400 RPM
-6.0
-3.0
jlim=-0.059-0.44
-4.0
-1/2
900 RPM
-8.0
C 2.0
-5.0
1.0
1600 RPM
"Tafel" slope of
-1
0.077 V decade
-6.0
-7.0
2500 RPM
A
0.0
0.6
0.8
0.0
-2
j / mA cm
-2
B
200 RPM
jlim / mA cm
0 RPM
log(jK/ mA cm )
-1.0
/ rad s
1.0
E / V (RHE)
0.2
0.4
0.6
0.8
1.0
E / V (RHE)
Figure 1. (A) Cyclic voltammetric profile for the oxygen reduction on a hanging
meniscus rotating disk Pt(111) electrode, in oxygen saturated 0.1 M HClO4 solutions.
Scan rate 50 mV s-1. Positive (dashed line) and negative–going (solid line) sweeps. (B)
Limiting current density (Levich equation). (C) Kinetic current density (KouteckyLevich equation).
ORR occurs at high overvoltages but the mass-controlled region appears very soon,
once the electrode potential approaches 0.75 V, Fig 1A. Thus, the potential range for the
activation region, in which the electron transfer mechanism can be really analyzed, is
quite limited to the foot of the wave. At low rotation rates, ω, the sharp peak at ~0.8 V
in oxygen free solutions, the so-called “butterfly” feature assigned to hydroxyl
adsorption, OHads, from water dissociation [25,26,27], superimposes on oxygen
reduction currents, Fig 1A. At faster ω, however, this peak broadens and cannot be
distinguished anymore, suggesting that, if contamination may be discarded, either the
OHads species is an ORR intermediate or the reactant and/or any intermediate species
may modify OHads adsorption dynamics [8]. Contrarily, the sharp spike at ~0.45 V in
H2SO4 solutions, associated with an order-disorder transition in the adsorbed (bi)
sulphate layer superimposes on the ORR current even at 2500 rpm in sulphuric acid
solutions TAMBIEN PASA EN LA REF 23 [22]. It should be recalled that the 0.45 V
spike in sulphuric acid is more sensitive to contamination than the 0.8 V peak in
perchloric acid.
A plot of jlim against ω-½, Fig. 1B, accurately gives a straight line, with a slope close to
the theoretical Levich slope for a four-electron process [8,17,19,22,23]. Thus, the
system hydrodynamic is well described by the Levich equation. Therefore, ORR curves
in Fig. 1A are commonly analyzed by using Koutecký–Levich plots [19,22,23], Fig. 1C.
In doing so, a first–order dependence, regarding O2 concentration, is found and so, the
first charge transfer step has been proposed as the rate-determining step (RDS) in the
overall mechanism [8,19]. However, experimental Tafel slopes, from 60 to 88 mV
between 0.8 to 0.9 V [6,8,17,19], are lower than 120 mV, the intrinsic value for a first
electron transfer as RDS. This deviation is explained in terms of either O2 adsorption
dynamics changes, due to changes in chemisorbed oxygen-containing species coverage
[23,28], or Temkin adsorption conditions for ORR reaction intermediates [3,19]. It has
been also proposed that the first electron transfer is not the RDS but precedes it.
Similar reaction dynamics has been reported for Pt(100), Pt(110) and stepped surfaces
in non-adsorbing electrolyte solutions [6,17,19,22,23,29]. However, the ORR activity is
sensitive to the crystallographic orientation and so, experimental Tafel slopes are larger
than on Pt(111) and the reaction rate (vs. RHE) depends on the solution pH at these
surfaces [17,20,22,23,30,31]. Figure 2A resumes ORR activities, expressed as the ration
between the reduction current at 0.9 V and jlim, j(0.9 V)/jlim, of different Pt single
crystals in non-adsorbing alkaline and acid solutions. Because the sweep direction can
play an important role in the shape of the cyclic voltammogram (CV) [8,23], only the
Pt(11 1 1)
Pt(100)
Pt(544)
0.1 M HClO4
0.1 M NaOH
Pt(211)
0.2
Pt(110)
0.4
Pt(111)
A
Pt(331)
Pt(221)
0.6
0.0
0.15
-40
Pt(11 1 1)
Pt(100)
Pt(211)
Pt(544)
0.30
Pt(111)
B
Pt(331)
Pt(221)
0.45
Pt(110)
PZTC / V (RHE)
j / jlim at 0.9 / V (RHE)
ratio j/jlim in the positive-going scan is given.
-20
0
20
Angle / grad
40
60
Figure 2. (A) Normalized reduction current density, j/jlim, at 0.9 V during the positivegoing scan for the oxygen reduction at different Pt single crystals and (B) Potential of
zero total charge as a function of the angle of the surface normal with respect to the
(111) direction in 0.1 M HClO4 (■) and 0.1 M NaOH (●). Scan rate 50 mV s-1.
The ORR inhibition by Hads, at E < 0.3 V, is also sensitive to the crystallographic
orientation [17,20,22,23,30,31]. In acidic solutions, it is lower on Pt(100) than on
Pt(111) and does not occur at all on Pt(110) [6,17,19]. Moreover, on Pt(100) and
Pt(111) vicinal surfaces the current density drop’s onset is shifted toward more negative
potentials and the total current drop decreases as the step density increases [22,23].
Similar crystallographic dependence has been also found for the HPRR at these latter
surfaces in this potential region [32,33].
