Evidence for Decoupled Electron and Proton Transfer in the

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Letter pubs.acs.org/JPCL

Evidence for Decoupled Electron and Proton Transfer in the Electrochemical Oxidation of Ammonia on Pt(100) Ioannis Katsounaros,*,†,‡,§ Ting Chen,§,∥ Andrew A. Gewirth,† Nenad M. Markovic,‡ and Marc T. M. Koper*,§ †

University of Illinois at Urbana−Champaign, Department of Chemistry, 600 South Mathews Avenue, Urbana, Illinois 61801, United States ‡ Argonne National Laboratory, Materials Science Division, 9700 South Cass Avenue, Lemont, Illinois 60439, United States § Leiden University, Leiden Institute of Chemistry, Einsteinweg 55, P.O. Box 9502, 2300RA Leiden, The Netherlands ∥ Shandong Jianzhu University, School of Science, 250101 Jinan, P. R. China S Supporting Information *

ABSTRACT: The two traditional mechanisms of the electrochemical ammonia oxidation consider only concerted proton−electron transfer elementary steps and thus they predict that the rate−potential relationship is independent of the pH on the pH-corrected RHE potential scale. In this letter we show that this is not the case: the increase of the solution pH shifts the onset of the NH3-to-N2 oxidation on Pt(100) to lower potentials and also leads to higher surface concentration of formed NOad before the latter is oxidized to nitrite. Therefore, we present a new mechanism for the ammonia oxidation that incorporates a deprotonation step occurring prior to the electron transfer. The deprotonation step yields a negatively charged surface-adsorbed species that is discharged in a subsequent electron transfer step before the N−N bond formation. The negatively charged species is thus a precursor for the formation of N2 and NO. The new mechanism should be a future guide for computational studies aiming at the identification of intermediates and corresponding activation barriers for the elementary steps. Ammonia oxidation is a new example of a bond-forming reaction on (100) terraces that involves decoupled proton−electron transfer.

T

formation.10−13 Density functional theory (DFT) calculations suggest that the square arrangement of atoms on the Pt(100) surface provide the active ensemble site for N−N bond-making during ammonia oxidation.14 This assessment is done in the awareness that a thermally annealed Pt(100) crystal still contains a large density of ad-islands of monatomic height on top of the flat terraces,15 and such ad-islands may in fact possess high(er) activity provided they preserve the square symmetry of the active site.14 Therefore, the AOR is an example of a reaction in which the presence of (100) sites is essential to make or break bonds, other examples including the NO2− reduction to N2 on Pt(100), the CO reduction to C2H4 on Cu(100), the CH3OCH3 oxidation to CO2 on Pt(100) and the O2 reduction to H2O on Au(100).14,16−19 Many of the aforementioned reactions exhibit remarkable pH dependence that originates from competitive steps in the reaction mechanism. For example, CO reduction to C2H4 on Cu(100) electrodes is favored at high pH values, which was attributed to a decoupled proton−electron transfer in the C2H4-forming pathway.20 By decoupled proton−electron

he electrochemical ammonia oxidation reaction (AOR) may find applications in the fields of energy conversion (“ammonia electrolysis” for hydrogen production), environmental protection (treatment of waste streams), analytical chemistry (ammonia sensors) or electrosynthesis (hydroxylamine production).1−4 Platinum is one of the best catalysts for this reaction as it can convert ammonia to dinitrogen at low potentials.5−7 Two mechanisms have been proposed to explain the formation of N2 during the AOR and they are summarized in Table 1.8,9 In the first mechanism (Oswin−Salomon) the N2 formation takes place by the dimerization of two Nad atoms, whereas in the second mechanism (Gerischer−Mauerer) the N−N bond forms by the recombination of two NHx,ad groups (where x = 1 or 2) and strongly adsorbed Nad is a poison for the N2 formation. Every electron transfer step in either mechanism is accompanied by a simultaneous proton (or rather hydroxide) transfer, thus both mechanisms predict that the reaction rate at a given potential is independent of the pH, in the pH-corrected reversible hydrogen electrode (RHE) scale. The ability of platinum to oxidize ammonia to nitrogen relies exclusively on the presence of (100) terraces. The reaction is so structure−sensitive that Pt(111) and Pt(110) are practically inactive, while the introduction of monatomic (111) or (110) steps to Pt(100) is detrimental to the N−N bond © 2016 American Chemical Society

