Electrocatalytic Efficiency of the Oxidation of Small Organic Molecules

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The Electrocatalytic Efficiency of the Oxidation of Small Organic Molecules under Oscillatory Regime Marcelo V. F. Delmonde, Loriz F Sallum, Nickson Perini, Ernesto Rafael Gonzalez, Robert Schloegl, and Hamilton Varela J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.6b06692 • Publication Date (Web): 06 Sep 2016 Downloaded from http://pubs.acs.org on September 7, 2016

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The Electrocatalytic Efficiency of the Oxidation of Small Organic Molecules under Oscillatory Regime M. V. F. Delmonde,1 L. F. Sallum,1 N. Perini,1,2 E. R. Gonzalez,1 R. Schlögl,2,3 H. Varela,1,* 1

Institute of Chemistry of São Carlos, University of São Paulo POBox 780, 13560-970, São Carlos, SP, Brazil. 2 Max Planck Institute for Chemical Energy Conversion Stiftstrasse 34–36, 45470 Mülheim an der Ruhr, Germany. 3 Fritz Haber Institute of the Max Planck Society, Department of Physical Chemistry, Faradayweg 4-6, D-14195 Berlin, Germany.

Abstract The electrocatalytic oxidation of small organic molecules is of general importance for energy related issues such as the fuel cells and electrochemical reform. The common emergence of current/potential oscillations in these reactions has implications on mechanistic aspects as well as on the overall conversion, and thus on the performance of practical devices. We investigate in this paper some general features of the electro-oxidation of formaldehyde, formic acid, methanol, and ethanol on platinum and in acidic media, with emphasis on the comparison of the activity under conventional and oscillatory regimes. The comparison is carried out by different means and generalized by the use of identical experimental conditions in all cases. In all four systems studied, the occurrence of potential oscillations is associated with excursions of the electrode potentials to lower values, which noticeably decreases the overpotential of the anodic reaction, when compared to that in the absence of oscillations. Quantitatively speaking, a two-fold enhancement in the power density was observed in an idealized fuel cell operated with formaldehyde. This aspect, together with spontaneous selfcleaning processes, presents important advantages to the use of autonomous oscillations to reach both higher and long-term activities. Finally, some mechanistic aspects of the studied reactions are also discussed.

* corresponding author: HV ([email protected], ++551633738059.)

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Introduction

The interconversion between chemical and electrical energies plays an undisputable role in the challenges of energy supply for a sustainable future.1 In this context, to master the chemistry of hydrogen and hydrogen-carriers is certainly a major target. In this respect, examples of relevant hydrogen-carriers include ammonia, borohydride, formaldehyde, formic acid, methanol, ethanol, glycerol, dimethyl ether, etc. Aside from their structural simplicity, some mechanistic aspects of the electro-oxidation of such molecules are presently unknown and still pose significant fundamental questions. In particular, the electro-oxidation of small organic molecules assumes a prominent role in the so-called Direct Liquid Fuel Cells (DLFCs), which consists mostly of polymer electrolyte membrane (PEM) systems that use liquid fuels without preliminary reform; and also in the electrochemical reform, where the fuel is oxidized instead of water at considerably lower potentials.25-6 Table 1 compares the thermodynamic data for the complete oxidation of hydrogen, formaldehyde, formic acid, methanol and ethanol. ∆Go is the free energy and n the number of electrons involved in the overall reaction. Uocell and Eoanode represent the thermodynamic open circuit voltage for the complete cell, and the theoretical standard potential for the anodic reaction, respectively.

Table 1: Thermodynamic data for the oxidation of H2, formaldehyde, formic acid, methanol, and ethanol at 25 oC and 1 atm.3,7-9 Theoretical cell reaction + 1⁄2 → () CH2 O + O → CO2 + H2 O HCOOH + 1⁄2 O → CO2 + H2 O

CH3 OH + 3⁄2 O → CO2 + 2H2 O CH3 CH2 OH + 3O → 2CO2 + 3H2 O

∆Go (kJ mol-1) -237.2 -521.7 -285.3

n 2 4 2

Uocell (V) 1.23 1.35 1.48

Eoanode (V) 0.00 - 0.12 - 0.25

-697.9 -1325

6 12

1.21 1.14

0.02 0.09

As seen in Table 1, the standard potential for the anodic reaction, Eoanode, for all organic species shown are rather similar to that of hydrogen. Nevertheless, in contrast to hydrogen, the oxidation of such small organic molecules proceeds only with considerable reaction rates far above the standard potential and thus high overpotentials results. This limitation is indeed the main cause of the restricted use of these species as fuels in DLFCs. As far as the reaction mechanism is concerned, it is irrefutable that the frequent occurrence of


