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Oscillatory Electro-oxidation of Methanol on Nanoarchitectured Pt pc /Rh/Pt Metallic Multilayer Ernesto Chaves Pereira ACS Catal., Just Accepted Manuscript • DOI: 10.1021/cs501652u • Publication Date (Web): 29 Dec 2014 Downloaded from http://pubs.acs.org on January 2, 2015
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Oscillatory Electro-oxidation of Methanol on Nanoarchitectured Pt pc /Rh/Pt Metallic Multilayer
Journal: Manuscript ID: Manuscript Type: Date Submitted by the Author: Complete List of Authors:
ACS Catalysis cs-2014-01652u.R1 Article 22-Dec-2014 Nagao, Raphael; IQSC/USP, Physical Chemistry Freitas, Renato; UFSCar, Chemistry; University of Mato Grosso, Chemistry Silva, Camila; UFSCar, Chemistry Varela, Hamilton; Instituto de Quimica de Sao Carlos, USP, Physical Chemistry; Fritz Haber Institute of the Max Planck Society, Physical Chemistry Pereira, Ernesto; Universidade Federal de S�o Carlos, Chemistry
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Oscillatory Electro-oxidation of Methanol on Nanoarchitectured Ptpc/Rh/Pt Metallic Multilayer
9 10 11 12 13
Raphael Nagao,1,2,† Renato Freitas,1,3 Camila Silva,1 Hamilton Varela,2,4 Ernesto Chaves
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Pereira1,* 1
15
Department of Chemistry, Federal University of São Carlos, POBox 676, 13565-905, São Carlos, SP, Brazil.
16 2
17
Institute of Chemistry of São Carlos, University of São Paulo, POBox 780, 13560-970, São Carlos, SP, Brazil.
18 3
19
Department of Chemistry, Federal Univeristy of Mato Grosso, POBox 78060-900, Cuiaba, MT, Brazil.
20 4
21
Ertl Center for Electrochemistry and Ctalysis, GIST,
Cheomdan-gwagiro 261, Buk-gu, Gwangju 500-712, South Korea
22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41
*corresponding authors:
[email protected] (ECP). †
current address: Department of Chemistry, Saint Louis University, 3501 Laclede Ave., St. Louis, MO 63103, USA.
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Abstract The oscillatory electro-oxidation of methanol was studied on polycrystalline Ptpc and Ptpc/Rh2.0/Pt1.0 metallic multilayers. The surfaces investigated consisted of 1.0 Pt outlayer surface deposited onto 2.0 Rh intralayers beneath a Pt outlayer and over the polycrystalline Ptpc substrate. Besides experimental studies, numerical simulations were performed using a dimensionless kinetic model for the electro-oxidation of methanol in order to provide a better understanding of the role played by the nanostructured metallic multilayer electrode in the electrocatalytic activity. A comparable electrochemical behavior found for the cyclic voltammetries in blank acidic media was observed in both electrodes. Remarkably, an increase of 90% in the peak current density around 0.88 V vs. RHE in the electro-oxidation of methanol appeared when Ptpc/Rh2.0/Pt1.0 was utilized. The numerical simulations suggested that this increase in the electrocatalytical activity for the metallic multilayers electrode is due to the prevention of carbon monoxide adsorption on the surface, and a consequent increase in the production of carbon dioxide from the direct pathway. Indeed, a decrease in the reaction rate constant of carbon monoxide formation resulted in an increase of the current density associated to CO2 formation in the potenciodynamic sweep, besides the decrease in amplitude and frequency of the oscillatory time-series. As the rate of carbon monoxide adsorption is suppressed by the presence of the metallic multilayers, the intrinsic drift usually found in the oscillatory electro-oxidation of methanol was enhanced and oscillations ceased earlier. Overall, the combination of electrochemical experiments and numerical simulations suggests that carbon monoxide acts as a poisoning species instead of a reaction intermediate in the electro-oxidation of methanol. keywords: oscillations, metallic multilayer, methanol, electrocatalysis.
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1. Introduction
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The search for more selective and efficient catalysts for complete electro-
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oxidation of methanol has been an active field of research in past decades due the
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possibility of use fuel cells as a green and renewable energy source.1,2 Several
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methodologies were employed in order to enhance the catalytic activity by
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modifications in their structural properties, and among them, the use of alloys,
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electrodeposits and nanoparticles have shown to be quite promising in this
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direction.3,4,5,6 Metallic multilayers are a new class of advanced engineered materials
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that are known to exhibited high structural stability, mechanical strength, high ductility,
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toughness and resistance to fracture and fatigue.7 Moreover, it has been widely explored
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mainly due its magnetic properties.8,9 Under the perspective of the electrocatalysis,
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Pereira and co-workers10,11,12,13,14,15,16 have extensively studied the preparation and
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characterization of metallic multilayers deposited onto polycrystalline platinum (Ptpc) as
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catalysts for the electro-oxidation of small organic molecules.