Received: November 15, 2015 Accepted: January 12, 2016 Published: January 12, 2016 387

DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392

Letter

The Journal of Physical Chemistry Letters Table 1. Proposed AOR mechanisms mechanism 1 (Oswin−Salomon)8 −

−

mechanism 2 (Gerischer−Mauerer)9 (1)

NH3,ad + OH(adδ−) ⇌ NH 2,ad + H 2O + δ e−

(5)

NH 2,ad + OH− ⇌ NHad + H 2O + e−

(2)

NH 2,ad + OH(adδ−) ⇌ NHad + H 2O + δ e−

(6)

NHad + OH− ⇌ Nad + H 2O + e−

(3)

NHx ,ad + NHy ,ad ⇌ N2Hx + y ,ad(x , y = 1 or 2)

(7)

Nad + Nad ⇌ N2

(4)

N2Hx + y ,ad + (x + y)OH(adδ−) ⇌ N2 + (x + y)H 2O + (x + y)δ e−

(8)

NHad + OH(adδ−) ⇌ Nad + H 2O + δ e−(Nad : poison)

(9)

NH3,ad + OH ⇌ NH 2,ad + H 2O + e

transfer, we mean that the transfer of the electron and the proton/hydroxide does not take place in the same elementary step (such as the steps in Table 1) but rather one-by-one in subsequent steps. Although the rate-potential relationship for concerted proton−electron transfer shifts by 60 mV per pH unit on the standard hydrogen electrode (SHE) scale, protondecoupled electron transfer always manifests with a shift different from 60 mV, the exact value depending on the sequence of the proton and electron transfer steps (for details and a mathematical model explaining this pH dependence see ref 21−23). Conversely, a shift different from 60 mV on the SHE scale does not have necessarily to be due to a decoupled proton−electron transfer. In this communication, we study the effect of the solution pH on the electrochemical oxidation of ammonia on Pt(100). We focus only on solutions with a pH higher than the pKa value of the NH4+/NH3 equilibrium (= 9.27)24 because ammonia is present in the form of NH4+ at lower pH. We will demonstrate the importance of a decoupled proton−electron transfer in the AOR mechanism, a crucial feature that is currently not incorporated in any of the existing mechanisms. The voltammetry of a thermally annealed Pt(100) electrode in argon-saturated solutions of different pH in the region 11.8− 13.5 is shown in Figure 1a. The solutions of different pH were prepared by dissolving the appropriate amount of KOH in 0.1 M HClO4. The exact voltammetric profile for Pt(100) in alkaline solutions depends strongly on the cooling conditions following annealing, and thus, it varies among studies; however, there is qualitative agreement that four peaks appear in the positive-going scan.25,26 The peaks are located (Figure 1a) at around +0.28 VRHE (Peak I), + 0.39 VRHE (Peak II), + 0.45 VRHE (Peak III), and +0.58 VRHE (Peak IV), and they have been attributed to hydrogen desorption and hydroxide adsorption at various types of sites (110-type defects, short- and long-range 100 terraces, etc.).25,26 It is difficult to assign each peak to a certain surface process, because the adsorption/desorption of hydrogen and hydroxide are in fact overlapping. The voltammograms for all solutions practically coincide irrespective of the solution pH, provided that the potential is expressed versus the pH-corrected RHE scale. Figure 1b shows that the position of all four peaks (with the potential expressed versus the SHE) shifts negatively by around 60 mV when the activity of hydroxide in solution is increased by 1 order of magnitude. In addition, the area associated with each peak is practically the same (Figure 1a). Assuming that different adsorption/desorption characteristics for hydrogen and hydroxide manifest by differences in the blank voltammetry, Figure 1 suggests that the surface coverage with Hupd and OHad does not change significantly at a given potential versus the RHE scale (under these transient conditions) for solutions in this pH range. This conclusion is limited to such a relatively narrow pH