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parallel pathways and the formation of byproducts, i.e. dissolved partially oxidized species, severely impact the desirable complete oxidation to carbon dioxide. There are several reports on the kinetic instabilities during the electro-oxidation of hydrogen and several small organic molecules on platinum,10-41 and platinum-based catalysts.42-47 Particularly for the electro-oxidation of hydrogen, there are reports of considerable performance enhancement in operational PEMFCs (polymer electrolyte membrane fuel cells) fed with H2 contaminated with tiny amounts (~ 100 ppm) of carbon monoxide.48-62 It has been shown that under oscillatory conditions the adsorbed carbon monoxide is periodically and autonomously oxidized. As a result, the overall performance of the PEMFC was significantly enhanced. The possibility of improving the total electrooxidation of small organic molecules to CO2 under oscillatory regime is therefore desirable, specially if one considers the similarities in the dynamics of electro-oxidation of H2/CO mixtures and of small organic molecules.23,63 Despite the growing interest in the use of kinetic instabilities to investigate reaction mechanisms36,38,40,64,65 and the enhanced performance of PEMFC fed with H2/CO mixtures under oscillatory regime, we are not aware of a systematic study comparing the oscillatory electro-oxidation of C1-C2 liquid fuels. Aiming at filling this gap, herein we report on some general features of the electro-oxidation of formaldehyde, formic acid, methanol, and ethanol on platinum, with focus on the comparison of the overall performance under conventional and oscillatory regimes. The comparison is carried out by different means and generalized by the use of identical experimental conditions in all cases. Namely the systems were kept as simple as possible and the following conditions were adopted: (electrochemically annealed) polycrystalline platinum as working electrode; all small organic molecules at 0.5 mol L-1 in a 0.5 mol L-1 H2SO4 aqueous solution; and T = 25.0 ± 0.1 °C. In order to provide a solid ground of the discussion that follows, the next section highlights some qualitative aspects of the electro-oxidation reactions considered under regular and oscillatory regimes and its relation with adsorbed species and different reaction pathways.

Reasoning Figure 1 displays schematically some relevant facets of the electro-oxidation of small organic molecules on a catalytic surface, normally platinum or platinum-based electrodes. Panel (a) represents a quasi-stationary potentiodynamic sweep and its general shape illustrates the initial negligible current in a considerable wide potential window, and thus the high overpotential usually found in such systems. The abrupt current increase coincides with the



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adsorption of oxygenated species from water, which assist the oxidation of poisoning species that block the electrode surface at lower potentials. The potential-dependent interaction between oxygenated species and the surface strengths as the potential increases and inhibits the adsorption of organic species; as a consequence, the current decreases and a peak or shoulder is observed. If the electrode potential is stepped from a small value to, say, the one indicated by the open circle in (a), the current evolution as a function of time can be represented as in (b). After the initial transient, the system relaxes to a stationary value. The stationary behavior at this point is a consequence of the time-independent reaction rates and, importantly, of the constancy of the coverage of different adsorbates. This situation is schematically given in terms of partially covered electrode surface by a given poisoning species P (filled circles), and reactions involving species A, B, and C. In this notation, A accounts for the small organic molecule under consideration, B for a partially oxidized soluble species, and C for the oxidized and desorbed poison. Reaction 1 represents the formation of the poisoning species P from A, reaction 2 the partial oxidation of A to B, and 3 to the oxidation of the poison P to the soluble species C; as it can be seen below for a more realistic, chemical, example.