From Fig, 2A, it is clear that, except for Pt(111), the ORR activity (vs. RHE) decreases
when moving from acid to alkaline solutions, i.e. while the ORR activity on Pt(111)
changes 0.059 mV per pH unit, the change on all other Pt single crystals is larger. This
is rather surprising and raises the question about why the (111) facet can maintain its
electro-catalytic activity regardless the solution pH, but not the other crystallographic
orientations. At first glance, it could be thought that the difference arrives either because
of changes in the surface electrode charge, since in alkaline solutions the electrode bears
an electronic charge density ~0.7 eV more negative than in acid solutions, or changes in
water structure, due to surface electrode charge and/or local pH changes [31]. Following
this idea, Fig. 2B resumes the experimental potential of zero total charge (PZTC) for
different Pt single crystals [34,35,36,37,38,39]. The role of water surface structure on
ORR activity will be discussed below.
Considering Fig. 2B, the increase on the PTZC of Pt(100) and Pt(110) surfaces in the
RHE scale, when moving from acid to alkaline solutions, can account for the lower
activity towards ORR in basic solution, however for Pt(11 1 1) electrode this is not true.
Moreover, the ORR activity does not apparently follow a systematic tendency with
electrode PTZC values, i.e. despite similar PZTC values for Pt(111) and Pt(110) in
alkaline solutions, or for Pt(331) and Pt(221) in acidic media, the ORR activity differs
at these surfaces. Therefore, electrode surface charge differences between acid and
alkaline solutions cannot completely explain the difference between ORR activities in
these media, and the origin of this fact is still not fully understood.
Other characteristic feature for the ORR at Pt single crystals is the dependence of the
reduction current at potentials close to the onset of the reaction with the direction of the
potential scan [6,8,17,23], and the value of the upper scan limit, Eup [8], in the CV.
Figure 3 shows j/jlim in the negative-going scan at 0.9 V, for Eup = 0.9 and 1.15 V, and
0.8 V, for Eup = 1.60 V, on Pt(111), Pt(221) and Pt(211) in 0.1 M HClO4. In a similar
way to polycrystalline Pt, surface oxides have been suggested to be at the origin of this
current hysteresis [6,17,23]. However, the different behavior between Eup = 0.9 - 1.15
V, and Eup = 1.60 V, and the strong reaction inhibition in this latter case, suggests a
more complex phenomenon. Thus, the formation of an aqueous ORR intermediate
species if Eup ≤ 1.15 V [8], and the reduction of PtO oxide species if Eup > 1.15 V, have
also been proposed for explaining differences between positive and negative-going
0.5
0.3
0.2
Pt(211)
0.4
Pt(111)
Eup = 0.90 V
Eup = 1.15 V
Eup = 1.60 V
Pt(221)
j / jlim at Ef / V (RHE)
scans.
0.1
0.0
-20
-10
0
10
Angle / grad
20
Figure 3. Normalized reduction current density, j/jlim, at 0.8 V (filled) and 0.9 V
(empty), during the negative going scan for the oxygen reduction at different Pt single
crystals and upper limit potentials, Eup, as a function of the angle of the surface normal
with respect to the (111) direction in 0.1 M HClO4 and 1600 RPM. Scan rate 50 mV s-1.
It is important to mention that, after the potential excursion till Eup = 1.15 V, in the
following positive-going scan, both Pt(221) and Pt(111) recover their ORR activity, but
the ORR activity at Pt(211) is lowered. However, after the potential excursion till Eup =
1.60 V, the surface ordering is destroyed, and the ORR activity in the next positivegoing scan at Pt(221) and Pt(111) increases, while the activity of Pt(211) decreases even
more. In consequence, while {110} disordered steps are more active than {110} ordered
ones, {100} disordered steps appear to be less active than the ordered ones and thus
potential excursions at high values, sometimes employed as cleaning procedures for Pt
nano-particles [40], may decrease ORR activity.
In the case of strong adsorbing electrolytes, i.e. solutions of dissolved species that could
strongly adsorb on the electrode surface, such as HSO4-, Cl- and Br- anions, the ORR
dynamics may change. In these cases, both the reaction rate and the main reaction
product, water or H2O2, can vary depending on the crystallographic electrode
orientation, which in turns defines the adsorption dynamics of these ionic species, also
called spectator species, on the surface [6,41,42,43]. Site availability and negative
energetic interactions between spectator and ORR intermediate species have been
suggested to be at the origin of these changes [6,28,41,42], however, this is still an open
question.
ORR activity in different surfaces: Volcano curves
Earlier works concerning ORR on different surfaces have shown the existence of
volcano-type curves when the electrode activity is plotted either as a function of the
oxygen adsorption bond strength, ΔGOads, [4], or of the electronic (Pt d–band vacancies)
and geometric (Pt–Pt bond distance) properties of Pt and Pt alloys [5]. Thus, finding
materials with proper ΔGOads values has been a common approach in the catalysts’
research for improving the ORR rate [5,11,12,13]. In agreement, theoretical studies
have predicted a volcano-type correlation between the reaction rate of any dissociative
reaction, such as the ORR, and the dissociative adsorption energy of the reacting
species, regardless the nature of the rate determining step. This is because the validity of
the Brønsted–Evans–Polanyi relation and the linear relationship between the adsorption
energy of reactive and intermediate species [44].
