H2O2 direct fuel cells

Electrochimica Acta 178 (2015) 163–170

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Perovskite cathodes for NaBH4/H2O2 direct fuel cells D.M.F. Santosa,b,* , T.F.B. Gomesa , B. Šljuki ca , N. Sousab , C.A.C. Sequeiraa , F.M.L. Figueiredob a b

Materials Electrochemistry Group, CeFEMA, Instituto Superior Técnico, Universidade de Lisboa, 1049-001 Lisbon, Portugal CICECO, Dep. Enga . Materiais e Cerâmica, Universidade de Aveiro, 3810-193 Aveiro, Portugal

A R T I C L E I N F O

A B S T R A C T

Article history: Received 3 June 2015 Received in revised form 16 July 2015 Accepted 26 July 2015 Available online 29 July 2015

Perovskite-type oxides are evaluated as cathodes for hydrogen peroxide (H2O2) reduction in direct borohydride fuel cells (DBFCs). The study is conducted on button-type ceramic electrodes of LaCoO3 (LC), La0.84Sr0.16CoO3 (LSC), La0.8Sr0.2Fe0.8Co0.2O3 (LSFC), and La0.7Sr0.3MnO3 (LSM), thus avoiding crosscontributions of carbon or nickel catalyst supports. Cyclic voltammetry shows that LSM has the highest activity for H2O2 reduction in alkaline solution, with the other three materials showing minimal (LC) to none (LSC, LSFC) electroactivity. The data also suggest that only LC and LSM are stable within the tested potential window, although the alteration of the samples surface is apparent in scanning electron microscopy images collected after the electrochemical measurements. The analysis of LSM by chronopotentiometry in light of the Sand equation indicates ca. 1 electron involved in the H2O2 reduction. A demonstration DBFC employing a single phase LSM ceramic cathode yields a peak power density of 8.2 mW cm
2 at 28 mA cm
2, at 45  C. This value is about 60% of that obtained with a commercial platinum foil electrode, which is a notable feature and demonstrates the potential of LSM as an alternative low cost cathode for DBFCs. ã 2015 Elsevier Ltd. All rights reserved.

Keywords: perovskite cathode hydrogen peroxide reduction direct borohydride fuel cell

H2O2 + 2e
! 2OH


1. Introduction The direct borohydride fuel cell (DBFC) is a particularly interesting technology that uses an alkaline solution of sodium borohydride (NaBH4) as a fuel fed directly to the anode, overcoming the transportation problems usually associated with hydrogen (H2). Moreover, DBFC reactions occur in alkaline media, where in principle non-precious metal electrocatalysts can be used without major performance loss [1]. The borohydride (BH4
) oxidation at the DBFC anode proceeds according to Eq. (1), with a standard electrode potential (E0) of
1.24 V vs. standard hydrogen electrode (SHE). NaBH4 + 8OH
! NaBO2 + 6H2O + 8e


(1)

The product of this reaction is sodium metaborate (NaBO2), inert and non-toxic, and can be recycled back to NaBH4 [2]. Typically, the cathodic reaction in DBFCs is the oxygen (O2) reduction but lately there has been much interest in using hydrogen peroxide (H2O2) as the oxidant. In that case, the cathodic process is described by Eq. (2).

* Corresponding author. E-mail address: [email protected] (D.M.F. Santos). http://dx.doi.org/10.1016/j.electacta.2015.07.145 0013-4686/ ã 2015 Elsevier Ltd. All rights reserved.

