The Role Platinum Nanoparticles Play in the Kinetic Mechanism of

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C: Energy Conversion and Storage; Energy and Charge Transport

The Role Platinum Nanoparticles Play in the Kinetic Mechanism of Oxygen Reduction Reaction in Nonaqueous Solvent Nelson A. Galiote, Ulderico Ulissi, Stefano Passerini, and Fritz Huguenin J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b02606 • Publication Date (Web): 26 Jun 2018 Downloaded from http://pubs.acs.org on June 26, 2018

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The Journal of Physical Chemistry

The Role Platinum Nanoparticles Play in the Kinetic Mechanism of Oxygen Reduction Reaction in Nonaqueous Solvent

Nelson A. Galiote1,2,3, Ulderico Ulissi2,3, Stefano Passerini2,3,*, Fritz Huguenin1,*

1

Departamento de Química, Faculdade de Filosofia, Ciências e Letras de Ribeirão Preto – Universidade de São Paulo, 14040-901 Ribeirão Preto (SP), Brazil 2

3

Helmholtz Institute Ulm (HIU), Helmholtzstrasse 11, 89081 Ulm, Germany

Karlsruhe Institute of Technology (KIT), P.O. Box 3640, 76021 Karlsruhe, Germany

*e-mail: [email protected] *e-mail: [email protected] ACS Paragon Plus Environment

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The Journal of Physical Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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ABSTRACT

Although lithium-oxygen batteries theoretically have high energy density, they are far from becoming viable and practical alternatives to current batteries because the oxygen reduction reaction (ORR) is slow, which limits the applicability of lithium-oxygen batteries, and these power sources require suitable catalysts for optimum performance. Understanding the chemical and electrochemical steps involved in the ORR mechanism is mandatory. This work investigates platinum nanoparticles dispersed in graphitized carbon as a positive electrode for lithium-air batteries, characterizes the materials by cyclic voltammetry and galvanostatic cycling, X-ray diffraction, and scanning electron microscopy coupled with energy-dispersive X-ray spectroscopy, and determines the kinetic constants for the reaction mechanism steps with the help of electrochemical impedance measurements and modeling. The electron transfer to oxygen molecules adsorbed onto platinum nanoparticles is the rate-limiting step. Li2O2 is preferentially produced via the electrochemical pathway instead of the chemical disproportionation reaction, leading to a device with improved reversibility and enhanced energy density.

KEYWORDS: Lithium-air batteries; oxygen reduction reaction; platinum nanoparticle; electrochemical impedance spectroscopy.

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INTRODUCTION Renewable sources, such as wind and solar power, could help to meet the growing energy demand of the modern society. Because these sources are intermittent by nature, high power and high energy storage systems are fundamental for their effective and efficient use1–3. Lithium-air batteries are a promising system with high theoretical energy density and could replace the state-of-the-art lithium-ion batteries in some applications4,5. However, despite their advantages, many challenges still have to be addressed before lithium-air batteries become a commercially viable option. Air cathode clogging, solvent/electrolyte stability, and low energy efficiency are just some examples of the factors that limit their (dis-)charge capability and cycle life2. Although much has been achieved over the past decades, metal-air batteries with enhanced performance can only be proposed after the ORR kinetics are elucidated and issues related to the air cathode are overcome. Most challenges regarding lithium-air batteries refer to the ORR end product, Li2O2: this compound has a wide band gap, which hinders electronic transport, and it is insoluble in practically all the solvents that are usually employed in these batteries, so it accumulates and passivates the electroactive sites. One approach to surmount the issues concerning Li2O2 consists in using heterogeneous electrocatalysts that can enhance sluggish ORR and oxygen evolution reaction (OER) rates during battery discharge and charge, respectively. Gasteiger et al.6 reported that, compared to electrodes without catalysts (Vulcan carbon), electrodes bearing catalysts (Au, Pt, Pd, or Ru) shift the onset potential of the ORR to more positive values (i.e., the overpotential was lower) in O2-saturated LiClO4 0.1 mol.L-1/ethylene glycol dimethyl ether. These authors also observed that the relative oxygen adsorption energies are strictly related to the electrocatalytic effect, which lowers the ORR overpotential (with the characteristic volcano trend), and they showed that Pt and Pd are the best electrocatalysts. McCloskey et al.7 employed electrodes including catalysts like Pt and Au as well as electrodes without catalysts (carbon electrodes), to find ACS Paragon Plus Environment