From a molecular point of view, several ORR mechanisms have been proposed to
explain experimental results [45,46,47,48,49,50,51,52,53,54]. In all of them, Oads and
OHads, and sometimes OOHads, are ORR intermediate species. Therefore, according to
recent theoretical studies, it would not be possible any further improvement in the ORR
performance, beyond a ΔGOads optimal value, because of the existence of a scaling
relationship between ΔGOads and the adsorption bond strength of OHads and OOHads
species, ΔGOHads and ΔGOOHads [45,55,56,57,58,59]. For example, in the “dissociative”
route, in which the O-O bond is broken upon oxygen adsorption, the simplest reaction
scheme is given by [45]
𝑂2 + 2𝑃𝑡 = 2𝑃𝑡𝑂
(5)
𝑃𝑡𝑂 + 𝐻 + + 𝑒 − = 𝑃𝑡𝑂𝐻
(6)
𝑃𝑡𝑂𝐻 + 𝐻 + + 𝑒 − = 𝑃𝑡 + 𝐻2 𝑂
(7)
wherein Oads and OHads are the only intermediate species. Similarly, in the “associative”
route, in which the O-O bond integrity is preserved upon adsorption and would only
break after the electron transfer, reaction steps would be [45]
𝑂2 + 𝑃𝑡 = 𝑃𝑡𝑂2
(8)
𝑃𝑡𝑂2 + 𝐻 + + 𝑒 − = 𝑃𝑡𝑂𝑂𝐻
(9)
𝑃𝑡𝑂𝑂𝐻 + 𝐻 + + 𝑒 − = 𝐻2 𝑂 + 𝑃𝑡𝑂
(10)
𝑃𝑡𝑂 + 𝐻 + + 𝑒 − = 𝑃𝑡𝑂𝐻
(11)
𝑃𝑡𝑂𝐻 + 𝐻 + + 𝑒 − = 𝑃𝑡 + 𝐻2 𝑂
(12)
with OOHads, Oads and OHads as intermediate species. Reduction of OOH* to H2O2,
which would react further in another surface site, instead of eqn. (10), has also been
proposed [46] and, in this case, the associative mechanism can also be termed as
“peroxo” –mechanism [45]. In this latter case, it is usually assumed that formed H2O2
dissociates in two OHads molecules [46,48,47,49,52] and so, again OOHads, Oads and
OHads are intermediate species.
Similarly, theoretical studies have suggested a volcano type response for the ORR
activity as a function of ΔGOHads or ΔGOads on Pt single crystals, with the (111) facet at
the top of this curve [18,60,61,62]. In this case, the theoretical ORR activity has been
calculated from a simple model of the electrode kinetics that only employs
thermodynamic data, according to [45,63]:
𝑗𝑘 (𝑉) = 𝑗̃
𝑙𝑖𝑚𝑖𝑡 𝑒
𝛼(∆𝐺0 −𝑒𝐸)
𝑘𝑇
(13)
with α = 1 being the transfer coefficient, ΔG0 the activation barrier for the rate-limiting
transfer step, equal to the larger of the free–energy differences, estimated from DFT
-2
calculations [45,63], and 𝑗̃
𝑙𝑖𝑚𝑖𝑡 = 96 mA cm , the current density achieved if all surface
reactions are exothermic (i.e. the highest possible turn-over frequency per site in an
electrochemical cell with minimal diffusion limitations), defined by [45,63]
𝑗̃
𝑙𝑖𝑚𝑖𝑡 = 𝑗0 𝑒
𝛼(𝑒𝐸0 −∆𝐺0 )
𝑘𝑇
(14)
Equation (13) would represent an upper bound to the ORR reaction rate. However, if
there is any additional barrier to the proton transfer, or if the coverage is far from ideal,
the reaction rate would be slower [45,63]. Figure 4 depicts theoretical and experimental
ORR activities at 0.9 V, as a function of the oxygen binding energy on Pt(111), Pt(100),
Pt(110), Pt(332) and Pt(211) surfaces, Fig. 4A and B respectively. Theoretical data were
calculated with the model given by eqn. (13) and by employing energetic values
reported from previous DFT calculations [45,55,58,62,64]. For sake of comparison,
experimental results are given for alkaline and acid solutions, Fig. 4B.
Pt(332)
Pt(111)
Pt(100)
-0.4
Pt(110)
-0.2
Pt(110)
A
Pt(211)
Theoretical
Activity (eV)
0.0
0.1 M HClO4
0.1 M NaOH
-0.2
1.3
1.4
1.5
Oxygen binding
energy (eV)
Pt(100)
-0.1
Pt(111)
Pt(110)
0.0
B
Pt(211)
Experimental
Activity (eV)
0.1
Pt(332)
-0.6
1.6
Figure 4. Theoretical (A) and experimental (B) volcano plots for the oxygen reduction
reaction on Pt single crystal (111), (110), (100), (332) and (211) surfaces. The y axis is
kTln(jk). All activities are calculated at a potential of 0.9 V vs. RHE. Theoretical data
are taken from refs. [45,55,58,62] (▲) and [64] (Δ), and theoretical activities, jk, are
calculated from eqn. (13) [45,63].
Following previously published analysis [45,63], the whole ORR theoretical curve for
-2
Pt(111) is calculated considering 𝑗̃
𝑙𝑖𝑚𝑖𝑡 = 96 mA cm , ΔG0 = 0.45 eV and α = 1, curve 1
in Figure 5. However, if surface charging for changing the electrode potential and
double layer effects are included into calculations, other ΔG0 values for eqns. (6) and
(7) can be obtained (see discussion in next section) [65,66]. In these cases, the
theoretical curve significantly differs from the experimental one, curves 3 to 5 in Fig. 5.
In addition, because eqn. (13) is quite sensitive to 𝑗̃
𝑙𝑖𝑚𝑖𝑡 , ΔG0 and α values, similar
adjusted curves can be obtained by assuming a different set of parameters, or including
more complex dynamics [28]. Contrarily, a different adjusted curve can be also
calculated if instead of α = 1, α is taken to be 0.5, curve 2 in Fig. 5.