(2)

Accordingly, the overall cell reaction can be written as Eq. (3). NaBH4 + 4H2O2 ! NaBO2 + 6H2O 0

(3)

The E of 0.87 V vs. SHE for the direct reduction of H2O2 (Eq. (2)) is higher than that for O2 reduction (0.40 V vs. SHE). Consequently, direct H2O2 reduction in the direct borohydride/peroxide fuel cell (DBPFC) yields larger theoretical cell voltages (2.11 to 3.01 V, depending on the pH) and specific energy (17 Wh g
1) compared to those of the DBFC (1.64 V and 9.25 Wh g
1). The reported significant increase of power density of a fuel cell with H2O2 as oxidant results from a lower activation barrier and hence faster kinetics of two-electron H2O2 reduction compared to the fourelectron O2 reduction [3]. Additional benefit of using H2O2 as oxidant comes from the fact that it is liquid and therefore its storage, transport, handling and controllable feeding to a fuel cell is easier than in the case of O2 gas. Still, H2O2 as oxidant in fuel cells has some drawbacks as it is prone to spontaneous decomposition into O2 and H2O in the presence of some metals and it is unstable at higher temperatures [9]. Furthermore, fuel cell membrane can deteriorate in the presence of such strong oxidant as H2O2. In general, the main operational issues of the DBPFC are the hydrolysis and thus incomplete BH4
anodic electrooxidation [4] and the

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mentioned chemical decomposition of H2O2 at the cathode [5]. Therefore, the development of anode and cathode electrocatalysts having high selective catalytic activities for BH4
oxidation and H2O2 reduction, respectively, has been a key point in DBPFC research [6–11]. Platinum-based catalysts, similarly to gold, palladium and silver, possess good catalytic activity for H2O2 reduction reaction (HPRR) but have the disadvantage of high cost [9]. Perovskite-type materials, with general chemical formula ABO3 where A is a lanthanide or an alkaline earth metal and B is a transition metal, are known to be good electrocatalysts for the oxygen reduction reaction (ORR) in fuel cells and other energy conversion devices, both at low and high temperatures [12–17]. Only few examples of their use in DBFC technologies have been reported, despite their acceptable ORR activity and BH4
tolerance in membraneless DBFCs [18,19]. Subsequent reports have confirmed the potential of La1
xSrxCoO3 [20] and LaNi0.8Co0.2O3 [21], as cathodes for regular BH4
/O2 DBFCs. Examples of perovskitetype cathodes in DBPFCs are not known, but the few and recent studies concerning the H2O2 electroreduction on perovskite electrodes suggest a potentially interesting electrocatalytic performance of La1
xSrxMnO3d [22,23], LaNiO3 [24], La1
xCaxCoO3
d [25], and LaCoO3 [23], with the manganites displaying somewhat higher activity [23]. To the best of our knowledge, the HPRR on ferrites has not been assessed so far. These studies are usually based on optimised electrode designs consisting of mixtures of the catalyst powders with carbon (of various kinds and in proportions approaching 50 vol.%) and an organic binder, which in some cases are pressed onto nickel foam current collectors [15,16,26], themselves also being able to catalyse the reaction. While it has not been demonstrated for the HPRR, there is growing evidence that the addition of carbon to transition metal oxides (e.g., LaCoO3 and La0.8Sr0.2MnO3d) can enhance the electrocatalytic activity for the ORR in alkaline media by up to one order of magnitude. The underlying mechanism seems to be the electroreduction of O2 to H2O2 on the carbon, which is then decomposed by the oxide [27,28]. This means that the identification of the exact role of the oxides on the performance of these otherwise composite electrodes comprising the carbon and the transition metal oxide particles, implies the knowledge of the behaviour of the oxides themselves. This aspect, certainly hidden by the best performance of the carbon-containing electrodes, remains unexplored in the literature, with few and very recent known reports [27,28]. The investigation of the complex relationships between the perovskite composition and their intrinsic electrocatalytic properties for the HPRR is of extreme importance to rationalise and optimise this type of electrodes, since the vast compositional variety potentially offered by the perovskite structure is difficult to cover on a basis of “synthesise and check”. Here we present an electrochemical study of four different perovskite-based materials, namely LaCoO3 (LC), La0.84Sr0.16CoO3 (LSC), La0.7Sr0.3MnO3 (LSM) and the ferrite La0.8Sr0.2Fe0.8Co0.2O3 (LSFC) with the objective of assessing the effect of the transition metal on the electroreduction of H2O2, and their performance as cathodes in a DBPFC. The study involves the characterisation by cyclic voltammetry and chronopotentiometry of button electrodes made from dense ceramic pellets to avoid cross-effects of other factors, namely porosity or the mentioned contribution of additional components, such as carbon or nickel supports. Finally, we compare the performance of two equivalent laboratory DBPFCs, one with a commercial platinum foil cathode, and another with the most electroactive perovskite oxide, which is a highlighted contribution of this work.