Figure 1. (a) Cyclic voltammograms for Ar-saturated solutions of different pH in the region 11.8−13.5, with the potential expressed versus the RHE. Scan rate: 50 mV s−1. (b) Potential (versus the SHE) of the peaks I−IV in the positive-going scan as a function of the solution pH. The slopes corresponding to all four peaks are between 60 and 64 mV pH−1.

region only; the adsorption/desorption processes are likely to be different in much less alkaline solutions or even acidic media (see for example ref 25), where also the nature of the electrolyte ions is, however, different. In the assessment of the curves in Figure 1a, one needs to keep in mind that small differences may originate from the fact that the concentration of KOH is not the same in each solution, and thus, the concentration of trace impurities present even in the highestpurity KOH is different. Figure 2a shows the positive direction of the hydrodynamic voltammograms recorded with rotating the Pt(100) electrode at 400 rpm in solutions of different pH in the region 10.7 to 13.8, which additionally contain 1 × 10−3 M NH4ClO4. We chose to record the voltammograms under hydrodynamic 388

DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392

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The Journal of Physical Chemistry Letters r Had ⇌ r * + r H+ + ne−

(10)

r * + NH3 ⇌ NH3,ad

(11)

r Had + NH3 ⇌ NH3,ad + r H+ + ne−

(12)

where the asterisk represents a vacant Pt site, r is number of hydrogen atoms that need to be desorbed for each ammonia molecule adsorbed, and n is the number of electrons exchanged per r protons. The Nernst equation for the equilibrium (12) is E = E0 −

a Hr ada NH3 2.303RT log nF a NH3,ada Hr +

(13)

where E is the equilibrium potential, E0 the standard equilibrium potential, R the universal gas constant (8.314 J mol−1 K−1), T the temperature (in K), F the Faraday constant (96 485 A s mol−1), and ai the activity of the species i. If we assume that the peak potential shows the same concentration dependence as the equilibrium potential, which would be rigorously correct for a reversible couple, eq 13 predicts that the peak potential shifts by −60 r/n mV (versus the SHE scale) for every unit increase of pH, at a given concentration of ammonia. Figure 2b shows that the peak potential shifts negatively by ca. 60 mV per pH unit increase, or in other words, the peak potential is the same for all solutions studied if the potential is expressed versus the RHE. Thus, an equal number of protons and electrons is exchanged in reaction 10 (r = n). This originates from the fact that the peak centered at +0.38 VRHE is controlled by hydrogen adsorption, and the latter is not dependent on the pH (on the RHE scale) as Figure 1 suggests. Note also that the peak current density and the associated charge are the same irrespective of the solution pH. In addition, eq 13 predicts that the peak potential shifts by −60/n mV (at a given pH) for an increase of the ammonia concentration by 1 order of magnitude. Given the fact that a 10-fold increase in the ammonia concentration leads to a negative shift of the prepeak by 30 mV (see Supporting Information, Figure S1), we propose that two hydrogen atoms need to be desorbed for each ammonia molecule adsorbed at the prepeak (n = r = 2). At higher potentials, the current increases until it reaches a maximum of ca. + 1.3 mA cm−2 at around +0.69 VRHE (for the pH = 13 solution, black curve) and then decreases (see Figure 2a). In this region, ammonia is oxidized after being transported from the bulk to the surface. Given the fact that these measurements were performed with a rotating electrode, the appearance of a current peak and the consequent decrease of the current cannot be attributed to limitations in the mass transport of ammonia but to a deactivation process. It should be noted that the amount of KOH dissolved in HClO4 was found to influence the peak current (but not the onset or the peak potential) in solutions of the same pH, which is probably an effect of trace impurities in KOH. Therefore, the small differences in the peak currents for the solutions with pH > 11.8 (Figure 2a) should not be interpreted as a true pH effect but as an unavoidable artifact related to the preparation methodology. Additional evidence is that practically no differences in the peak current are observed if the same experiments are performed in stagnant electrolytes, where the transport of trace impurities is slow. Contrary to the pH independence of the prepeak on the RHE scale, the onset potential for ammonia oxidation, as well as the potential of the corresponding peak current shift to lower