0.0
j / mA cm
-2
-1.0
-2.0
-3.0
-4.0
-5.0
0.8
1 5 2
6
3
0.9
1.0
1.1
4
1.2
E / V (RHE)
Figure 5. Experimental (solid) and theoretical (dashed) kinetic current densities, jk, for
the ORR. jks are calculated using eqn. (13) and 1) ΔG0 = 0.45 eV, α = 1; 2) 0.45 eV/0.5;
3) 0.22 eV/1; 4) 0.11 eV/1. In calculating 5 (⋆) and 6 (×) the larger of the free–energy
differences for eqns. (6) and (7), according to DFT data from Ref. [65,66], is used ΔG0
= 1,229-min{E0,6, E0,7} (crosses in Fig. 8), with α = 1 and 0.5, respectively.
Theoretical kinetics trends shown in Fig 4A, estimated from DFT calculations taken
from [45,55,58,62] (solid symbols), suggest the sequence Pt(111) >> Pt(100) >>
Pt(110) > Pt(211) for the ORR activity, with surface steps significantly less active for
the reaction than the Pt(111) surface. Experimental results, instead, suggest a different
sequence: Pt(211) > Pt (110) > Pt(111) > Pt (100) for acid solutions, and Pt(111) >
Pt(211) > Pt (110) > Pt (100) for alkaline solutions. Hence, the theoretical results only
agree with experiments in alkaline solutions on predicting Pt(111) on the top of the
volcano curve. Nonetheless, the experimental activity of Pt(211) is still in contradiction
with theoretical predictions, being more active than the other two basal planes, Pt(100)
and Pt(110). Besides, the theoretical activity’s variation between Pt(111) and Pt(211)
surfaces predicted from these data is almost ten times greater than that observed in the
experiments.
Under this theoretical framework, it has been proposed that active sites on nanoparticles are located on the terraces, particularly the (111) facet. More open facets, such
as steps, edges and kinks, provide a negligible contribution to the activity because their
tendency to bind more strongly Oads and OHads species [18,60,61,62]. Discrepancies
between theoretical and experimental activities at Pt stepped surfaces in acidic media,
Fig. 4B, are explained in terms of either a symmetry breaking of the hexagonal
overlayer on the (111) facet because of the steps [61], or the existence of specific
surface terrace sites on stepped surfaces that bind O-containing species more weakly
than Pt(111) [40], related, perhaps, to their surface reconstruction [18] in acidic but not
in alkaline solutions. However, in acid media surface steps increase ORR activity even
on Pt(s)[n(111)x(111)] surfaces with the longest possible terrace widths, although these
latter surfaces have a defect percentage low enough to preserve the (111) long-range
surface ordering and do not suffer faceting or reconstruction, if the appropriate
experimental protocol is fulfilled. [67,68].
In contrast, a recent theoretical work, combining solvation and electric field effects by
the electric double layer with DFT calculations, reported that Hads, OHads and Oads are
bound less strongly to (110)-step on Pt(332) than to terrace on Pt(111), while this trend
is opposite in vacuum [64]. By employing these new calculations, the theoretical
activity on the stepped Pt(332) surface is now greater than on Pt(111) and Pt(110)
electrodes (empty symbols in Fig. 4A), as found in experiments in acid solutions, and
the theoretical activity’s variation between the different surfaces is much lower.
Nevertheless, the activity trend for Pt(111) and Pt(110) follows the ORR activity in
alkaline and not in acid media. These results illustrate how different simplifications
behind theoretical studies can affect extrapolated conclusions, and highlight the
necessity of refining current theoretical models to get a better understanding of the
molecular processes involved in the ORR mechanism, likely considering surface
charge/potential effects.
Importance of the Oxygen-containing species coverage on the ORR mechanism
Different experimental and theoretical studies have proposed the OHads desorption, or
Oads, reduction as the main bottleneck in the ORR mechanism [55,56,69]. Hence,
because eqns. (6) and (7) are the same as eqns. (11) and (12), both dissociative and
associative mechanisms can occur [45]. A theoretical model postulates that OHads has
site blocking and negative energetic effects, and so the ORR kinetics is determined by
the number of free Pt sites available for O2 adsorption and the adsorption energy of
ORR intermediates [6,28]. This view is supported by some studies in which an
increased ORR activity is accompanied by a weakening of the OHads bond to the
catalyst surface [11,12,13]. However, several works have also reported an enhanced
ORR activity due to an increase in the oxygen coverage [70], and smaller improvements
in the ORR activity, or even inhibition [10], than those expected from the measured
decrease in the OHads coverage [71,72,73,74].
Indeed, OHads reduction is also considered a reaction step in the H2O2 reduction (HPRR)
on Pt [32,75,76,77] but, contrarily to the ORR, this is a mass–controlled reaction up to
~0.95 V [24,32,76]. Therefore, the real role of O-containing species in the ORR
mechanism at Pt surfaces is still unknown. In this sense, as a first step toward the
understanding of it, Figure 6 resumes the integrated charge, after double-layer charging
corrections, from CVs of different Pt single crystals in O2-free solutions, from the PZTC
to different final potentials, 0.78 V (Fig. 6A) and 0.9 V (Fig. 6B). Although the exact
identity of adsorbed species, OHads or Oads, cannot be determined from this graph, this
charge is related to O-containing species coverage on the electrode surface and, because
the 2D-adsorbed structures in O2-free and O2-saturated solutions are comparable [78],
similar O-containing species coverage would be expected during the ORR.