2. Experimental 2.1. Preparation and characterisation of the electrodes Commercial powders of the four tested perovskites (LaCoO3
d, La0.84Sr0.16CoO3
d, La0.8Sr0.2Fe0.8Co0.2O3-d and La0.7Sr0.3MnO3d, all from Seattle Specialty Ceramics, now Praxair Specialty Ceramics) were shaped into discs by applying uniaxial pressure (70 MPa) in order to obtain powder compacts (8 mm in diameter and 1– 2 mm thick) with a fractional density equal or higher than 50% of theoretical value. These compacts were sintered in air at 1425  C to obtain ceramic samples with fractional density higher than 92%, thus free from percolating porosity. Silver paint (SPI, high purity) was used to glue each ceramic disc to a copper wire, with a thin glass tube protecting the wire. The discs were then mounted in an epoxy resin in a proportion of 5 parts of resin (EpoxiCure Epoxi Resin 20-8130-032) to 1 part of hardener (EpoxiCure Epoxi Hardener 20-8132-008). Prior to the electrochemical measurements the surface of the electrodes was polished using SiC abrasive papers (Struers) of decreasing grit sizes (1200, 2400, 4000) until reaching perfectly smooth surfaces. The geometric area of the electrodes was determined to be 0.35 cm2. The surface and the cross-section of the ceramic electrodes before and after the electrochemical measurements were observed by scanning electron microscopy (SEM) on a Hitachi SU-70 microscope equipped with Bruker QUANTAX 400 energy-dispersive Xray spectroscopy (EDS) detector. The structural studies were carried out on X-ray diffraction (XRD) patterns collected on the surface of the ceramics with a Rigaku Geigerflex D/Max-C series diffractometer using CuK radiation and 2u = 3  min
1. The patterns were analysed with PowderCell v2.4 [29] by refining the scale factor, zero shift, background, lattice parameters and peak profile parameters. Two series of patterns were collected, one on the surface of freshly polished samples and another after the electrochemical measurements. Finally, the room temperature electrical conductivity (s) of the ceramic pellets was confirmed by 2-probe AC method. The pellets were painted with a commercial silver paste (Algar) on opposite surfaces, and placed on an alumina jig with platinum (Pt) wires to measure the electrical resistance (R) with an Agilent 2980A meter using a test signal amplitude of 100 mV and a frequency range between 20 Hz and 2 MHz. The impedance of the short-circuited cell was subtracted to the measurements in order to correct for the resistance of the wires, and also to minimise the effect of the inductance in the platinum wires. The R values were estimated as corresponding to the high frequency intercept with the real axis. The conductivity was obtained through s = L(RS)
1, where S is the area of the electrodes and L is the thickness of the pellets. The results, presented in Table 1, were found to be in good agreement Table 1 Electrical conductivity, lattice parameters and unit cell volume (indexed on R-3c space group) estimated from the XRD patterns before and after the electrochemical measurements. Electrode

a, b / Å

c/Å

V / Å3

Electrical Conductivity / S cm
1

LC

before after

5.426(6) 5.428(0)

13.061(8) 13.049(1)

333.11 332.96

0.048

LSC

before after

5.431(7) 5.431(1)

13.245(1) 13.261(3)

338.42 338.76

0.832

LSFC

before after

5.442(0) 5.445(5)

13.239(5) 13.238(4)

339.56 339.97

0.029

LSM

before after

5.504(9) 5.504(8)

13.362(6) 13.362(1)