Figure 2. (a) Positive-going scans (50 mV s−1) in Ar-saturated KOH solutions of different pHs that additionally contain 1 × 10−3 M NH4ClO4. Rotation rate: 400 rpm. The voltammograms represent the first cycle, as this corresponds to the freshly annealed, well-ordered (100) electrode. (b) Potential versus the SHE of the prepeak (black points) and of the AOR onset (defined as the potential when the current density is 0.2 mA cm−2, red points), as a function of the solution pH.27

conditions because changes in the local concentration of reactants and products at the solid−liquid interface are less pronounced compared to stagnant solutions and they can be modeled as the mass transport regime is well defined. The presence of ammonia in solution results to the appearance of a peak at around +0.38 VRHE (Figure 2a). The exact origin of this peak has not been fully resolved even though it has been previously observed in alkaline ammonia solutions on Pt(100).10−13,28 We assign this peak to adsorption of ammonia that is possible only after hydrogen desorption, given that hydrogen is known to be adsorbed on Pt(100) in the ammonia-free KOH solution at potentials lower than +0.3 VRHE,29 and the currents in the solutions with and without ammonia coincide at potentials lower than of the prepeak. The fact that ammonia adsorption cannot take place unless hydrogen desorbs above ca. + 0.3 VRHE, implies that the Pt(100)−H bond is stronger than the Pt(100)−NH3 bond, in agreement with DFT calculations of the corresponding adsorption energies.30 We believe that ammonia is not dehydrogenated upon adsorption, because the area under the peak corresponds only to the charge of hydrogen desorption from Pt(100) (ca. 200 μC cm−2); if ammonia was adsorbed in the form of NH2, an additional charge should have been observed. Therefore, the following scheme can describe the processes taking place at the prepeak region: 389

DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392

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The Journal of Physical Chemistry Letters