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that the former electrodes perform better during the ORR, while the electrodes containing Pt also have the lowest charge potential during the OER. Previous results also indicated that bimetallic nanoparticles composed of Au and Pt exhibit high bifunctional activity with high catalytic activity for the ORR and the OER8. Nevertheless, recent studies have shown that CO2 evolution due to electrolyte decomposition predominates during the charging step when Pt-based electrodes are employed7. This suggests that Pt is not selective for Li2O2 decomposition, which favors unwanted side reactions. Notwithstanding this lack of selectivity, Pt-based electrodes can still be considered as a viable alternative because lithium superoxide and peroxide species can easily bind to oxygen defects in oxides, to diminish their reactivity toward side reactions9. This is the case of Pt nanoparticles dispersed on zirconia support. The strength of Li2O2 adsorption onto the catalyst surface can alter the reaction product morphology and particle size, which leads to low electrooxidation potentials and minimizes secondary product formation. This strategy promotes high specific capacity and cyclability for electrodes consisting of copper oxide supported on Pt nanosheets 10. Given that Pt-based electrodes with suitable design can be a viable alternative to lithium-air batteries, kinetic studies on Pt nanoparticles (Pt-nps) can improve understanding of the ORR mechanism in aprotic media. Some studies have attempted to model this reaction by considering the electrode material, solvent type, and discharge rate when solving transport equations based on timedependent techniques. However, these same studies have failed to account for all the details in a mathematical model, which has prevented them from gathering reasonable kinetic information to circumvent the issues involving these systems11–13. Here, we have investigated how Pt-nps influence the ORR mechanism by using electrochemical impedance spectroscopy (EIS). This approach enabled us to examine the ORR under the minimal effect of side reactions because the potential and frequency ranges were specifically chosen for this purpose. Through comparison with our earlier results for Pt ACS Paragon Plus Environment

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bulk electrodes 14, we have been able to interpret results associated with the use of nanoparticles, which proved to be more suitable for practical batteries. In addition, we have determined the rate-limiting step, the rate constant for each elementary step, preferential reaction pathways, and the number of a transferred electron per reduced oxygen molecule for Li2O2 formation, which will contribute to the development of efficient cathodes for practical lithium-air batteries. EXPERIMENTAL Dry ethylene glycol dimethyl ether (water content at 4 ppm as detected by Karl Fisher titration) was purchased from BASF. Lithium perchlorate (99.99% purity), phosphorous pentoxide ACS reagent (purity ≥ 99%), tetraethylene glycol dimethyl ether (purity ≥ 99%), platinum nanoparticles (Pt-nps, size ≤ 5 nm) on graphitized carbon with 40 weight (wt) % loading, 1-Methyl-2-pyrrolidone (NMP, anhydrous, purity ≥ 99.5%), and anhydrous 2-propanol were acquired from Sigma-Aldrich and used as received. Alcohol-based Li+-exchanged Nafion® binder 10 wt % (LITHion®) was supplied by Ion power, Germany. Ether-based electrolytes were prepared by dissolving LiClO4 to a final concentration of 0.1 mol L-1. The electrolytes were stored inside a glove box MBRAUN with water and oxygen content lower than 0.1 ppm. Pt-np ink was prepared according to a procedure described elsewhere15. Briefly, Pt-nps on graphitized carbon were mixed with the LITHion® dispersion and diluted with 1 mL of 2-propanol. The ink-like dispersion was kept under continuous stirring. The final catalyst/binder ratio was 5:1. With a microsyringe, 1-2 µL of this solution was deposited on alumina-polished glassy carbon electrodes, which were later used in a three-electrode glass cell investigation. Electrodes for coin cell assembling were prepared by thin film casting of slurries on SIGRACET® Gas Diffusion Layer (GDL) 35bc (thickness = 325 ± 25 µm, areal weight = 110 ± 10g m-2) with a doctor blade. Electrodes without catalysts consisted of 80 wt % of C-NERGYTM SUPER C65 (Imerys) and 20 wt % of polyvinylidene