350.68 350.66

0.117

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165

with the literature data, except the case of LSC, which displays a somewhat lower conductivity [30–32]. 2.2. Electrochemical measurements Cyclic voltammetry (CV) and chronopotentiometry (CP) measurements were carried out in a conventional three-electrode cell, with the perovskite disc serving as working electrode, a saturated calomel electrode (SCE) reference (Hannah Instruments) and a large Pt mesh (Johnson Matthey, 100 cm2) counter electrode. The electrochemical cell was connected to a PAR 273A computercontrolled potentiostat (Princeton Applied Research, Inc.) with the associated PowerSUITE software package. The electrode potential values are reported against the SCE reference. The electrochemical glass cell had an effective volume of 125 ml. A thermostatic water jacket was used for controlling the electrolyte solution temperature (0.1  C), set by a recirculating water bath (Ultraterm 6000383 P-Selecta). The test solutions contained H2O2 (35 wt.%, Merck) concentrations ranging from 0.05 to 0.4 M, in 2 M sodium hydroxide (NaOH, 99 wt.%, AnalaR) electrolyte. This low concentration range was selected for the preliminary studies in order to minimise the peroxide decomposition to H2O and O2. CV scans were carried out at rates between 5 and 200 mV s
1, in the 25–55  C temperature range. Five CV cycles were run for each experiment with only the 5th cycle being considered, as stability was observed after the 4th cycle. Additional CV measurements in the absence of H2O2 were recorded in O2-saturated 2 M NaOH solution and in deaerated 2 M NaOH solution. CP measurements were carried out in the 25–55  C temperature range, with current densities ranging from 5 to 10 mA cm
2, using a 0.3 mA cm
2 interval step. A laboratory fuel cell was assembled using a LSM electrode as the cathode and a Pt mesh as the anode. The anolyte solution contained 1 M NaBH4 (96 wt.%, PS, Panreac) in 4 M NaOH supporting electrolyte, while the catholyte was 5 M H2O2 in 4 M NaOH, which we found to be the optimal conditions in terms of fuel cell performance [5]. It should be noticed that practical BH4
/H2O2 fuel cells use similar amounts of H2O2 in the catholyte [5,6]. Each cell compartment contained 70 ml of the corresponding electrolyte, which were separated by a Nafion N117 membrane (30 cm2 active area, DuPont, Wilmington, DE). Nafion N117 was selected for the separator as previous studies on cation- and anion-exchange membranes (CEM and AEM) for DBPFCs showed that the highest power density is obtained with this type of separator [33,34]. Crossover by diffusion of species such as NaOH, NaBH4 and H2O2, gases or organic compounds, can occur due to their concentration gradient between the anode and cathode compartments. It has been suggested that CEMs could be more efficient than AEMs in the suppression of BH4
crossover due to the negative charge of this ion [33]. Yet, for long time operation CEMs can lead to precipitation of borate and depletion of OH
near the anode, and to a build-up of OH
near the cathode, thereby reducing the fuel cell efficiency. Thus, though a CEM was used for demonstration of the DBPFC concept, for a real, practical system an AEM would be required. The cell polarisation and corresponding power density curves were obtained as a function of temperature, and directly compared to a second DBPFC assembled with benchmark Pt foil cathode (Metrohm 6.0305.100) with 1 cm2 geometric area. 3. Results and discussion 3.1. Cyclic voltammetry measurements CV was used to investigate the activity of the perovskite materials towards H2O2 electroreduction. The CVs were run in

Fig. 1. CV scans of the four tested perovskite electrodes at 25  C. Scan rate: 25 mV s
1.

three different electrolyte media, namely deaerated 2 M NaOH solution, O2-saturated 2 M NaOH solution, and 0.4 M H2O2 in 2 M NaOH solution. Fig. 1 compares the 5th cycle obtained for the four studied perovskite electrodes in the three media at 25  C.