versus the SHE at a given current density by 59 mV per one pH unit. The above explanation does not rule out the participation of adsorbed hydroxide in the oxidation of ammonia. Instead, the fact that the NH3 oxidation is initiated in a region where adsorbed hydroxide is present at the surface might imply that the first step of NH3 oxidation possibly involves OHad. However, one elementary step in the reaction mechanism has to involve OH− rather than OHad, either at or before the ratedetermining step, to explain the observed pH effect. In particular, in a mechanism explaining the results shown in Figure 2, either reaction 14 must be considered irreversible and rate-determining, or reaction 14 is reversible and in preequilibrium, with reaction 15 being irreversible and ratedetermining. At more positive potentials, another oxidation peak appears at around +1.0 VRHE, for the solutions of pH higher than ca. 12.5. The position of this peak coincides with the region at which NO oxidation to nitrite is taking place on Pt(100) in alkaline solutions;33 therefore, we assign it to the oxidation of adsorbed nitric oxide that is formed as a product of the AOR. It should be noted that NO formation is not considered in the two dominant AOR mechanisms (Table 1) and it needs to be incorporated in the ammonia oxidation mechanism, especially since it may play an essential role in the mechanism as recently suggested also by Guay and co-workers.28 The comparison of the area under the NO oxidation peak at +1.0 VRHE in solutions of different pH reveals that it becomes more prominent by increasing the pH, whereas it vanishes in solutions less alkaline than ca. pH 12.5. Rodes et al. have shown that neither the peak potential (versus the RHE) nor the oxidative charge for adsorbed NO oxidation depend on the pH, in the entire pH region.33 Thus, the pH effect on the NO oxidation peak in ammonia solutions is not attributed to an effect on the rate of NO oxidation, but on the rate of NO formation by the AOR. This indicates that the pathway for NO formation also involves a deprotonation step before the electron transfer (namely that NO is formed after reactions 14 and 15). The nature of the (NxHy−1)ad− species is uncertain but we believe that we can still outline some likely candidates based on the following reasoning. Possible values of x and y could be x = 1 and y = 1−3, or x = 2 and y = 1−5. As described above, the pathways for both N 2 and NO formation involve a deprotonation step preceding the electron transfer. It does not seem likely that there are two such steps in the AOR mechanism, and thus, we assume that the N2 and the NO formation pathways share the same intermediates, at least until the formation of the negatively charged species (reaction 14) and the subsequent electron transfer step (reaction 15). If this assumption is correct, the proton-decoupled electron transfer step precedes the N−N bond formation, and thus the most likely negatively charged species are NH2,ad−, NHad− or Nad−. In that scheme, the formation of (possibly poisoning) adsorbates that follow the proton-decoupled electron transfer, for example, Nad, is also pH-dependent. In summary, we have shown that the ammonia oxidation to nitrogen and to nitric oxide on Pt(100) in alkaline solutions involves a rate-determining deprotonation step without simultaneous electron transfer. This implies that the pathways that lead to N2 and NO formation include the intermediate formation of a negatively charged surface-adsorbed species, which has not been incorporated in any of the two proposed mechanisms for the AOR, as they both consider only concerted proton−electron transfer elementary steps.8,9 Therefore, we

values by increasing the pH, on the RHE scale. If the potential is expressed versus the SHE, a non-Nernstian slope of around 90−100 mV pH−1 is observed for the onset potential, defined as the potential at which the measured current density is 0.2 mA cm−2 (red data points in Figure 2b). We do not aim to interpret the meaning of this slope quantitatively, but the fact that the onset and the peak potentials clearly do not follow a 60 mV pH−1 dependence (or the fact that the voltammograms do not coincide on the RHE scale) is a striking observation that implies a complex impact of the pH on the reaction. It should be noted that the confidence in the determination of the onset potential was within 2 mV, thus insignificant compared to the differences observed between different solutions. This observation could be the result of a change of the local pH at the solid−liquid interface, due to the consumption of hydroxide ions during the AOR. If we follow the methodology presented in ref 31, we can calculate based on Fick’s law that at the rotation rate of our experiments, the measured current densities (proportional to the rate of hydroxide consumption) are sufficient to alter significantly the local pH only when the solution pH is lower than ca. 11.2 (see Supporting Information for details). This is, for example, the case for the measurement in pH 10.7 (cyan curve) where a plateau is observed at ca. 0.65 mA cm−2, and it is related to a limitation in the diffusion of OH−: the mass transport calculations suggest that the local pH at this current density has decreased to ca. 9.7 (see Supporting Information). It must be noted, though, that our local pH calculations are based on the assumption that the surface is homogeneously active for ammonia oxidation and, thus, that the rate of hydroxide consumption is the same at any electrode location. Moreover, Figure 1 shows that no significant differences in the potential-dependent OHad coverage can be evidenced in the pH range studied. Therefore, the observed pH effect cannot be explained by a higher availability of OHad species, which might promote the ammonia oxidation in a bifunctional mechanism. In addition, the impact of the pH on the potential of zero total charge (pztc) for Pt(100) (a 5 mV positive shift on the RHE scale when the pH increases by one unit)32 is practically insignificant compared to the observed impact of the pH on the onset of ammonia oxidation. Finally, the enhancement of the AOR by increasing the pH cannot be attributed to the higher conductivity of more alkaline solutions, because we compensated for the electrolyte resistance (see Experimental Methods section). Given the fact that none of the explanations above is compatible with the experimental data, we suggest that the observed enhancement of the AOR in the more alkaline solutions can be explained if one considers a step in the reaction sequence in which the proton (hydroxide) transfer precedes the electron transfer, for example NxHy ,ad + OH− ⇌ (NxHy − 1)−ad + H 2O