50
-40
-20
0
Pt(100)
Pt(544)
Pt(331)
Pt(221)
Pt(110)
100
0.90 V
Pt(111)
-2
Q / C cm
0.1 M HClO4
0.1 M NaOH
B
150
Pt(100)
50
Pt(111)
Pt(544)
Pt(331)
100
0
200
Angle / grad
0.78 V
Pt(11 1 1)
A
Pt(221)
150
Pt(110)
Q / C cm
-2
200
20
40
60
Angle / grad
Figure 6. Integrated charge, after double-layer charging corrections, from cyclic
voltammograms of different Pt single crystals, associated with the oxygen-containing
species adsorption, as a function of the angle of the surface normal with respect to the
(111) direction, in O2-free 0.1 M HClO4 (■) and 0.1 M NaOH (●) solutions. Integration
is done from PZTC to 0.78 V (A) and 0.90 V (B). Scan rate 50 mV s-1.
From Fig. 6, it is not a clear tendency between ORR activity and O-containing species
coverage (see Fig. 2A). At 0.78 V, for example, although the less active surfaces,
Pt(100) and Pt(11 1 1), have the highest O-containing species coverages, the most active
ones do not have the lowest coverages: Pt(331), Pt(221) and Pt(544) in acid and Pt(111)
in alkaline solutions. Indeed, at 0.9 V, surfaces with the highest coverages apparently
are the most actives, in both acid and alkaline media, with exception of Pt(100).
Therefore, according to these results, the removal of O-containing species is not
necessarily the main problem in the ORR dynamics. In this respect, recent experiments
on PtNi alloys highlight a more important role of the electrode pre-treatment processes
in determining the ORR activity than a reduction on the OHads adsorption rate and
coverage [10]. Moreover, the ORR activity at Pt single crystals in alkaline solutions can
only be explained considering that OHads covered surface is also active for the reaction
[31].
Of course, because of the validity of Sabatier principle [4,5,44] and the linear
relationship between ΔGOHads and ΔGOads [45,55,56,57,58,59], it is expected that a
proper ΔGOHads value may increase ORR activity in some cases, but not always.
However, the ORR mechanism is so complex that a change in ΔGOHads may modify
more than just the OHads reduction kinetics in the whole reaction scheme. Here, it is
important to keep in mind that volcano plots do not usually predict rate, or potential,
determining steps, or the main bottleneck in the reaction scheme. As mentioned, besides
PZTC changes, differences in ORR activity depicted in Fig. 2A should be more related
to the surface structure of water and water dissociation products, OHads and Oads, at Pt
single crystals than to the total coverage of O-containing species, in addition to the
particular identity of these latter species: OHads, Oads or platinum oxides [6,8,20].
For all Pt surfaces, interactions between H2O, OHads, Oads and the surface are important,
as confirmed by changes in the work function and the PZTC of Pt(111) surfaces upon
H2O, OHads and Oads adsorption [65,66,74], and can significantly affect the global ORR
reactivity on these surfaces. In O2-free solutions, the existence and stability of these
species depend on the electrode potential [26,27,79,80], and the same tendency is
expected in O2-saturated solutions [78]. Specifically, three main regions are identified in
the CV of Pt(111), Figure 7. At E < ~0.35 V, hydrogen adsorption, Hads, is responsible
of pseudo-capacitive currents, while the potential region between ~0.6 and 0.8 V has
been attributed to OHads adsorption from water dissociation [25,26,27], in both acid and
alkaline solutions. At higher potentials, Oads formation occurs, depending on the
140
j / A cm
-2
solution pH.
0.1 M HClO4
I
III
II
70
Hads
0
OHads+Oads
OHads
A
-70
0.1 M NaOH
0
III
II
I
50
j / A cm
-2
100
Hads
OHads OHads+Oads
-50
B
-100
0.0
0.2
0.4
0.6
0.8
1.0
1.2
E / V (RHE)
Figure 7. Stable voltammetric profile of a well ordered Pt(111) at 50 mV s-1, in 0.1 M
HClO4 (A) and 0.1 M NaOH (B).
A proposed oxidation dynamics behind the oxide growth process on Pt(111), supported
by both experimental [26,27,80] and theoretical works [65,66,81], considers the
formation of stable H2Oads/OHads and H2Oads/OHads/Oads mixed layers at the second and
third regions depicted in Fig. 7, respectively. Initially, OHads is adsorbed, fast and
reversibly, till the attainment of a stable surface state, identified by the butterfly feature
in Fig. 7. In acidic, but not alkaline, media, the process is followed by a wide potential
region in which no faradaic charge is transferred, Fig. 7A. Increasing the potential
promotes OHads adsorption, and this extra OHads coverage destabilizes the stable adlayer
completed in the butterfly, producing other surface adlayers of varied composition.
Finally, a new, relatively stable, H2Oads/OHads/Oads adlayer is formed, at the end of the
second oxidation peak in Fig. 7 [27,80]. This latter layer evolves rapidly to generate
more stable surface species, at increasing time or potential, presumably PtO oxide
species [27,80]. However, beyond 1.15 V the surface will start to disorder and would
not be more a flat, well ordered, close packed Pt(111) [82,83,84].
In HClO4, a nucleation and growth process, N&G, has been identified in the rising part
of the Oads formation peak, E2p,a, ~1.06 V in Fig. 7A, together with a small reversible
step. This latter step is similar to one observed earlier in flame annealing studies, when
the first potential scan runs in the positive direction from the rest potential [85],
suggesting that chemical reaction steps, involving the so-called thermal oxides, could
also give rise to a significant charge fraction of this peak [80]. A similar decrease in
current in E2p,a has also been reported in O2-saturated solutions and so, dissolved O2
may populate surface sites corresponding to E2p.a, despite of the lack of reduction
current in this region, acting like an extra Oads source and modifying the oxide growth
dynamics [8].