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Figure 1: Schematic views of the electro-oxidation of a typical small organic molecule on platinum under different regimes: (a) potentiodynamic sweep, (b) constant applied potential (chronoamperometry), and (c) constant applied current (chronopotentiometry), in which the electrode potential undergoes spontaneous temporal oscillations. The poisoning species P is represented by the filled circle, see text for details. If a constant current of magnitude illustrated in the open circle in panel (a) is applied through the interface, the system might undergo an oscillatory instability. As a consequence, the electrode potential oscillates as given in Figure 1(c). In this situation, the coverage of different adsorbates and the reaction rates cause the observed oscillations in the electrode potential. Grosso modo, the cycle can be explained as follows. The electrode potential increases as the coverage of poisoning species increases in order to allow the flow of electrons imposed through the interface. When the potential reaches a certain value, the adsorption of oxygenated species starts and the poison P is oxidized, freeing thus the surface and causing the decrease in the electrode potential, and the cycle repeats itself. The cartoons of the interface along the oscillations illustrate the situation of lowest and highest (shortly before the oxidation of P at high potentials) coverages of the poisoning species. In this way, the coverage of poisoning species (and of other adsorbates), as well as the rate of all reactions



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involved switches between two limiting values and the overall dynamics can be followed by the oscillating electrode potential. The nature of species A, B, C, and P given in Figure 1 can be discussed in terms of the chemistry underlying the mechanism of a given particular reaction. In order to exemplify a plausible situation, if we take methanol as the starting species A, the following set of reactions can be written. Reaction 1 is the formation of the surface poison P, i.e. COad, from a methanol molecule (A): Pt + → PtCO + 4H+ + 4e-

(1)

In reaction 2, the partial oxidation of methanol (A) can, for instance, result in dissolved formic acid (B): Pt + + O → Pt + HCOOH +4H+ + 4e- .

(2)

Finally, the oxidation of adsorbed CO (P) to solution CO2 (C), can be represented as, PtCO + → 2 + C .

(3)

Of course, the reactions listed in this example are not intended to represent elementary steps. For a more realistic scheme of the electro-oxidation of methanol, the reader is addressed to references 36 and 66.

Experimental Polycrystalline platinum flags were used as working electrodes (WE), a platinized platinum flag with a high surface area served as counter electrode, and a reversible hydrogen electrode (RHE, with respect to which all potentials are quoted) prepared with identical concentration than that used in the support electrolyte was used as reference electrode. For all experiments, the WE was flame-annealed, with a butane/oxygen flame, for at least one minute and cooled inside the electrochemical cell, above the electrolyte, under argon flow. After that, the WE was cycled several times in the base electrolyte, from 0.05 V to 1.50 V, to achieve a stable and reproducible response. The active surface area of the WE was established from cyclic voltammograms recorded at 50 mV s-1 in argon-purged electrolyte by integrating the


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charge in the hydrogen ad/desorption regions, that is, from 0.05 V to 0.44 V. The active surface area was found to amount to 0.38 cm2, 0.26 cm2, 0.26 cm2, and 0.84 cm2, for experiments with formic acid, formaldehyde, methanol, and ethanol, respectively. The differences in the areas accounts for the exposed geometric areas of the independent experiments. Nevertheless, the electrode geometry and roughness was very similar in all cases. Prior all experiments, with the electrolyte containing 0.5 mol L-1 of the appropriate organic molecule, the WE was cycled 15 times at 0.10 V s-1, between 0.05 and 1.50 V. All systems were studied in a conventional three electrode glass cell, kept at 25.0 ± 0.1 °C by means of a thermostatized bath. The experiments were carried with an Autolab (302N) potentiostat/galvanostat. All solutions were prepared with high purity water (Milli-Q system, Millipore, 18.2 M

Ω cm), H2SO4 (Merck, 95-97%), HCOOH (Sigma-Aldrich, ≥ 98%), H3COH (J.T. Baker, 99.97%), CH3CH2OH (J. T. Baker, 99.95%). Before each experiment, the formaldehyde (Mallinckrodt Baker) was conditioned for 30 minutes at 60 °C in order to remove traces of methanol.

Results and Discussion

Initial Characterization. Figure 2 summarizes the potentiodynamic sweeps and typical potential oscillations for the electro-oxidation of (a) formaldehyde, (b) formic acid, (c) methanol, and (d) ethanol. The potentiodynamic runs (curves in black) were performed at 0.002 V s-1 and represent quasi-stationary responses. In these curves, the activity, as inferred by the peak current, decreases in the following sequence: formaldehyde ~ formic acid > methanol > ethanol. All current-voltage curves are relatively similar and a small modulation or shoulder prior to the main peak is apparent in most cases. In contrast to the thermodynamic potentials indicated in Table 1, detectable reaction rates under potential control are discernible only at relatively high potentials and an overpotential of at least 500 mV can be associated to the electro-oxidation of these molecules on platinum. This high overpotential is due to the formation of adsorbed species at low potentials, which are oxidized only at high potentials, as previously mentioned. This kinetic aspect critically limits the use of these organic molecules in low temperature fuel cells, and several strategies to improve the activity of platinum surfaces have been reported.67-77 Moreover, based on the these curves it is not possible to infer on the trends in the conversion to carbon dioxide, as partially oxidized species are commonly