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For instance, the authors observed that, although Ptpc/Metal/Pt (Metal = Rh, Ru,
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Bi and Ir) electrode has a Pt outlayer surface, the electrocatalytic activity for such
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metallic multilayers is higher than that observed for Ptpc electrode. Oliveira et al.10
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experimentally observed increases in the peak current density between about two and
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four times for the electro-oxidation of methanol, ethanol, formaldehyde and formic acid
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on Ptpc/Rh/Pt with respect to that on Ptpc. Interestingly, a high selective oxidation of
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ethanol to carbon dioxide was obtained when the reaction was carried out on Ptpc/Ir/Pt
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nanostructured electrode,14 as reflected by the absence of the infrared bands relative to
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acetic acid and acetaldehyde molecules.
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In despite of much effort devoted to this topic, the role played by the metallic
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multilayer electrodes in the improvement of the electrocatalytic activity is not
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completely understood. Herein we provide some mechanistic evidences based on a
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combination of experiments, modeling and numerical simulations, under both
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conventional and oscillatory regimes, that the increase in the current density observed in
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the electro-oxidation of methanol over the metallic multilayer electrode is mainly
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caused by the production of carbon dioxide from the oxidative decomposition of the
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active intermediate, i.e. direct pathway. This argument is supported by the highly
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selective electronic effect caused by nanostructured metallic multilayer in preventing
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the adsorption of carbon monoxide. Potential oscillations registered under galvanostatic
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control were utilized as a mechanistic probe, corroborating the scenario of suppression
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of adsorbed carbon monoxide.
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2. Experimental Procedure
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Solutions were prepared using ultrapure water (Millipore, 18.2 MΩ cm), HClO4
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(Sigma-Aldrich, 70%), H2SO4 (Mallinckrodt, 99.8%), RhCl3.3H2O (Alpha-Aesar,
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99.9%), H2PtCl6.6H2O (Sigma-Aldrich, 99.9%) and H3COH (J. T. Baker, 99.9%).The
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electrochemical cell, all glassware and electrodes were cleaned in a solution of KMnO4
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and NaOH for 12 h. Subsequently, all electrochemical apparatus were left soaked for 1
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h in a bath of H2O2 and H2SO4 concentrated solution. Finally, the system were carefully
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and exhaustively rinsed with ultrapure water and boiled in it for at least three times.
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This cleaning procedure resulted in voltammograms free from any impurities.
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In order to prepare the metallic multilayer electrode, a polycrystalline Ptpc (0.3
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cm2, Sigma-Aldrich, 99.99%) substrate was utilized. It was mechanically polished with
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diamond paste down to 1.0 µm, washed with acetone and a large amount of ultrapure
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water. The potential of deposition was 0.05 V (vs. a reversible hydrogen electrode,
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RHE) for 1800 and 120 s leading to about 2.0 and 1.0 Rh and Pt monolayers (ML),
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respectively. The procedure was as follows: Rh intra-layer was electrodeposited from
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the precursor solution 2.0 10-5 mol L-1 of RhCl3.3H2O in 0.1 mol L-1 of HClO4 and the
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charge related to electrodeposition process was used to obtain the number of Rh ML.
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After that, the electrode was washed kindly in ultrapure water and transferred to another
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electrochemical cell with 5.0 10-4 mol L-1 of H2PtCl6.6H2O solution in 0.1 mol L-1 of
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HClO4 where Pt outlayer could be electrodeposited leading to Ptpc/Rh2.0/Pt1.0
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nanostructured electrodes.
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The electrochemical characterization of the metallic multilayer electrodes was
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carried out using a potentiostat/galvanostat Autolab PGSTAT 302N equipped with
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SCANGEN module. Voltammetric curves were measured in 0.5 mol L-1 of H2SO4
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solution in the potential range between 0.05 and 1.40 V. A platinized-Pt plate (Sigma-
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Aldrich, 99.99%) was used as auxiliary electrode. Prior to the experiments, the solutions
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were sparged with N2 (White Martins, 5N) for 30 min. The electrolyte was kept free of
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atmospheric pollutant-gas by continuously introducing N2 into the region of the
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electrochemical cell above the electrolyte. Prior each galvanostatic time-series, the
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working electrode was cycled between 0.05 to 1.40 V in 0.5 mol L-1 of H3COH and 0.5
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mol L-1 of H2SO4 along 30 cycles. Soon after, the galvanostatic time-series for several
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current values were recorded.