Different H2Oads/OHads and H2Oads/OHads/Oads coverage combinations have been
suggested to illustrate how changes the adlayer stability [65,66]. For H2Oads/OHads
mixed layers, calculated reactions are [65,66]
⅔𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 ⇄ ⅙ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 + ½ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=0.63 V (15)
½ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ⅙ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ⅓ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 + ⅓ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=0.56 V (16)
⅓ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ⅓ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ½ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 + ⅙ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=1.30 V (17)
⅙ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ½ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ⅔ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=1.21 V (18)
Similarly, for H2Oads/OHads/Oads mixed layers, calculated reactions are
½ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ⅙ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ⅙ 𝑀𝐿 𝑂𝑎𝑑𝑠 + ½ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=1.01 V (19)
⅓ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ⅓ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ⅙ 𝑀𝐿 𝑂𝑎𝑑𝑠 + ⅙ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 +⅓ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 − Erev=1.12 V (20)
⅙ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + ½ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 ⇄ ⅙ 𝑀𝐿 𝑂𝑎𝑑𝑠 + ⅓ 𝑀𝐿 𝑂𝐻𝑎𝑑𝑠 +⅙ 𝑀𝐿 𝐻2 𝑂𝑎𝑑𝑠 + 𝐻 + + 𝑒 −
Erev=0.8 V (21)
As can be seen from eqns. (15) to (21), the stability of these adlayers depends on the
electrode potential, which in turns defines the H2Oads, OHads and Oads coverages [65,66].
This is because the formation, or not, of hydrogen bonding structures between product
and reactant [65,66]. The interconversion between these adlayers is evident from the
standard potential value of the different electrochemical equilibria. In this scheme, the
OHads formation begins at ~0.63 V, eqn. (15), and OHads adsorption is more favorable
than Oads adsorption, eqn. (16) vs. (19) and (20). OHads coverages lower than 0.5 V, eqn.
(17), such as those measured in the butterfly step, can be stable below ~1.0 V. At E >
1.0 V, Oads can adsorb, eqn. (20), and the stability of the adlayer becomes particularly
increased at high OHads coverages, eqn. (21). Because of this, it is highly probable that
Oads adsorption will occur first in alkaline than in acid solutions, Figs. 6 and 7, despite
the same ORR activity, Fig. 2 [8,31].
Figure 8 resumes equilibrium potentials for mixed H2Oads/OHads and H2Oads/OHads/Oads
adlayers in the oxidation process described above, represented by eqns. (6) and (7).
Crossed curve in this figure represents the largest of the free–energy differences for
eqns. (6) and (7), assuming a linear correlation between E0 and O-containing species
coverage. If the OHads, or Oads, reduction were the bottleneck in the ORR mechanism,
this latter energetic profile would dictate the kinetic ORR current density, jk. However,
as seen from Fig. 5, when this energetic profile is used to calculate the theoretical ORR
activity (curves 5, with α = 1, and 6, with α = 0.5) reduction currents higher than those
experimentally measured (solid lines) are obtained and thus, the main ORR drawback
should be other than the OHads, or Oads, reduction. Nevertheless, this energetic profile
predicts an interaction between the ORR kinetics and the H2Oads/OHads layer in the
butterfly region, because of the high stability of this layer, as experimentally observed
(see Fig. 1A).
0
E / V (RHE)
1.4
1.2
Eqn. (6)
Eqn. (7)
1.0
0.8
0.6
0.0
0.2
0.4
0.6
0.8
Oxygen-containing
species coverage
Figure 8. Reversible potential of mixed H2Oads/OHads/Oads, E0,6, (eqn. 6) layers and
H2Oads/OHads, E0,7, (eqn. 7) on Pt(111) as function of the O-containing species coverage.
Values are shown considering equilibrium layers with the maximum possible amount of
H2Oads, according to ref. [65,66]. Crossed curve highlights the larger of the free–energy
differences for eqns. (6) and (7) expected during the ORR, assuming a linear correlation
between E0 and O-containing species coverage.
For stepped surfaces, in contrast, ordered H2Oads/OHads or H2Oads/OHads/Oads hexagonal
adlayers are expected to occur only on {111} terraces, and so, increasing the step
density will break the long-range ordering. Figure 9 depicts stable CVs at Pt(554) and
Pt(544) surfaces in 0.1 M HClO4. On {111} terraces, Hads occurs at potentials lower
than 0.35 V, whereas OHads formation takes place between 0.6 and 0.85 V. Regarding
steps, both {111} and {100} monatomic steps have a characteristic voltammetric peak
at 0.13 V and 0.27 V, respectively. Albeit its origin is not entirely clear, it has been
commonly attributed to Hads on step sites [68]. Instead, OHads and Oads adsorption
potentials on the step still are under controversy. Theoretical and experimental results
suggest that while OHads is less stable at steps than at terraces, because of the lack of an
stable H2Oads/OHads network, Oads adsorb on steps at lower potential than on terraces
[64,86].
(111)
150
(554)
s1
j / A cm
-2
100
50
(554)
(544)
(544)
(111)
s2
T
T
s2 s
1
s
s1,2 1
0
-50
s2
-100
-150
0.0 0.2 0.4 0.6 0.8 0.0
E / V (RHE)
0.3
0.6
0.9
1.2
E / V (RHE)
Figure 9. Stable voltammetric profile of a well ordered Pt(111), Pt(554) and Pt(544) at
50 mV s-1, in 0.1 M HClO4. Upper limit potentials, Eup, = 0.9 V (A) and 1.20 V (B).