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observed, and the use of in situ and on line approaches is mandatory for mechanistic investigations.78

Figure 2: Slow potentiodynamic sweeps (black lines) for the platinum electrode in 0.5 mol L-1 of (a) HCHO, (b) HCOOH, (c) CH3OH, and (d) CH3CH2OH, in 0.5 mol L-1 H2SO4 aqueous solution. From top to bottom, galvanostatic curves (red lines) obtained at constant applied current: (a) 2.6 mA cm-2, (b) 2.7 mA cm-2, (c) 0.77 mA cm-2, and (d) 1.1 mA cm-2. The mean potential at each cycle is given in green. T = 25 ºC. Black curves registered at dU/dt = 0.002 V s-1. As extensively reported for this class of electrochemical oscillators,23,63 spontaneous potential oscillations emerge for a certain range of applied currents. The curves in red in Figure 2 represent galvanostatic oscillations at selected applied currents within the oscillatory window,27 and the green points account for the mean potential at each cycle, Um, defined as:

1 ( )


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The first aspect to be pointed out is the fact that, in spite of the nearly indistinguishable potentiodynamic profiles, oscillations greatly differ in each case, c.f. Figure 4, discussed below. The particular characteristics of the oscillatory dynamics bring much more mechanistic information than, for instance, voltammetric or regular chronoamperometric curves. Actually, oscillations are much more sensitive to the reaction mechanism because the relative weights of the different reaction steps involved change as the surface coverage of different adsorbates autonomously oscillate. In contrast, conventional, non-oscillatory, stationary current and potential responses reflect a time-invariant contribution of distinct elementary steps involved. This aspect will be further explored below. The time scales of each oscillatory series are indicated in the vertical bars. Although all controllable experimental parameters are kept constant, a spontaneous drift is observed and the system develops in time due to, very, slowly evolving surface deactivation processes.19, 30,37,79-82

The oscillating time for each system depends, among other factors, on the applied

current. Nevertheless, the total duration or stability of all series presented in Figure 2 are comparable, and the shorter time observed for ethanol seems to result of the critical role played by partially oxidized species in the oscillatory dynamics in this system, and also of the accumulation of long-lived adsorbates.83

Electro-oxidation efficiency. Importantly for the electrocatalytic efficiency discussed is that oscillations considerably delay the surface inactivation, as for instance due to CO poisoning during the electro-oxidation of methanol.32 Therefore, the long-term performance is generally benefitted if the system is operated under oscillatory conditions. Oscillations in all cases presented in Figure 2 develop around the quasi-stationary current-potential curves. As the latter represent nearly stationary reaction rates, the potential window visited during oscillations clearly probes different surface coverages and reaction rates, presumably including also the state seen in the regular current-potential curve. Indeed, as self-organized coverage changes spans a considerably wide parameter range, it opens the possibility of periodic autonomous self-cleaning or oxidation of surface poisoning.84 As already mentioned in the introduction, the autonomous surface cleaning can promote noticeably increase of overall performance of polymer electrolyte membrane fuel cells (PEMFC) operated under oscillatory regime and fed with H2/CO mixtures.48-62 In spite of different possibilities of writing the efficiency (ε) of a fuel cell,85,86 it can be generally expressed in terms of the actual potential U(i) 86:



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( )

() "! (|$%&"'! | + |$ "! "!

%("'! |

+ ))

(1),

in which U(i) is the cell voltage under working conditions, i.e. at a certain current i, and Uocell is the thermodynamic open circuit voltage already presented in Table 1. The cell voltage U(i) depends on the flowing current and it is reduced with respect to Uocell because of cathodic and anodic overpotentials, ohmic drops from different origins and so on.86 For half-cell experiments, and neglecting all losses except the one due to the anodic reaction, one can write the efficiency as, " , "! *() +%&"'! () " !

(2).