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3. Modeling the electro-oxidation of methanol
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Electrochemical experiments were compared to numerical simulations in order
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to shed light in some mechanistic aspects in the electro-oxidation of methanol on the
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modified surfaces. The modeling procedure was based on ten elementary steps, r1 to
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r10, in a dimensionless model previously proposed by Nagao et al.17 As widely
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accepted, the electro-oxidation of methanol proceeds through the so-called dual
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pathway mechanism.18 The indirect pathway is described by the adsorption of the
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methanol molecule with consecutive dehydration steps which results in the formation of
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adsorbed carbon monoxide,19,20 i.e. COad. See steps r1 and r2, respectively,
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H3 COH + 3Pt → HCOad + 3H+ + 3e-
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HCOad → COad + H+ + e- + 2Pt
k2
(r1) (r2)
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In our model we considered only the formation of OHad as the oxygenated
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species by the oxidation of water. In this case, step r3 represents the oxidation reaction
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while r4 is related to the reduction one. This simplification has been used successfully
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in different modeling procedures for the electro-oxidation of small organic
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molecules.17,21,22 In addition, COad reacts with adsorbed oxygenated species in a
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Langmuir-Hinshelwood mechanism, step r5, resulting in CO2.23
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H2 O + Pt → OHad + H+ + e-
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OHad + H+ + e- → H2 O + Pt (r4)
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COad + OHad → CO2 + H+ + e- + 2Pt (r5)
(r3)
k4
k5
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Occurring in parallel, the direct pathway takes place by the formation of soluble
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products, such as formic acid and formaldehyde which can also form CO2 from the
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oxidative decomposition of the active intermediates,24,25,26 and methylformate that is
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easily detected by means of differential electrochemical mass spectrometry
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(DEMS).27,28,29 The nature of the active intermediates is still under debate, but it is
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currently believed to proceed in a very fast path.30 In this context, we adopted the
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production of adsorbed formate, i.e. HCOOad in step r7 as the active intermediate 5 ACS Paragon Plus Environment
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formed only from formic acid dehydration, and ignored the formation of HCOOCH3 by
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the homogeneous reaction between HCOOH and H3COH in solution27 and the
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esterification reaction between HCOad and H3COH.29 Nevertheless, methylformate can
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be followed by HCOad time-evolution as it is a precursor to its formation.17
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HCOad + Pt + H2 O → HCOOHad + H+ + e- + 2Pt
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HCOOHad → HCOOad + H+ + e-
k7
k8
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HCOOad + H+ + e- → HCOOHad
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HCOOad + Pt → CO2 + H+ + e- + 3Pt
k9
(r6) (r7) (r8) (r9)
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We attributed the steps r6 to r9, as also the responsible for the production of
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CO2 derived of re-adsorption process of soluble species. In this case, only formic acid is
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incorporated in the model and all molecules form adsorbed formate. An excellent
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agreement between experiments and simulations based on these assumptions was
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recently observed.31 Finally, formic acid might also result in COad from a reduction step
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of HCOOad, as proposed by Cuesta et al.32
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k10
HCOOad + Pt + H+ + e- COad + H2 O + 2Pt (r10)
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Each coverage variable was converted in a generic representation xn according
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to: x2 = θHCO , x3 = θOH , x4 = θCO , x5 = θHCOOH , and x6 = θHCOO ; xe = eωϕ as the electrical
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component for faradaic reactions, and ω as the transfer coefficient. The surface free
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sites are denoted as x1 and can be written as,
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x1 = 1 - 3x2 - x3 - x4 - 2x5 - 2x6 207 208
applying the kinetic laws in each elementary step, it is obtained,
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v1 = k1 x21 xe (e1)
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v2 = k2 x2 xe (e2)
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v3 = k3 x1 xe (e3)
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v4 = k4 x3 xe (e4)
-1
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v5 = k5 x3 x4 xe (e5)
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v6 = k6 x1 x2 xe (e6)
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v7 = k7 x5 xe (e7)
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v8 = k8 x6 xe (e8)
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v9 = k9 x1 x6 xe (e9)
-1
-1
v10 = k10 x1 x6 xe
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(e10)
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Additional simplifications as the assumption of two free active sites33 in
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Equation e1 and the independency of the number of electrons produced in xe were
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adopted, and as previously observed, do not result in significant changes compared to
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experimental observations.17,34 Hence, the electrochemical model can be written as a set
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of six autonomous ordinary differential equations, e11 to e16,
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x 2 = v1 - v2 - v6 (e11)
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x 3 = v3 - v4 - v5 (e12)
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x 4 = v2 - v5 + v10 (e13)
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x 5 = v6 - v7 + v8 (e14)
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x 6 = v7 - v8 - v9 - v10 (e15)
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ϕ = i - vF (e16)
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with,
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vF = 3v1 + v2 + v3 - v4 + v5 + v6 + v7 - v8 + v9 - v10 (e17) i =
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u - ϕ r
(e18)
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In Equation e18, i might be interpreted as the total current, u as the applied
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potential and r as the ohmic drop between the working and reference electrodes. The six
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ordinary differential equations were numerically integrated using the tool ode15s in the
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Matlab software. When not specified, the initial conditions were considered as zero.