(14)

This deprotonation step leads to a short-lived negatively charged species, which is then discharged by electron transfer step, for example (NxHy − 1)−ad ⇌ NxHy − 1,ad + e−

(15)

Such a decoupling of the electron and proton transfer is not included in the existing AOR mechanisms (Table 1); both of them consider only concerted proton−electron transfer at every elementary step, which would lead to a shift of the potential 390

DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392

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The Journal of Physical Chemistry Letters propose that a revised mechanism for the ammonia oxidation reaction is required that includes the findings of this study. Providing a more complete mechanism of the ammonia oxidation incorporating such a decoupled proton−electron transfer step is our aim in future works. In addition, computational studies that focus on possible intermediates and corresponding activation barriers of the elementary AOR steps,14,34−37 should also consider such a charged adsorbate species as a precursor for N2 and NO. Finally, the ammonia oxidation on Pt(100) is a new example of a bond-making reaction that takes place exclusively on (100) terraces and that involves decoupling of the proton and the electron transfer in the reaction mechanism.14,22

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was supported by a Marie Curie International Outgoing Fellowship within the seventh European Community Framework Programme to I.K. under Award IOF-327650, and by the U.S. Department of Energy, Office of Science, Materials Sciences and Engineering Division (contract DE-AC0206CH11357). T.C. acknowledges support from the Chinese Scholarship Council (award number 201306220113). The experimental work presented in this manuscript was performed at the Argonne National Laboratory.

EXPERIMENTAL METHODS A three-electrode Teflon FEP cell was used with a platinum wire as the counter electrode and a saturated Ag/AgCl (BASi, RE-6) as the reference electrode separated from the main cell compartment with a salt bridge. The potential is expressed with respect to either the reversible or the standard hydrogen electrode potential. The current in the voltammograms was normalized to the geometric area of the electrode (0.2826 cm2). The potential was controlled with an Autolab PGSTAT 302N potentiostat and rotation was controlled with a Pine AFMSRCE electrode rotator. A rotation rate of 400 rpm was chosen for the hydrodynamic voltammetry as a compromise between controlled mass transport conditions and a slow transport of trace impurities from the solution to the surface. Positive feedback was used to compensate for the electrolyte resistance and the remaining resistance was no more than 5 Ω in any experiment. All electrochemical measurements were performed in an argon-saturated electrolyte and at room temperature. Prior to each measurement, the Pt(100) single-crystal (Princeton Scientific Corp., 6 mm diameter) was annealed using induction heating at 1100 °C for 10 min in a hydrogen/ argon flow (3%/97%), followed by cooling down to room temperature for 10 min in the same gas environment. The crystal was then immediately covered with a drop of ultrapure water, assembled into a rotating disc electrode (RDE) configuration, and transferred to the electrochemical cell. The annealed electrode was immersed in the electrolyte at +0.05 (±0.01) VRHE, and a sweep was recorded starting from this potential to the positive direction. All voltammograms shown in this manuscript represent the first cycle, as this corresponds to the freshly annealed, well-ordered (100) electrode. The solutions were freshly prepared using ultrapure water (Millipore, 18.2 MΩ, TOC < 4 ppb) and chemicals of the highest purity from Sigma-Aldrich. The addition of ammonium perchlorate to the KOH solution was always done just before recording the voltammograms, to minimize the concentration drop that occurs with time in alkaline solutions due to the NH4+/NH3 equilibrium and subsequent NH3 evaporation. The concentration of ammonia used was not enough to alter the bulk pH of the solution in any of the experiments. Solutions of different pH were prepared by dissolving the appropriate amount of KOH in 0.1 M HClO4. The gases used were 6N quality purchased from Airgas Inc.