We could do the same analysis with a time varying efficiency estimated under oscillatory regime via the oscillating electrode potential, U(t). It can be clearly seen that, in any case, the smaller the difference (U – Eoanode), the higher the overall efficiency. As given in Figure 2, the lower potentials visited during oscillations noticeably show a decrease in the overpotential of the anodic reaction. This aspect, together with the reported delay in the surface poisoning,32 presents important advantages to the use of autonomous oscillations to reach both higher and long-term activities. These aspects are obviously relevant for the performance improvement in energy conversion systems, which include the electro-oxidation of hydrogen carries, such as direct liquid fuel cells and electrochemical reforming. Moreover, these advantages remain important even if the time-averaged potential, i.e. the mean potential at each cycle (see the green points in Figure 2), remains higher than the stationary values under regular regime, because of the mentioned self-cleaning of the electrode surface and the consequent slower surface deactivation. In order to further illustrate the efficiency increase observed under oscillatory conditions, Figure 3 shows a power density versus current density plot for the electrooxidation of formaldehyde in identical conditions to that presented in Figure 2. The curve in red accounts for the galvanodynamic sweep, along which the electrode potential oscillates nearly over the whole current window. In contrast to that in Figure 2(a), where a constant current was applied, the current is slowly varied in this case. For comparison, the quasistationary potentiodynamic curve displayed in Figure 2(a) is also presented. For these plots, the overpotential, η, was estimated in terms of the equilibrium potential of a hypothetical fuel


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cell using formaldehyde and oxygen. Therefore, η = 1.35 - U, where 1.35 V is the thermodynamic open circuit voltage for the complete cell, c.f. Table 1, and U is the potential of the anodic reaction, measured versus a RHE, c.f. Figure 2. By doing so, all limitations in this idealized situation are attributed to the anodic reaction and the oxygen reduction is assumed to occur strictly at its equilibrium value of 1.23 V. Although recognizably simplified, this procedure allows comparing the power density under oscillatory and regular regimes; a more than two-fold enhancement is observed in the oscillatory case.

Figure 3: Power density versus current density curves for the electro-oxidation of formaldehyde under galvanodynamic (red) and potentiodynamic (black) regimes. dI/dt = 5 µA s-1. Remaining conditions as in Figure 2. Oscillation’s Features. As already mentioned, oscillations are more sensitive than nonoscillatory data to characterize an electrocatalytic system as they probe distinct reaction steps, spread over the corresponding potential window. Nevertheless, to obtain specific information on the nature and population of surface oscillating species, reactions involved, and so on, it is also required the employment of auxiliary in situ and on line techniques, vide infra. Still, the self-organized time-evolution observed during oscillations can be used as a diagnostic tool. For instance, the transition from low to high coverage of poisoning species schematically illustrated in Figure 1 can be inferred from the rate of increase of the potential U.46 In fact, at constant applied current, as the coverage of poisoning species increases in time, the electrode



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potential increases accordingly in order to permit the flow of electrons through the interface; consequently, the rate dU/dt at which the electrode potential increases reflects the rate of surface poisoning. Figure 4 shows typical oscillatory cycles as well as the rates dU/dt for the electro-oxidation of all molecules studied. The horizontal dashed lines at each U/t curves accounts for the mean potential, Um, for certain oscillatory cycle and the vertical dashed lines for the time window used to calculated the dU/dt rate in each case that accounts the minimum and maximum potential of each cycle.

Figure 4: U versus t (upper row) and dU/dt versus U (lower row) for one oscillatory cycle. Data extracted from Figure 2 for (a) and (e) formaldehyde, (b) and (f) formic acid, (c) and (g) methanol, and (d) and (h) ethanol. For all molecules studied, the rate dU/dt is low when the electrode potential is in its lower value, increases up to a maximum and then decreases; for formaldehyde and ethanol two maxima are observed. In terms of absolute values, the term dU/dt decreases in the sequence: formaldehyde > formic acid > methanol > ethanol. Except of the much higher dU/dt rates found for formaldehyde with respect to formic acid, this trend is comparable to that given above for the activity in the potentiodynamic profiles. In fact, high electrocatalytic activity to some extent implies high susceptibility to poisoning.28,78 We have used this concept to discuss the effect of tin modification on platinum towards the electro-oxidation of formic