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4. Results and Discussion
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Figure 1 depicts the electrochemical characterization carried out by linear
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potentiodynamic sweeps at dU/dt = 0.10 V s-1 on a polycrystalline platinum electrode,
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Ptpc as the blue curves, and metallic multilayer nanostructured electrode, Ptpc/Rh2.0/Pt1.0
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as the red ones; experiments were performed in sulfuric acid 0.1 mol L-1 in the (a)
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absence and (b) presence of methanol 0.5 mol L-1. Numerical simulations presented in
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(c) were performed at du/dt = 0.01 with changes in k2 and k3 in order to qualitatively
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reproduce the experimental observations. In this case, the set of reaction rate constants,
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k2 = 5.0 and k3 = 1.0 represents the blue curves, while k2 = 0.7 and k3 = 0.9 the red ones.
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Figure 1. (a) cyclic voltammetry at dU/dt = 0.10 V s-1 under Ptpc (blue curves) and Ptpc/Rh2.0/Pt1.0 (red curves) in sulfuric acidic media; (b) potential linear sweep at same rate measured in electrochemical experiments and (c) numerically calculated in the electro-oxidation of methanol. k2 = 5.0 and k3 = 1.0 (blue curve); k2 = 0.7 and k3 = 0.9 (red curve) in numerical simulations at du/dt = 0.01, u(0) = -0.3 and x4(0) = 0.9. Additional parameters are: k1 = 6.0, k4 = 4.0, k5 = 0.079, k6 = 50, k7 = 600, k8 = 30, k9 = 300, k10 = 0.1, ω = 15, r = 0.05.
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A typical voltammetric behavior can be observed in Figure 1(a) for Ptpc and
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Ptpc/Rh2.0/Pt1.0 electrodes in sulfuric acid solution,35 where the electrochemical process 9 ACS Paragon Plus Environment
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on the cyclic voltammograms can be associated to: i) hydrogen adsorption and
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desorption between 0.05 to 0.40 V, ii) double layer region between 0.40 to 0.80 V, iii)
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formation and reduction of platinum oxides, between 0.80 to 1.40 V and 1.4 to 0.40 V,
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respectively. As evident in Figure 1(a), the voltammetric profile for polycrystalline
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platinum is thoroughly recovered when Ptpc/Rh2.0/Pt1.0 structure is formed by the
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deposition of a Pt layer over the Rh surface, which implies in a surface composed
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mostly by platinum atoms. XPS measurements have confirmed this experimental
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evidence.10,12 The electrochemical surface areas for Ptpc and Ptpc/Rh2.0/Pt1.0 electrodes
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were calculated using the procedure well established in the literature36 which considers
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a charge density of 210 µC cm-2 as being equivalent to desorption of one hydrogen
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monolayer. The electrochemical surface area for Ptpc and Ptpc/Rh2.0/Pt1.0 electrodes
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differ about 5% and then all current densities were normalized by its values. These
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results are in agreement with in situ AFM experiments which revealed the same
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roughness mean square values for both electrodes.11,12,13
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Analyzing the electro-oxidation of methanol over Ptpc and Ptpc/Rh2.0/Pt1.0
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electrodes in Figures 1(b), an increase of c.a. 90% was observed in the peak current
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density for Ptpc/Rh2.0/Pt1.0 along the linear potentiodynamic sweep. Similar behavior of
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the increased catalytic activity in the electro-oxidation of different small organic
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molecules have been showed by Pereira and co-workers10,11,12,13,14,15,16. As commonly
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accepted, the electro-oxidation of methanol proceeds through parallel pathways,18 as
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already presented in Section 3. One controversial point to be stressed is the role played
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by carbon monoxide during the electro-oxidation process. Due to its strong influence in
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fuel cells efficiency, it has triggered a continuous study of its removal in low potentials
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by the presence of different metals such as Ru, Rh, Sn, etc.37,38,39 In the literature, a
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great effort has been done in the development of more selective catalysts based on the
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bi-functional mechanism and electronic effects.40,41,42,43
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Considering the absence of Rh ad atoms in the surface of metallic multiyear
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electrode, justified by the no observation of the characteristic bands of the second metal
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in XPS measurements10,12 and the similarity between the profiles for the hydrogen
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underpotential deposition in the cyclic voltammetries in Figure 1(a), the mechanism that
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would explain the experimental results in Figure 1(b) might be attributed to: (a) the
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electronic interaction of the exposed Pt outlayer and the Rh interlayer beneath,
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weakening the back-donation between the CO anti-ligand orbitals 2π*, with energies
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above the Fermi level, and d band from the Pt outlayer;42,43 (b) the compressive strain 10 ACS Paragon Plus Environment
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due to the electrodeposition of a monolayer between two distinct substrates and
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interatomic distances, resulting in an enlargement of d band, and consequently, lower
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energies of adsorption.44,45 Both electronic effects can, indeed, prevent the formation of
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COad on the surface. Experimental results have showed the catalytic effect of Rh for the
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electro-oxidation of methanol,10,38,46 but as previously discussed, in this set of results we
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can discard the contribution from the bi-functional mechanism and the formation of
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rhodium oxides.