On {100} steps, it has been suggested that Oads adsorption begins first in the butterfly
region, ~0.78 V, through a diffusive process of OHads from the terrace, and slowly
continues between 0.85 and 1.0 V, with a small peak at 0.92 V [86], Fig. 9. Similarly,
Oads adsorbs on step sites with {110} orientation between 0.85 and 1.0 V, with two
small peaks at 0.92 and 1.01 V. In contrast, current contributions from OHads formation
on steps have not been clearly identified. A recent work suggested possible
contributions of OHads from water dissociation in voltammetric peaks at 0.13 V and 0.27
V [87]. However, this picture cannot explain the nature of the diffusive process at 0.78
V on {100} steps. In the case of {110} steps, a theoretical study suggests that the OHads
formation overlaps with OHads adsorption on the terrace [64], maybe at ~0.82 V.
Comparing CVs with Eup = 0.9 V, Fig. 9A, and Eup = 1.2 V, Fig. 9B, it can be seen that
high Eup affects the surface order of steps, specially with {100} symmetry, which are
converted to {110} surface sites. Because Pt oxides are responsible of surface
disordering, it could be suggested that, in acidic media, PtO oxides are first formed at
steps rather than at terraces, and they are more strongly bound to {100} than to {110}
steps. This would explain why the decrease in the ORR reactivity in the negative-going
scan is higher at stepped surfaces than at Pt(111) when increasing Eup, Fig. 3, and the
lower ORR activity of {100} disordered steps. In consequence, similar to what was
suggested on Pt basal planes in alkaline solutions, the surface PtO oxides coverage, but
not OHads, or Oads, decreases ORR activity at Pt single surfaces in acid solutions, and its
formation depends on the interaction of water, and water dissociation products, with the
surface.
Hydrogen Peroxide Reduction and Oxidation Reactions on the ORR
As mentioned, H2O2 can be oxidized to oxygen (HPOR) or reduced (HPRR) to water,
following two different reactions, eqns. (2) and (3), or it can disproportionate into water
and oxygen, eqn. (4), in a chemical reaction. According to the equilibrium potential for
these reactions, for E > 0.695 V HPOR and HPRR would compete and the overall
current would be defined by the specific electrode kinetics. In the context of the ORR,
H2O2 is one of the most probable reaction intermediate species on various metals and
so, a large HPRR reaction rate would be desirable. However, in most of metals, HPRR
is exceedingly slow irreversible process with a high overpotential. To date, one of the
best electro-catalyst for this reaction is Pt, but even in this electrode the reaction occurs
with a high overpotential, η > 0.7 V [76,88,89].
On (Poly)Pt and Pt single crystals and non-adsorbing acid media, HPRR and HPOR
exhibit a very complex, interrelated dynamics. Both reactions are fast, almost diffusionlimited, and limiting reduction current is immediately followed by a continuous
transition to the limiting oxidation current at increasing potentials [24,32,76,77],
crossing zero around ~0.9 V [24], Figures 10A and 11A. In addition, because HPRR
and HPOR are two different reactions, CVs in peroxide-containing solutions can be
arbitrarily decomposed into the sum of HPOR and HPRR processes, by using
conventional equations for S-shaped electrochemical processes given by
Figure 10. Cyclic voltammograms of different Pt single crystals, in 0.1 M HClO4+2 mM
H2O2 at 2500 rpm, for HPOR and HPRR (solid lines) during positive (a1) and negative
(a2) going scans. Dashed and dotted lines correspond to adjusted separated curves and
the sum of fitted branches, respectively. (b1) and (b2) are fitted E½ values. Arrows
indicate the sweep direction. Scan rate 50 mV s-1. Data were taken from Ref [32].
Figure 11. Cyclic voltammograms of different Pt single crystals, in 0.1 M NaOH+2 mM
H2O2 at 2500 rpm, for HPOR and HPRR (solid lines) during positive (a1) and negative
(a2) going scans. Dashed and dotted lines correspond to adjusted separated curves and
the sum of fitted branches, respectively. (b1) and (b2) are fitted E½ values. Arrows
indicate the sweep direction. Scan rate 50 mV s-1. Data were taken from Ref [33].
𝐸 = 𝐸1/2 + 𝑚log(
𝑗𝑙𝑖𝑚 −𝑗
𝑗
)
(22)
where m is a parameter that would depend on the particular charge transfer mechanism
and E½ is the half-wave potential, i.e. the potential at which the current density is one
half of the corresponding limiting current. Following this approach, CVs for different Pt
single crystals in H2O2-containing solutions have been fitted by eqn. (22), in such a way
that the addition of both HPRR and HPOR contributions should agree as much as
possible with the overall experimental curve, Figs. 10B and 11B [32,33].
In acidic media, the electrode activity for the reaction increases with the step density for
Pt(111) vicinal surfaces, and the activity of basal planes decreases according to Pt(110)
> Pt(100) > Pt(111) for HPRR and Pt(111) > Pt(110) > Pt(100) for HPOR, Fig. 10B1
[32]. In alkaline media, Pt(111) is at the top of HPRR activity and the other basal
planes, Pt(110) > Pt(100), have the lowest activity. Instead, the Pt(111) has the lowest
HPOR reactivity and Pt(100) and Pt(110) have the highest activity, Fig. 11B1 [33].
Contrarily, on oxide covered surfaces all Pt surfaces, except Pt(111) in acid media,
decrease their HPRR activity and increase their HPOR activity, and electrodes with
large (111) terraces are now better electro-catalysts for HPRR but worse for HPOR,
Figs. 10B2 and 11B2.