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acid,46 and observed considerable smaller dU/dt rates for Pt3Sn intermetallic electrodes, which were more active than pure platinum. This contradiction with our present data is just apparent. The high activity induced by the Pt-Sn intermetallic electrode is mainly due to the facilitated oxidation of adsorbed carbon monoxide because of the formation of oxygenated species at lower potentials than that on platinum. Therefore, the activity enhancement observed in the potentiodynamic curves is reflected in the slowing down of the COad poisoning rates dU/dt.46 Otherwise, the case of the reaction of distinct species on identical catalysts is characterized by the different electro-oxidation mechanisms and, despite the nature of some adsorbates be the same for most of small organic molecules, c.f. COad, the adsorption state and the amount critically depend on the identity of the molecule. For instance, Jusys and Behm78 reported COad coverages of 87, 36, and 10% (of that of a saturated CO adlayer under identical conditions) after adsorption of formaldehyde, formic acid, and methanol on platinum, respectively. Although specific for the adsorption conditions reported, these results indicate a general trend. Table 2 compiles the results of the maximum dU/dt rate for each peak and the corresponding potential peak Up, the oscillatory amplitude ∆U, and of the oscillatory frequency f, for all cases studied. In addition, the times needed to reach the maximum potential along one cycle, tmax are also shown.

Table 2: maximum dU/dt rate and Up for an oscillatory cycle , the oscillatory amplitude (∆U), the time required for the electrode potential reach its maximum (tmax), and of the oscillatory frequency (f) for typical oscillatory cycles. Molecule Formaldehydea,b1 Formaldehydeb2 Formaldehydeb3 Formic acida Methanola Ethanola

dU/dt (Vs-1) 50.7, 40.3 19.1, 28.2 6.4, 8.5 8.9 1.2 0.4, 0.23

Up(V) 0.39, 0.75 0.37, 0.75 0.40, 0.74 0.75 0.75 0.47, 0.74

∆U (V) 0.12-0.90 0.29-0.92 0.37-0.94 0.57-0.83 0.59-0.81 0.47-0.82

a

t max (s) 0.054 0.17 0.28 5.3 0.80 4.1

f (Hz) 3.5 3.3 3.2 0.19 0.82 0.21

Data obtained from the selected regions presented in Figure 4. Data extracted before, around and after the abrupt variation in the mean potential, respectively, c.f. Figure 2 (a). b1, b2, b3



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As already mentioned, it seems undisputable that the spontaneous evolution of the electrode potential presented in Figure 4 carries more information on the electro-oxidation reaction than, for instance, the potentiodynamic curves presented in Figure 2. This is due to the fact that, under oscillatory regime, the importance of different reaction steps continuously changes along oscillations; whereas under regular regime, a single rate determining step dominates the observed profile. In fact, parameters described in Table 2 vary dramatically according to the reaction under consideration. Parameters depicted in Table 2 reflect: (a) the nature and (b) the population of adsorbed species, (c) their rates of adsorption and (d) the rates of reaction among adsorbate and/or solution species. Although they are intrinsic to each system, it is possible to find some regularity between them in the light of some mechanistic aspects. For formaldehyde, the oscillatory cycle showed in Figure 4 (a) starts at 0.12 V, passing through a maximum at 0.90 V, cf. ∆U in Table 2. The self-cleaning process during the oscillatory regime is probably more efficient than that for formic acid, methanol and ethanol, as suggested by this larger oscillation amplitude and oscillation frequency. The corresponding dU/dt curve in Figure 4 (e) showed two peaks for formaldehyde, one close to 0.40 V, presumably due to the high CO adsorption rates,87,88 at low potentials, and other peak close to 0.75 V, probably due to the adsorption of oxygenated species. In addition, the adsorption of HCOOad must be considered in the second peak, since its eventual decomposition to CO2 is rather slow.87-89 The smallest value of tmax, 0.054 s, found for formaldehyde can be rationalized as follows. Between 0.12 and 0.40 V, at the first dU/dt peak, hydrated formaldehyde (methylene glycol) dehydrogenation to COad has a fast kinetics,90 as suggested by the high value of dU/dt for the first peak, c.f. Table 2. If compared to those of formic acid and methanol, the COad adsorption from formaldehyde is the less susceptible to the hydrogen adsorption inhibition,78,91 and at this potential range (i.e. 0.12-0.40 V) COad accumulates, since the formation of oxygenated species is negligible.92 It must be taken into account that the accumulation of COad strongly inhibits the oxidation of hydrated formaldehyde,90,91 and the system readily loses its charge transfer capacity to maintain the applied current at low potentials, and thus the potential increases very fast. A second poisoning process involving HCOOad and oxygenated species is initiated at high potentials, causing a further abrupt increase of overpotential, giving rise to a second peak at about 0.75 V, as already mentioned.