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Numerical simulations were also carried out and a good correspondence with the
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electrochemical measurements was obtained when the reaction rate constants of
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formation of adsorbed COad and OHad were changed, as it can be seen in Figure 1(c).
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Although a suppression for carbon monoxide layer formation is enhanced in the
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presence of metallic multilayer electrode, interestingly, just a slight shift in the current
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density peak for more negative potentials has been observed. According with our
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simulations, this behavior is mostly attributed to the decrease of k2 values which results
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in the increase of catalytic activity while k3 affected drastically the potential range
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where the maximum current occurs (see Figure S1 in Supporting Information for
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additional details). Hence, the enhancement of the catalytic activity in the electro-
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oxidation of methanol might be understood in terms of the removal of carbon monoxide
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from the surface.
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Galvanostatic experiments were carried out in different applied currents and
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concomitantly compared with numerical simulations. In Figure 2, potential time-series
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measured at j = 0.32 mA cm-2 on (a) Ptpc (blue curves) and (b) Ptpc/Rh2.0/Pt1.0 (red
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curves) working electrodes were compared with numerical simulations at i = 0.20 with
324
(c) k2 = 5.0, k3 = 1.0, k6 = 50 (blue curves) and (d) k2 = 0.5, k3 = 0.8, k6 = 23 (red
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curves).
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Figure 2. Experimental (item (a) and (b)) and calculated (item (c) and (d)) potential time-series in the electro-oxidation of methanol at j = 0.32 mA cm-2 and i = 0.20, respectively. The blue curves represent Ptpc and k2 = 5.0, k3 = 1.0, k6 = 50; while red ones represent Ptpc/Rh2.0/Pt1.0 and k2 = 0.5, k3 = 0.8, k6 = 23; working electrodes in electrochemical experiments and numerical simulations. Additional parameters are: k1 = 6.0, k4 = 4.0, k5 = 0.079, k7 = 600, k8 = 30, k9 = 300, k10 = 0.1, ω = 15.
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Electrochemical oscillations presented an usual profile described in the
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literature.47,48 As observed in Figure 2, a good adjustment between experiments and
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simulations was obtained when the formation of COad and OHad were disfavored. It
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reinforces the evidence that the metallic multilayer electrode might prevent more
340
selectively the formation of carbon monoxide on the surface. Additionally, qualitative
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agreement was even improved when k6 values were decreased from 50 to 23. In spite of
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a significant increase of the current density in voltammetric experiments, the
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electrochemical variables do not reveal explicitly the presence of the second metal,
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nevertheless some hints become clearly evident in oscillatory regime. As already
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pointed out, a rich variety of dynamic behavior found during the electrochemical
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constants.49 In this specific case, the amplitude shift for lower potential values followed
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by the decrease of the period is a good indicative of a surface modification by a second
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metal.50,51
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The experimental and numerical results support the proposition that the
351
oxidation of adsorbed carbon monoxide by oxygenated species in a Langmuir-
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Hinshelwood mechanism is not the predominant reaction step for the CO2 formation,
353
and consequently, the faradaic current. In other words, our results suggest that adsorbed
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carbon monoxide acts as a poisoning species instead of a reaction intermediate in the
355
electro-oxidation of methanol in a wide potential window. This hypothesis has already
356
been investigated52,53,54 and it is in agreement with our observations. This controversial
357
behavior played by carbon monoxide have been discussed with the aid of spectroscopic
358
and spectrometric techniques coupled in the electrochemical cell.52,53,54,55,56,57,58,59 The
359
classification as a reaction intermediate was attributed to the high sensitiveness of CO2
360
production when different anions were present in the electrolyte. As the formation of
361
adsorbed carbon monoxide by the surface dehydration of methanol requires a larger
362
number of active sites, CO oxidation was considered a non-negligible elementary step
363
in the overall carbon dioxide production,56,57,58 e.g. above 0.6 V vs. RHE,
364
approximately. In contrast, recent results have showed a strong impact of soluble
365
intermediates in the total production of CO2 and the faradaic current.31,60,61 In fact,
366
Nagao et al.34 have demonstrated that sulfate anions in electrolyte impede more
367
effectively the CO2 production from the direct than the indirect pathway during the
368
oscillatory electro-oxidation of methanol.