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Dependence of the prepeak versus the concentration of ammonia in solution, and details on the calculation of the local pH at the solid−liquid interface during the ammonia oxidation reaction. (PDF)

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REFERENCES

(1) Rosca, V.; Duca, M.; de Groot, M. T.; Koper, M. T. M. Nitrogen Cycle Electrocatalysis. Chem. Rev. 2009, 109, 2209−2244. (2) Vitse, F.; Cooper, M.; Botte, G. G. On the Use of Ammonia Electrolysis for Hydrogen Production. J. Power Sources 2005, 142, 18− 26. (3) Marinčić, L.; Leitz, F. B. Electro-Oxidation of Ammonia in Waste Water. J. Appl. Electrochem. 1978, 8, 333−345. (4) López de Mishima, B. A.; Lescano, D.; Molina Holgado, T.; Mishima, H. T. Electrochemical Oxidation of Ammonia in Alkaline Solutions: Its Application to an Amperometric Sensor. Electrochim. Acta 1998, 43, 395−404. (5) Wasmus, S.; Vasini, E. J.; Krausa, M.; Mishima, H. T.; Vielstich, W. DEMS-Cyclic Voltammetry Investigation of the Electrochemistry of Nitrogen Compounds in 0.5 M Potassium Hydroxide. Electrochim. Acta 1994, 39, 23−31. (6) Gootzen, J. F. E.; Wonders, A. H.; Visscher, W.; van Santen, R. A.; van Veen, J. A. R. A DEMS and Cyclic Voltammetry Study of NH3 Oxidation on Platinized Platinum. Electrochim. Acta 1998, 43, 1851− 1861. (7) de Vooys, A. C. A.; Koper, M. T. M.; van Santen, R. A.; van Veen, J. A. R. The Role of Adsorbates in the Electrochemical Oxidation of Ammonia on Noble and Transition Metal Electrodes. J. Electroanal. Chem. 2001, 506, 127−137. (8) Oswin, H. G.; Salomon, M. The Anodic Oxidation of Ammonia at Platinum Black Electrodes in Aqueous KOH Electrolyte. Can. J. Chem. 1963, 41, 1686−1694. (9) Gerischer, H.; Mauerer, A. Untersuchungen Zur anodischen Oxidation von Ammoniak an Platin-Elektroden. J. Electroanal. Chem. Interfacial Electrochem. 1970, 25, 421−433. (10) Vidal-Iglesias, F. J.; García-Aráez, N.; Montiel, V.; Feliu, J. M.; Aldaz, A. Selective Electrocatalysis of Ammonia Oxidation on Pt(100) Sites in Alkaline Medium. Electrochem. Commun. 2003, 5, 22−26. (11) Vidal-Iglesias, F. J.; Solla-Gullón, J.; Montiel, V.; Feliu, J. M.; Aldaz, A. Ammonia Selective Oxidation on Pt(100) Sites in an Alkaline Medium. J. Phys. Chem. B 2005, 109, 12914−12919. (12) Vidal-Iglesias, F. J.; Solla-Gullón, J.; Feliu, J. M.; Baltruschat, H.; Aldaz, A. DEMS Study of Ammonia Oxidation on Platinum Basal Planes. J. Electroanal. Chem. 2006, 588, 331−338. (13) Rosca, V.; Koper, M. T. M. Electrocatalytic Oxidation of Ammonia on Pt(111) and Pt(100) Surfaces. Phys. Chem. Chem. Phys. 2006, 8, 2513−2524.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpclett.5b02556. 391

DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392

Letter

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DOI: 10.1021/acs.jpclett.5b02556 J. Phys. Chem. Lett. 2016, 7, 387−392