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The time-series for formaldehyde can be divided in three regions: the first located before the abrupt variation of mean potential and the other ones around and after it, cf. Figure 2 (a). As the profiles of dU/dt are similar for the oscillations situated in all three regions, we present in Figure 4 (a) just one oscillation cycle of the first region and its corresponding dU/dt profile is displayed in Figure 4 (e). The specific characteristics of the oscillation cycles of each region are showed in Table 2. The time of transition from low to high overpotentials, tmax, increases throughout these three regions because the minimum potential also increases, cf. ∆U in Table 2, probably due to a slower kinetics of COad formation. In accordance, the maximum dU/dt related to the first peak (0.40 V) also decreases throughout the time-series and the same tendency is observed for dU/dt associated with the second peak (0.75 V). Since the rate of poisoning is decreasing, evidently other pathways must be favored in order to support the applied current, as for example, the oxidation of methylene glycol to formic acid, via bridge-bonded adsorbed formate. Of course, the COad oxidation above 0.55 V92 can not be underestimated and, therefore, the structure of the poisonous species, the kinetics of oxygenated species adsorption and of COad oxidation must have great impact in the oscillatory cycles, but, in situ and on line approaches are indispensable to do inferences in this case. These slowing down poisonous rates and the consequent increase of the contribution of other non CO pathways90 for the overall current clearly cause an improvement in terms of overpotential as can be deduced by the considerable decrease of the mean potential along the time-series, cf. Figure 2 (a). On the other hand, the decrease of the amplitude ∆U diminishes the efficiency of the surface reactivation process. For formic acid and methanol, the oscillation cycles start at around 0.60 V, reaching a maximum nearly of 0.80 V, cf. ∆U in Table 2. The dU/dt achieve a maximum close to 0.75 V most due to accumulation of oxygenated species and HCOOad. In the electro-oxidation of formic acid, the oscillatory cycle has a wide quasi timeindependent potential region centered around 0.6 V, before the abrupt increase of potential, cf. Figure 4 (b) and (f). Osawa and coworkers25 attributed the slow variation of the potential between 0.50 and 0.60 V to the slow rate of COad adsorption.25 Behm and coworkers93 have been reported a contribution of less than 1% of CO pathway for the overall current reaction around 0.60 V, at specific conditions. The abrupt increase of potential during the oscillatory cycle occurs due to the fast adsorption of HCOOad and oxygenated species, which suppress the direct oxidation of formic acid via reactive intermediate.94,95 The value for the time of transition from low to high overpotential of formic acid was larger than those observed for formaldehyde and methanol, cf. tmax in Table 2. This is an



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evidence of its superior capacity of charge transfer via reactive intermediate during its slow poisoning process. This behavior benefits the electro-oxidation of formic acid in terms of overpotential as can be observed by the mean potential in Figure 2 (b). Although the potential oscillations of methanol and formic acid happen approximately at the same range, cf. ∆U in Table 2, the short time interval between low and high overpotential found for the oscillatory cycle for methanol indicates a lower capacity of this system to maintain the applied current during its specific poisoning process, despite the low value of dU/dt. Since the site requirement for the dehydration of formic acid and methanol9698

is the same,99 there must some other features that justifies the differences in the oscillation’s

features. An important aspect to be considered is the non-electrochemical formation of carbon monoxide during the oxidation of formic acid. In contrast, in the electro-oxidation of methanol, all steps have charge transfer, including the CO formation that relives four electrons. This idea must be tested in numerical simulations. Autonomous changes in the coverage of carbon monoxide and bridge bonded formate in the range of 0.20-0.30 monolayer (ML) and 0.10-0.20 ML, respectively, have been observed in the electro-oxidation of formic acid on platinum.25 For methanol, changes in the coverage of carbon monoxide of about 0.24-0.37 ML have been reported.32 In the former case, the short-term experiments showed that the coverage of HCOOad increases and COad decreases as the applied current increases, and similar behavior was reported for formaldehyde.88 The long-term evolution of oscillations in the case of methanol evidences that the coverage of CO decreases in time. Those results suggested that the surface is not completely covered with COad and HCOOad. Since the potential oscillations occur in a high potential range, the oxygenated species could also play a role in the poisoning surface process. We have recently found40 that the electro-oxidation of 1 mol L-1 ethanol containing acid solution over platinum, at oscillatory regime, results in a rather small production of carbon dioxide, if compared with the production of acetaldehyde. Moreover, we showed some evidences that acetic acid is also produced as a partially oxidized species, and it is likely that adsorbates such as CHx,ads, COads, OHads, CH3CHOads,40,83,100-107 might take part in the oscillatory cycle. Despite of scarce mechanistic studies on the oscillatory electro-oxidation of ethanol, it is possible to establish a trend in the poisoning rate by comparison with ethanol and C1 molecules. As in the case of formaldehyde, the dU/dt profile of ethanol showed in Figure 4 (h) has two peaks: one close to 0.47 V and the other passing through a maximum at 0.75 V. The