369
Based on this considerations, we utilized modeling and numerical simulations to
370
have some insights of the relative weight of production of CO2 from the direct and
371
indirect pathways. Figure 3 shows the production of CO2 from the elementary steps r5
372
and r9 in the (a) linear sweeps described in Figure 1(c); and (b) potential time-series in
373
Figure 2(c) and 2(d). Table 1 depicts a more quantitative relationship of the Figure 3 in
374
terms of the maximum current found in the linear potential sweep, and the mean
375
production during the oscillations.
376
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377 378 379 380 381 382 383 384
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Figure 3. (a) simulated linear sweeps: k2 = 5.0 and k3 = 1.0 (blue curves); k2 = 0.7and k3 = 0.9 (red curves), both at k6 = 50, du/dt = 0.01, u(0) = 0.3 and x4(0) = 0.9; (b) potential time-series with k2 = 5.0, k3 = 1.0 and k6 = 50 (blue curves); and k2 = 0.5, k3 = 0.8 and k6 = 23 (red curves), both at i = 0.20. Additional parameters are: k1 = 6.0, k4 = 4.0, k5 = 0.079, k7 = 600, k8 = 30, k9 = 300, k10 = 0.1, ω = 15, r = 0.05.
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385
Under potentiodynamic control, vide Figure 3(a1) to (a3), the decrease of the
386
reaction rates k2 and k3 resulted in a substantial increase of the formation of CO2,
387
represented by v5 + v9, from 0.230 to 0.335, which is in agreement with the Figures 1(b)
388
and 1(c). However, the enhancement in the electro-catalytic activity is clearly related to
389
the higher contribution from the direct pathway. It is mainly occasioned by the
390
consumption of COad, releasing free sites to the oxidative decomposition of the active
391
intermediate. Otherwise, in oscillatory conditions, Figure 3 (b1) to (b3), the decrease of
392
k2 and k3 also resulted in an inversion for the global contribution of CO2 formation:
393
decrease of v5 from 0.020 to 0.007 and a consequent increase of v9 from 0.012 to 0.032,
394
respectively. Nevertheless, it does not exhibit dramatic changes in the mean production,
395
see Table 1 for a more detailed description. Differently from observed in
396
electrochemical experiments coupled with a mass spectrometer,17,31,34 these oscillations
397
do not end up with CO2 peak splitting. Harmonic oscillations, favored in lower applied
398
currents, tend to generate period-1 carbon dioxide time-series, instead of relaxation ones
399
where period-3 and 4 have been observed.17,31,34
400 401
Table 1. Relative production of carbon dioxide extracted from Figure 3. v5 blue maximum current in linear potential sweep mean production during oscillations
v9 red
blue
v 5 + v9 red
blue
red
0.190 0.151 0.042 0.240 0.230 0.335 0.020 0.007 0.012 0.025 0.032 0.032
402 403
The removal of carbon monoxide from the surface seems to be the key
404
parameter to increase the activity. Clearly, the metallic multilayer is very efficient in
405
this purpose. This idea was also tested using the electrochemical oscillations in long
406
term dynamics as a probe to infer about the reaction mechanism. A general overview of
407
the effect of Ptpc/Rh2.0/Pt1.0 working electrode in the oscillatory electro-oxidation of
408
methanol is better visualized in Figure 4.
409 410 411
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Figure 4. Period (a) and amplitude (b) of the galvanostatic oscillations experimentally measured as a function of number of cycles in the electro-oxidation of methanol over Ptpc (blue curves) and Ptpc/Rh2.0/Pt1.0 (red curves). The dashed arrow indicates the direction of the applied current density from the lowest to the highest values: j = 0.32, 0.48, 0.64 and 0.81 mA cm-2.
421
The oscillatory period and amplitude were accounted as function of the number
422
of cycles over Ptpc and Ptpc/Rh2.0/Pt1.0 surfaces and under different applied current
423
densities, e.g. j = 0.32, 0.48, 0.64 and 0.81 mA cm-2. Overall, the period and amplitude
424
were not significantly influenced neither by the variation in applied current and the use
425
of the metallic multilayer electrode. As observed in Figure 4(a) and (b), the decrease of j
426
resulted in slower oscillations with smaller amplitudes, while the use of Ptpc/Rh2.0/Pt1.0
427
electrode enhanced this behavior in all range of current densities studied. In contrast to
428
previous experimental findings,62 the relation between period and amplitude in the
429
oscillatory electro-oxidation of methanol does not seem to be straightforward. This 16 ACS Paragon Plus Environment
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430
aspect is very important and will be discussed later with the aid of numerical
431
simulations.