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The LC perovskite displays some activity towards H2O2 reduction (Fig. 1a) as evidenced by an increase of current density in the presence of H2O2, but with no distinct feature observed in the CV at the expected potential region. In agreement with the literature, the negligible currents observed in the absence of H2O2 in O2-saturated 2 M NaOH solution indicate a poor activity for the ORR [35]. The constant cathodic current also shows that the material remains stable within the studied potential window. Considering the relatively close composition of LC and LSC, and also the higher electronic conductivity and the higher bulk (possibly, also surface) oxide vacancy concentration of LSC in comparison to LC, it is somewhat surprising that the former material shows no activity for H2O2 reduction (Fig. 1b). This result is in apparent contradiction with the 60–80 mA cm
2 cathodic currents reported for a closely related La1
xCaxCoO3
d (x = 0.2, 0.4, 0.5) perovskite-based electrode consisting of a mixture of the oxide with 15 wt.% carbon black and 10 wt.% polytetrafluoroethylene (PTFE) supported on Ni foam [25]. These data, however, do not show any clear trend with the level of Sr substitution, and the authors suggest the onset of the H2O2 electroreduction at 0.23 V vs. SHE based on CVs. In fact, the cathodic scans are nearly flat (for x = 0.4) or similar to that in Fig. 1a (for x = 0.2). Thus, direct comparison with [25] is not possible, where the data do not seem to support a correlation between the high current density and the electroreduction of H2O2. In fact, the LSC bulk electrode (Fig. 1b) shows identical behaviour in the three electrolytes, with a marked increase of current at potentials more negative than
0.6 V, suggesting that the material may be reduced. This is not totally unexpected since the partial substitution of La3+ for 30 at.% Sr2+ leads to an increase of the oxide vacancy concentration compensating the negative Sr’La defects and destabilising the structure [36]. A slightly higher activity for the ORR has indeed been reported for La0.8Sr0.2CoO3-d in comparison to LaCoO3 [37], but the overall currents measured herein are rather low in both cases. Even so, the results suggest that LSC is not active towards H2O2 reduction. The LSFC perovskite electrode showed cathodic currents at potentials more negative than
0.5 V, even in deaerated solution, making LSFC unsuitable for the envisaged application. Higher cathodic currents could be observed in the presence of H2O2 than in the O2-saturated or in the deaerated electrolytes (Fig. 1c), but voltammograms obtained with the LSFC electrode in H2O2 solution at different potential scan rates did not show any differences among them, which is a somewhat unusual behaviour. The apparent instability of the LSFC is also not totally expected in comparison to LSC, since the partial substitution of Co by Fe should lead to increased stability [14]. This suggests that pH may change the formation energies of the perovskites. To the best of our knowledge, Pourbaix diagrams exist only for the individual components of the perovskite structure (La, Sr, Mn, Co and Fe) [38]. The intrinsic redox behaviour of the simple oxides with respect to the perovskites may obviously be different due to the different chemical environment, but still worth considering in the present context. Lanthana and strontia readily form hydroxides in contact with water. At pH = 14, La(OH)3 can be reduced to La at E <
2.8 V vs. SHE; and Sr(OH)2 to Sr at E <
2.9 V vs. SHE, with formation of SrH2 at E <
1.55 V vs. SHE. Strontium peroxide can actually be stabilised under anodic polarisation at E > 0.65 V vs. SHE (but this is not relevant for the present study). From the three transition metals, Fe displays the most stable oxidation states, with Fe2O3 appearing as the stable phase down to E 
0.6 V vs. SHE (always at pH = 14), which is roughly the cathodic limit we have used in our study. Metallic Fe is formed at E  -0.9 V vs. SHE. On the contrary, Co shows the least stable oxides and the highest tendency to form the hydroxide (Co(OH)2), which tends to be stable between E  0.1 V vs. SHE and E 
0.75 V vs. SHE, and is reduced to Co for