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first poisoning process probably is caused by strongly adsorbed residues coming from the C-C bond break at low potentials,83,103,105,106,108 and the second one manly due to oxygenated species adsorption at more positive potentials, a common intermediate of the oscillatory mechanism of small organic molecules. Ethanol has a relatively slower poisoning process with a maximum absolute value close to 0.4 V s-1 and 0.23 V s-1, for the first and second peaks, respectively, c.f. Table 2. Besides these two kinds of poisoning process, a third is also seen: the production of soluble acetic acid creates an adsorption/desorption equilibrium that blocks surface sites.106 The oscillation cycle of ethanol electro-oxidation is the only one that spends several seconds in the superior potential limit, Figure 4 (d). This indicates a greater capacity to support the applied current through the interface at high potentials, and can be attributed to the fact that: (i) the poisoning process caused by oxygenated species is slow; (ii) ethanol residues formed in low potential range hinder the oxygenated species adsorption; and/or (iii) adsorbed oxygenated species are consumed to produce soluble acetic acid.

Conclusions In order to compare the activity under conventional and oscillatory regimes, we have performed experiments under identical conditions for the electro-oxidation of formic acid, formaldehyde, methanol, and ethanol, on polycrystalline platinum surfaces and in acidic media. The reaction rates under non-oscillatory regime were estimated by means of a quasistationary potentiodynamic sweep, and the activity, determined by the main current peak value, was found to decrease in the sequence: formaldehyde ~ formic acid > methanol > ethanol. Overpotentials of at least 500 mV were observed in all cases. Under oscillatory regime, the lower potentials visited during oscillations noticeably decreases the overpotential of the anodic reaction in any case. This results in a considerable increase in the overall conversion and thus in the system performance. An estimation of a hypothetical fuel cell operated with liquid formaldehyde showed a two-fold increase in the power density. Under optimized conditions and at higher temperatures this difference is expected to be considered increased. Studies in this direction with direct liquid fuel cells are in progress and will be available soon. Besides the performance increase, it must also be mentioned that, under oscillatory regime, the self-cleaning process, slows down the surface deactivation, favoring the long-term performance. These aspects are relevant for the activity enhancement in energy conversion systems, which include the electro-oxidation of hydrogen carriers, such as direct



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liquid fuel cells (DLFC) and electrochemical reforming. From the mechanistic point-of-view, the spontaneous potential oscillations in each system are greatly different, bringing additional information about reaction mechanism and surface coverage of different adsorbates when compared to conventional regime. This is due to the fact that under oscillatory regime more than one step is important to maintain the current imposed through the interface and their features depends on the nature and population of adsorbed species, their rates of adsorption and the rates of reaction among adsorbates and/or solution species. The poisoning rates under oscillatory conditions decrease in the sequence: formaldehyde > formic acid > methanol > ethanol. Except for the much higher rates founds for formaldehyde, this trend is comparable to that given above for the activity in the potentiodynamic profiles. As general perspectives, further developments in this direction would include similar investigations at higher temperatures (close to those where DLFCs are operated), spectroscopic characterization of adsorbate dynamics under oscillatory conditions, and, of course, extensive tests in real DLFCs.

Acknowledgements MVLD (Grant No. 160511/2011-9) and HV (Grant No. 304458/2013-9) acknowledge Conselho Nacional de Desenvolvimento Científico e Tecnológico (CNPq) for financial support. LFS (Grant No. 2013/24162-0), and HV (Grants Nos. 2012/24152-1 and 2013/16930-7) acknowledge São Paulo Research Foundation (FAPESP) for financial support. HV acknowledges Prof. E. A. Ticianelli for fruitful discussions.

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Electrocatalytic Efficiency of the Oxidation of Small Organic Molecules

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