432
Another interesting effect of the nanostructured metallic multilayer electrode in
433
the oscillations is the temporal stability of time-series. Perini et al.51 have demonstrated
434
the highly efficient way to attenuate the intrinsic drift during potential oscillations in the
435
electro-oxidation of formic acid using Sn as a second metal on the surface. The authors
436
were able to keep a robust time-series for a very long period of time, e.g. 40 h. Based on
437
mechanistic considerations, the rationalized method was performed aiming the
438
prevention of OsubPt (sub means subsurface oxygen35) formation, in its turn, leads Sn as
439
the main responsible to form oxides that might be used in the oxidation with
440
carbonaceous species. The efficiency of the Ptpc/Rh2.0/Pt1.0 electrode towards the
441
stabilization of time-series in the electro-oxidation of methanol was analyzed as well.
442
Table 2 summarizes some experimental data collected in Figure 4 as the measured drift
443
in terms of ∆Um.∆cycles-1 and the total number of cycles. Both were reproduced by
444
numerical simulations. In this set of results the drift was simulated by a linear increase
445
of the applied current from initial i of 0.05, rate of 2.00 10-3 and 2.24 10-3. This
446
characteristic has been proved experimentally.48,63
447
Table 2. Effect of the metallic multilayer Ptpc/Rh2.0/Pt1.0 and reaction rate constant of COad formation on the total number of oscillatory cycles. Simulations consider: k1 = 6.0, k3 = 1.0, k4 = 4.0, k5 = 0.079, k6 = 50, k7 = 600, k8 = 30, k9 = 300, k10 = 0.1, ω = 15.
experiments
448 449 450
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electrode
drift
ratio in drift
Ptpc
0.130 mV cycle-1
Ptpc/Rh2.0/Pt1.0
0.117 mV cycle-1
Ptpc
0.190 mV cycle-1
# of cycles 669
1.11
Ptpc/Rh2.0/Pt1.0
0.169 mV cycle
k2 = 5.0
2.24 10-3
-1
1.64 406 502
1.12
2.16 232 132
1.12 k2 = 1.0
2.00 10
-3
ratio in # of cycles
2.16 61
451 452
As observed in all experimental results with constant applied current, the drift
453
evolves faster in the Ptpc related to Ptpc/Rh2.0/Pt1.0 electrode (see Table S1 in Supporting
454
Information for additional values for ∆Um.∆cycles-1 and the total number of cycles). 17 ACS Paragon Plus Environment
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455
However, instead of compare the drift at same applied current, where the available area
456
of the electrode plays a pivotal role, we conducted the analysis normalizing the drift in
457
terms of ∆Um.∆cycles-1. With a ratio of approximately 1.1 in the drift, the number of
458
cycles was roughly twice higher in the presence of Ptpc. Remarkably, the same behavior
459
was obtained when the reaction rate of COad formation was increased from 1.0 to 5.0.
460
This is a clear evidence that the Ptpc/Rh2.0/Pt1.0 electrode does not have a stabilizing
461
effect as observed by Perini et al.51 even with the presence of a second metal. The major
462
difference by their experiments and ours is that in our case the second metal is located
463
beneath the platinum outlayer in the Ptpc/Rh2.0/Pt1.0 electrode. As the bi-functional
464
mechanism can be discarded in this electrode configuration, the surface oxidation still
465
occurs on platinum active sites, but in this particular situation, the formation of carbon
466
monoxide is much slower. Considering that COad is the main specie responsible to
467
consume platinum oxides on the surface, formed in high potential values, it is expected
468
a lesser number of cycles when the metallic multilayer is used as working electrode.
469
The intrinsic drift might also induce morphological changes in period and
470
amplitude simultaneously.48,63 Figure 5 depicts an amplitude vs. period diagram for
471
experiments (filled circles in bold axes) and numerical simulations (open circles in italic
472
axes). The blue curves represents the Ptpc electrode and k2 = 5.0; while the red ones is
473
associated to Ptpc/Rh2.0/Pt1.0 electrode and k2 = 1.0, for electrochemical experiments and
474
numerical simulations, respectively. The surface drift was implemented by slow
475
changes in the applied current with initial i of 0.05 and a linear sweep of 0.002.
476 477 478
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479 480 481 482 483 484 485 486 487 488 489 490 491 492
Figures 5(a) and (b) show a typical characteristic of quadratic relationship
493
between the oscillatory amplitude and bifurcation parameter, along the period drifting,
494
i.e. applied current. As it is changing slowly during the time, we can state the
495
proportional relation: amplitude ∝ applied current which is predicted theoretically.64
496
However, differently from item (b), Figure 5(a) shows a distinct behavior in the onset of
497
the time-series. At first, there is an inverse tendency of increase between the period and
Figure 5. Amplitude vs. Period diagram represented as filled circles for the experiments (amplitude / V vs. period / s in bold axes) at j = 0.64 mA cm-2 and open circles for numerical simulations (a vs. p in italic axes). (a) Electrochemical reaction on Ptpc and k2 = 5.0 for numerical simulations (blue curves); (b) Ptpc/Rh2.0/Pt1.0 and k2 = 1.0 (red curves). The black arrows indicate the direction of increasing the number of cycles from initial i of 0.05 and drift of 0.002. Both simulations consider: k1 = 6.0, k3 = 1.0, k4 = 4.0, k5 = 0.079, k6 = 50, k7 = 600, k8 = 30, k9 = 300, k10 = 0.1, ω = 15.