lower potentials. The trends for Fe and Co might suggest an intermediate behaviour for the LSFC due the mixed composition. Finally, Mn shows an intermediate behaviour and, while it seems more reducible than the Fe oxides (Mn2O3 is reduced to Mn3O4 at E 
0.2 V vs. SHE), Mn3O4 is expected to prevail down to E 
0.42 V vs. SHE, before forming Mn(OH)2. From this analysis, LSM could be the most stable compound. The LSM electrode exhibited the highest activity for H2O2 reduction, displaying a current density of about 2 mA cm
2 at
0.6 V (Fig. 1d), which is about 10 times higher than that observed for the LC electrode. Moreover, only residual currents were observed in the electrolyte solutions where H2O2 was not present, showing that the electrode is stable even at high overpotentials. This confirms previous results obtained for electrodes consisting of La1
xSrxMnO3 powders mixed with carbon and supported on Ni foams [22,23]. This result is in line with the reported higher electrocatalytic activity for the ORR of manganites in comparison to cobaltites or ferrites [12,39]. The effect of the potential scan rate on the H2O2 reduction behaviour of LSM was observed in the 25–55  C temperature range. Fig. 2 shows linear scan voltammograms obtained for the LSM perovskite electrode scanning the potential from the open circuit potential to
0.9 V using scan rates ranging from 5 to 200 mV s
1 at 35  C. As expected, the voltammograms showed increasing current densities for increasing potential scan rates. 3.2. Structure and microstructure Fig. 3 shows a series of SEM micrographs taken at the surface of the pellets after the electrochemical measurements. The four materials show clear signs of modification of the surface microstructure, which reveals a contrasted pattern of small, rounded particles with a highly heterogeneous distribution of size. In contrast, the bulk of the pellets appears clean with the micrographs revealing the look of typical fracture surfaces of fairly dense ceramics with grain sizes in the 2–4 mm range. The EDS analysis confirmed the nominal composition of the samples both on the surface and the cross-sections, which suggests that the heterogeneities observed on the surface are not coupled to major compositional differences with respect to the bulk of the pellets. This is also supported by the XRD analysis of the pellets before and after the electrochemical measurements. The XRD patterns collected on the surface of fresh pellets confirmed the expected single phase for all materials, showing only reflections due to the perovskite structure (bottom grey lines in Fig. 4). The formation of additional crystalline phases suggesting a major modification of the materials is not apparent in any of the

Fig. 2. Effect of the potential scan rate on the CVs obtained with the LSM electrode in 0.4 M H2O2 + 2 M NaOH at 35  C.

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167

Fig. 3. SEM micrographs taken at the edge of a fracture revealing the surface and the cross-section of the ceramic electrodes after the electrochemical measurements.

patterns collected after the electrochemical measurements. All patterns can be indexed on the R-3c rhombohedral space group. One can observe for the case of LC and LSC a slight shift of the entire pattern towards higher angles and some differences in the relative intensity of some reflections (also noticeable for LSFC). These can be accounted for by sample displacement or preferential orientation of the grains, both likely to occur since the patterns were collected on the surface of sintered pellets. Lastly, we note that the peaks breadth is larger than expected for a sample with the micron-sized particles displayed in Fig. 4. But since the XRD patterns collected before and after the electrochemical measurements are very similar, the peak broadening should result more from sample displacement and/or preferential orientation than from diffracting nano-sized crystallites.

The lattice parameters estimated from the refinements are similar before and after the electrochemical measurements (Table 1), thus suggesting that the microstructural changes observed by SEM are either related to the loss of crystallinity of the material, to more subtle modifications of the short range order, or simply to sample removal by dissolution of the surface. The unit cell volume is actually identical for the case of LSM, which may indicate a slightly enhanced stability of the manganite in comparison to the other materials, as suggested by the Pourbaix diagrams mentioned in the previous section. Obviously, differences may indeed exist at a length scale smaller than the spatial resolution of the technique, which is in this case determined by the penetration of the X-ray beam (Cu Ka radiation at 40–45 kV) in the sample (about 1 to 2 mm). This means that any

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possible to determine the transition time, t, corresponding to the time step necessary to consume the H2O2 at the electrode surface until complete depletion. It can be seen that t decreases with increasing current density. Moreover, plotting t1/2 as a function of j
1 led to straight lines with very good fitting to Sand equation (Eq. (4)), with R2 values of 0.99 both at 25  C and 35  C (Fig. 6). Therefore, Sand equation (Eq. (4)) can be applied in this case in order to determine the number of exchanged electrons, n, from the slope of the plots of Fig. 6 [40],

t1=2 ¼ nFCðpDÞ1=2 =2j

(4)
1

Fig. 4. XRD patterns collected at the surface of the electrodes before (grey lines) and after the electrochemical measurements (black lines).