As it can be seen in Figs. 10B1 and 11B1, E½ for the HPRR at Pt(111) and its vicinal
surfaces follows the ORR activity tendency depicted in Fig. 2A, both in acid and
alkaline solutions. However, E½ is always more positive than the E½ for ORR
[17,20,32,33]. Therefore, if H2O2 is a stable intermediate species, it could not be
detected under reaction conditions, because any H2O2 formed will immediately reduce
to water. In oxide covered surfaces, however, this would not always hold, because the
lower HPRR activity in these surfaces [32], especially in alkaline solutions [33]. This
fact would explain why H2O2 is detected in the negative-going scan during the ORR on
Pt(100) and Pt(110) surfaces in alkaline solutions [17,20].
In contrast, in strong adsorbing electrolyte solutions, kinetic limitations are introduced
in both HPRR and HPOR [77]. Although the global mechanism for these reactions is
unclear [90,91], theoretical results indicate that H2O2 dissociation on Pt(111) is always
possible, regardless of the coverage of spectator species. However, H2O2 adsorption
becomes strongly endothermic, and desorption highly exothermic, at high coverage of
adsorbing anions [77], proving that surface reactivity and the availability of surface sites
are key points in the reaction dynamics [77]. Incidentally, H2O2 has been detected as a
stable intermediate, or final product, during the ORR at Pt surfaces only when the
HPRR is kinetically limited [77], indicating an incomplete electron transfer. It has been
also detected when anions are strongly adsorbed on the electrode surface [6,41,42,43],
specifically with Cl- and Br- [6,41,42,43], or underpotentially deposited Hads adatoms,
Fig. 1A [6,8,17,19]. H2O2 has been also detected under high mass transport conditions
[15,16], and/or in slightly contaminated solutions [2,92,93].
All these results strongly suggest an active role of H2O2 species in the ORR dynamics.
In this scenario, there would be another serious drawback in the ORR dynamics, in
addition to a slow first electron transfer step. This is because any effort for decreasing
the present ORR overpotential, beyond the potential region where the HPRR occurs,
would imply not only an increase in the reaction rate of the first electron transfer step
but also of the HPRR, and this latter reaction could be the real problem in the ORR
mechanism at high potentials.
Toward an understanding of ORR mechanism: Synergistic Theoretical and
experimental approaches
Undoubtedly, ORR is a catalyzed reaction and so, the electrode surface has a pivotal
role on its dynamics. Therefore, it is expected that at least one of the reaction steps
involves adsorbed species. However, it does not necessarily imply that all electron
transfer steps must correspond to inner-sphere electrode reactions, as proposed by
current accepted theoretical mechanisms, eqns. (5) to (12). It is possible than one of the
ORR elemental steps may involve outer-sphere electrode reactions with a dissolved
species directly participating in the reaction elemental step. Hence, the ORR mechanism
could be composed by a mix of inner and outer-sphere electron transfer steps.
Indeed, the superoxide anion, O2*-, has been suggested as the primary radical produced
after the first electron transfer during the ORR on Pt surfaces in alkaline solutions [94].
However, in acidic environments the picture is not clear yet, although early works
already suggested the possible formation of aqueous hydroxyl radicals, OH*, during the
ORR on Pt surfaces, through a similar mechanism proposed for the Haber–Weiss
reaction [88,89,95]. In agreement, a recent experimental study has suggested the
production of aqueous OH* during the ORR on (Poly)Pt, but not on gold, surfaces [7].
Similarly, from an electrochemical study on Pt(111) in acid solutions, the reduction of
an aqueous intermediate species has been proposed as RDS. In this latter case the
hydroperoxyl radical, OOH*, was suggested [8], according to
𝑂2 + 𝑃𝑡𝐻2 𝑂 = 𝑃𝑡𝑂𝐻 + 𝑂𝑂𝐻 ∗
(23)
𝑂𝑂𝐻 ∗ + 𝐻 + + 𝑒 − = 𝑃𝑡𝐻2 𝑂2
(24)
Unfortunately, available experimental techniques have not been able to undoubtedly
identify the identity of ORR intermediate species [6,9]. The reasons behind this fact can
be diverse. It could be, for example, that the physical, or chemical, properties of ORR
intermediates cannot be measured with current experimental techniques or because the
equipment sensitivity is not enough to detect them. Hence, it is necessary the
development of new experimental techniques, together with better and more realistic
theoretical models, to reach a fully understanding of the ORR mechanism. In this sense,
it is clear that full agreement should exist between experiments and theoretical
calculations for model surfaces describing the same processes, as a first step to
understand the electro-catalysis at more complex surfaces, such as dispersed nanoparticles employed on real applications.
Concluding remarks
In this perspective, the oxygen reduction dynamics at Pt single surfaces is reviewed, and
a special emphasis is given to the role of the electrode surface charge, the oxygen
containing species coverage, the surface structure of water, and water dissociation
products, and the nature of adsorbed species, such as H2Oads, OHads, Oads and PtO
oxides. Additionally, the hydrogen peroxide oxidation and reduction dynamics on Pt
single crystals have also described and compared to the ORR kinetics.
It is shown that, the main problem in the ORR mechanism is not uniquely a problem of
OHads, or Oads, coverage, or competence for active sites. Instead, other factors, such as
the electrode surface charge and the nature of adsorbed species in the electrode may
significantly modify the electrode activity. Moreover, H2O2 can be an intermediate
species in the ORR mechanism, even if it is not measured under some experimental
conditions, and thus, the main bottleneck in the ORR at high potentials could be the
H2O2 oxidation.
Acknowledgements
Notes and references
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