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Page 20 of 26
498
amplitude, followed by concomitant increment after a critical point. In Figure 4, the
499
increase of the applied current resulted in faster oscillations with larger amplitudes. As
500
verified experimentally,62 larger amplitudes are usually connected with larger periods. A
501
suitable explanation for this dynamic behavior is that higher potential values must be
502
reached in order to promote an effective surface reactivation when the trajectories of the
503
limit cycle have taken a slow excursion in low potentials. Experimental and numerical
504
results65,66 converge to the idea that poisoning species might block the surface, slowing
505
down the reaction rates. Considering the electro-oxidation of methanol, the
506
consumption of carbon monoxide by oxygenated species is the elementary reaction
507
which mostly influence the oscillatory period.65 Based on this statement, we can clearly
508
make a distinction between Figure 5(a) and (b) by the amount of adsorbed carbon
509
monoxide formed on the surface. The former initially speed-up the formation of
510
carbonaceous intermediates which can be latter used in the oxidation of oxygenated
511
species, prolonging the time-series, and consequently, the overall number of cycles.
512
Nevertheless, Figure 5(b) presents a monotonic and gradual decrease by the slow
513
platinum oxidation (see Figure S2 in Supporting Information for the time evolution of
514
x4). Similar behavior has been reported in the literature.48,63
515
In summary, the experimental and numerical results described here allow us to
516
propose that a metallic multilayer is indeed a new electrocatalytic material. In the
517
present case, it has improved properties for the electro-oxidation of methanol in acidic
518
solution, leading important changes in the mechanistic reaction compared to bulk
519
polycrystalline platinum.
520 521
5. Conclusions
522
Metallic multilayer electrodes have been utilized as a new kind of
523
nanostructured materials with the aim of increasing the catalytic activity in the electro-
524
oxidation of small organic molecules. In this work we provided additional mechanistic
525
evidences of the role played by Ptpc/Rh2.0/Pt1.0 surface in the electro-oxidation of
526
methanol by a combination of electrochemical experiments, modeling and numerical
527
simulations. A similar voltammetric profile in acidic solutions was observed in both
528
electrodes, i.e. Ptpc and Ptpc/Rh2.0/Pt1.0. Remarkably, an increase of about 90% in the
529
peak current density around 0.88 V vs. RHE has appeared in an anodic linear
530
potentiodynamic sweeps for the Ptpc/Rh2.0/Pt1.0 electrode, with respect to that for Ptpc.
531
Based on spectroscopic measurements previously reported10,12 and similar profiles for 20 ACS Paragon Plus Environment
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532
the hydrogen underpotential deposition in the cyclic voltammetries, the bi-functional
533
mechanism and the formation of rhodium oxides on the surface can be discarded as
534
responsible for increasing the catalytic activity. Consequently, it might be attributed
535
solely to the highly selective electronic effect in the prevention of adsorbed carbon
536
monoxide.
537
Additionally, a considerable increase of period and decrease of amplitude of the
538
oscillations were observed in a wide range of applied current densities, and in contrast
539
to previous studies, the presence of a second metal enhanced the intrinsic drift
540
embedded in the time-series. A good qualitative agreement between experiments and
541
simulations was obtained only when lower values of reaction rate constant of carbon
542
monoxide adsorption were considered in the numerical model. These evidences
543
highlight the effect of metallic multilayer electrodes on the prevention of carbon
544
monoxide layer formation. In other words, carbon monoxide seems to act as a poisoning
545
specie instead of a reaction intermediate that contribute to the faradaic current and total
546
CO2 production.
547 548
6. Associated Content
549
Supporting Information
550
Simulated voltammetric profiles for i and CO2 besides the time-series for COad and CO2
551
in the electro-oxidation of methanol. Number of cycles and drift in terms of
552
∆Um.∆cycles-1 measured experimentally under different j. This material is available free
553
of charge via the Internet at http://pubs.acs.org.
554 555
7. Acknowledgments
556
The authors would like to thank the Brazilian research funding agencies: National
557
Council of Scientific and Technological Development (CNPq, grants: 229171/2013-3
558
and 304458/2013-9) and São Paulo Research Foundation (FAPESP, grants:
559
2010/05555-2, #2013/16930-7 and 2013/16930-7) for financial support.
560 561
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