modification of the surface during the electrochemical measurements must be limited to the atomic planes close to the surface, which represent a very small fraction of the diffracting sample volume and hence not detectable by XRD. Instead, other methods specific for surface analysis such as XPS must be used in order to clarify the stability of these materials in strong alkaline media and under polarisation. 3.3. Chronopotentiometry measurements In order to gain some insight in the charge-transfer process, the work was pursued with a more detailed study of the LSM electrode by chronopotentiometry (CP). Fig. 5 shows CP curves for the reduction of 0.05 M H2O2 in 2 M NaOH, using current densities ranging from 6.9 to 9.4 mA cm
2. As expected, the increase of the applied current densities led to a gradual decrease of the electrode potential step for H2O2 reduction. By identifying the inflexion points of the CP curves it is

Fig. 5. CP curves for H2O2 reduction at LSM electrode at 25  C.

where F is Faraday’s constant (96485 C mol ), D is the diffusion coefficient of H2O2 (cm2 s
1) and C is its molar concentration (mol cm
3). The electrochemical surface area is assumed to be close to the geometric area of the simple button electrodes used in this work. The D values for H2O2 must be estimated from the literature. Here we assume measurements from Borggaard [41], who calculated D as a function of temperature, yielding values of 1.30  10
5 (25  C), 1.73  10
5 (35  C), 2.15  10
5 (45  C), and 2.56  10
5 cm2 s
1 (55  C). The lack of the exact knowledge of the electrochemical area and the D values may indeed limit the applicability of the Sand equation, but the fact that it provides a first approximation to the mechanism involved fully justifies its use in the present context. The obtained n values were found to be 1.2 and 0.7 for 25  C and  35 C, respectively. The higher n for 25  C may be justified by the higher stability of H2O2 alkaline solutions at room temperature. Additionally, the fact that these n are lower than the theoretically expected 2 electrons (Eq. (2)) may be due to several factors that do not seem to be related to the electronic or ionic conductivities of the perovskites (approximately varying in the range 0.03–1 S cm
1, Table 1). The low electron transfer observed can, in fact, be related to the complex nature of the perovskite heterostructures with strong electron correlations [42], complemented by the dissociative O2 chemisorption and the regeneration of OH
on the surface [12,23,37], or even a diffusion controlled process other than the diffusion of H2O2 in the NaOH solution. The results obtained with the Sand equation must be confirmed by more detailed studies of the mechanism of H2O2 discharge on these materials (e.g., by chronoamperometry with ultramicro disc electrodes). This is even more so in face of the promising results obtained with the LSM cathode on a DBPFC, as presented in the following section. 3.4. Fuel cell testing Typical polarisation curves of the DBPFC with the LSM cathode obtained at different temperatures are presented in Fig. 7. The open circuit voltages (OCVs) of these alkaline fuel cells were ca. 1.0 V. The observed OCVs were lower than the theoretical cell voltage of 2.1 V, but in line with those previously reported in the literature [43]. The corresponding power density curves are also shown in Fig. 7, with peak power densities attaining values of 6.1, 6.7, and 8.2 mW cm
2 for the temperatures of 25, 35, and 45  C, respectively. Enhancement of cell performance with the temperature increase could be clearly observed, as evidenced by higher current densities and higher power densities. The performance of this DBPFC employing LSM cathode was compared to a DBPFC using a benchmark Pt cathode operating under the same experimental conditions. As shown in the inset of Fig. 7, at 25  C the assembled Pt-cathode DBPFC had a peak power density of 11.3 mW cm
2 at a cell voltage of 0.25 V and current density of 45 mA cm
2. Under the exact same conditions, the LSM-cathode DBPFC had a peak power density of 6.1 mW cm
2 at a cell voltage of 0.3 V and current density of 20 mA cm
2. It has been reported [22,44–46] that some

