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. 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