D.M.F. Santos et al. / Electrochimica Acta 178 (2015) 163–170

perovskite-based electrodes can catalyse H2O2 disproportionation, generating O2 in solutions of high H2O2 concentration. As mentioned earlier, H2 evolution may also occur at the anode side due to borohydride hydrolysis. Hence, these possibilities need to be considered in the design of DBPFCs, implying additional tests to assess to which extent the above chemical reactions would take place in such cells. The performance of the LSM-based cell is about 60% of that of the DBPFC with Pt cathode. Although promising, the present results still do not confirm a potential for practical application of LSM, but they demonstrate the existence of a significant activity for the H2O2 reduction that needs to be understood in order to be further improved. As starting point, one may consider the design principles recently proposed to improve the ORR activity of transition metal perovskite oxides [12] as a guide to optimise the composition of LSM-type materials by changing the electronic configuration of the metal (so that the d eg orbital is filled with 1 electron) and by promoting a higher covalency of the metal– oxygen bond. This can be achieved by suitable doping or by the control of the processing conditions and/or high temperature isothermal annealing under variable oxygen partial pressure. It is also appropriate to note that mixing these oxide electrocatalysts with carbon could enhance the overall performance of the electrode by enhancing the catalytic or electrocatalytic activity for the electroreduction of H2O2 (besides increasing the electronic conductivity), as observed for the ORR. In any case, the clarification of the role of carbon is also necessary for the technological application of these perovskite-type electrocatalysts in DBPFCs and other related technologies. 4. Conclusions The possibility of using perovskite-based materials as cathodes for the electroreduction of H2O2 in DBPFCs was investigated by assessing the intrinsic electrocatalytic activity of pure ceramic specimens, as opposed to the usual powdered catalyst/carbon mixtures. LaCoO3 (LC) and La0.7Sr0.3MnO3 (LSM) electrodes show catalytic activity for H2O2 reduction in alkaline media, with the latter showing the most promising results. Chronopotentiometric analysis of LSM indicate a number of 1.2 electrons involved in the H2O2 reduction at 25  C, but this must be confirmed by more detailed and precise studies. One key factor determining the potential application of these materials is the stability under the strongly alkaline conditions of DBPFCs. SEM observations of the samples combined with XRD data collected on samples before and after the H2O2 reduction experiments revealed a significant modification of the surface that is related to the loss of crystallinity. While there is no apparent

169

Fig. 7. Effect of temperature on the polarisation behaviour and corresponding peak power density curves for a DBPFC using LSM cathode. Inset shows direct comparison of the LSM and Pt cathodes at 25  C.

correlation between this observation and the electrochemical activity of the various electrodes, it is a key conclusion of the present work and implies a future detailed study by more appropriate techniques (e.g., TEM, Raman and XPS), particularly because this type of information remains scarce in the literature. Simple studies based on the post-mortem analysis of powdered samples after immersion in alkaline solutions, complemented with the elemental analysis of the remnant solutions in search for dissolved metals, may provide valuable information on the dissolution rate of these materials in strong alkaline media. Finally, a DBPFC operating at 45  C employing the LSM cathode was demonstrated. This fuel cell attained a peak power density of 8.2 mW cm
2 and a current density of 28 mA cm
2 at a cell voltage of 0.3 V, which is about 60% of the values obtained using a Pt cathode, in the exact same conditions. Considering the prohibitive cost of Pt, this result is very promising and demonstrates the potential of lanthanum manganites as a low cost alternative cathode for DBPFCs. Acknowledgements The authors would like to thank FCT, the Portuguese Foundation for Science and Technology, for postdoctoral research grant no. SFRH/BPD/77768/2011 (B. Šljuki c) and PhD grant no. SFRH/BD/ 89670/2012 (N. Sousa). FCT is also acknowledged for contract no. IF/01084/2014/CP1214/CT0003 under IF2014 Programme (D.M.F. Santos), for Investigador FCT 2013 contract no. IF/01174/2013 (F.M. L. Figueiredo), and for the funding through project CICECO-FCOMP01-0124-FEDER-037271 (Ref. FCT PEst-C/CTM/LA0011/2013). References

Fig. 6. t1/2 vs. j
1 plots for H2O2 reduction at LSM electrode.

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