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fluoride 6020 (Solef Solvay). The powders were thoroughly mixed with mortar and pestle, dispersed in NMP, and cast as a thin film on GDL. The average loading of SUPER C65 was 0.4 mg cm-2. After coating, the electrodes were dried at 60 °C, cut (diameter = 16 mm), dried under vacuum at 120 °C overnight, and transferred to a glove box without contact with air atmosphere. These electrodes were designated as SUPERC65. Analogously, the Pt-np electrodes were prepared by casting the slurry (5:1 catalyst/binder, see above for details). After drying at room temperature (few minutes for complete 2propanol evaporation), disk electrodes (on GDL) were cut (diameter = 16 mm), dried under vacuum at 75 °C for 24 h, and transferred to a glove box without contact with air atmosphere. The average total mass loading (Pt-nps+graphitized carbon+binder) on GDL was ≈ 0.9 mg cm-2, which will be designated as the total mass loading when the specific current density or specific capacity is concerned. XRD measurements were performed on a Bruker® D8 Advance diffractometer equipped with a CuKα source (λ = 0.154 nm). 2θ ranged from 30° to 80°; the step size was 0.007°, and the time per step was 2.75 s. The complete scan was obtained after 5 h. Before measurements, coin cells were disassembled inside the glove box; cathode electrodes were washed several times with dimethyl carbonate (DMC, Merck battery grade, water content less than 10 ppm) to remove electrolyte traces and transferred to a closed sample holder to avoid any contact with air moisture. The PDF files used to identify the products were: 01-074-0115 for Li2O2, 00-004-0708 for LiOH, 00-001-0996 for Li2CO3, and 00-046-0902 for PtCl2. Scanning electron microscopy (SEM) images were obtained with a Zeiss LEO 1550 VP microscope equipped with an energy-dispersive X-ray spectroscopy (EDS) detector (XMaxN 50, Oxford Instruments, analyzed area = 30 x 30 µm2). Before analysis, discharged and charged electrodes were rinsed with DMC, dried under vacuum inside a glove box, and transferred to the SEM chamber without contact with air. The GC-MS analyses were accomplished on a Shimadzu QP2010 (Shimadzu Corporation, Kyoto, Japão) chromatograph linked to a Shimadzu QP2010Plus mass spectrometer system equipped with an ACS Paragon Plus Environment

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electron ionization source (EI-EM) operating at 70 eV. Chromatographic separation was carried out on a fused silica Rtx5-MS (Restek) capillary column (30 m x 0.25 mm i.d., 0.25-µm film) composed of 5% diphenylsiloxane and 95% dimethylpolysiloxane. The temperature was programmed to rise from 50 to 280°C at a rate of 10°C·min-1. Helium was used as carrier gas. The following parameters were employed: split ratio = 1:100, injector temperature = 250 °C, and ion source temperature = 250 °C. The mass spectra were recorded from 40 to 600 Da with a scan interval of 0.5 s. Ultrapure N2 (99.999% purity) or O2 (99.998% purity) from Linde-AGA was purged into the solution for 20 min before the experiments; the gas flow was maintained during the electrochemical measurements. Cyclic voltammetry (CV) and electrochemical impedance spectroscopy (EIS) measurements were conducted with a potentiostat/galvanostat (Solartron 1287) coupled with an impedance/gain-phase analyzer (Solartron SI 1260). A conventional three-electrode cell containing 10 mL of electrolytic solution was employed; the working electrodes were Pt-nps deposited on aluminapolished 3-mm-diameter glassy carbon (BAS, Japan). The reference electrode was assembled by immersing a silver wire in AgNO3 (99.9999% metal basis, Sigma Aldrich) 0.01 M and tetrabutylammonium perchlorate (TBAClO4, Sigma Aldrich) 0.1 M in acetonitrile (99.8% anhydrous solvent further dried over molecular sieves, Sigma Aldrich). All the potentials were converted to the standard Li/Li+ electrode potential (3.04 V vs standard hydrogen electrode). The program Maple® v13.0 was used to fit the EIS data. Electrochemical charge/discharge tests were carried out with the MACCOR Battery tester 4300 at room temperature by holding the cells in a thermostatic climatic chamber (20 ± 2 °C). Galvanostatic tests started with a three-hour rest step to ensure that the electrolyte wetted the electrodes and separator homogeneously. Pt-np and SUPERC65 electrodes were investigated in Li-O2 coin cells, assembled inside an argon-filled glove box (MBRAUN) with oxygen and water content lower than 1 ppm. Meshed CR2032 coin cell cases (MTI Corp.), glass fiber separators (Whatman GF/A), and 14-mm-


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diameter metallic lithium battery grade from Rockwood Lithium (≥ 99.8 % purity) as anode material was used. After assembly, the cells were transferred to an airtight glass tube (Figure S1) and purged with O2 for 20 min prior to testing. RESULTS AND DISCUSSION Figure 1(a) shows the first voltammetric profile in O2-saturated LiClO4 0.1 mol L-1/ethylene glycol dimethyl ether obtained for the Pt-np and Pt foil electrodes. The inset graph shows the magnification of the onset potential region. Compared to the Pt bulk electrode, the Pt-np electrode enhances the current density upon the ORR and OER while shifting the ORR onset potential to higher values (2.88 and 2.7 V for the Pt-np and Pt foil electrodes, respectively; see inset in Figure 1a). The shift in the onset potential indicates that Pt-nps have higher catalytic activity; i.e., they increase the energy density of a lithium-air battery. These values also agree with literature data6,16.


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200

a

Pt-np Pt

j / µ A.cm-2

100 0 0 -10

-100

-20 -30

-200

-40

-300

-50 2.4 2.6 2.8 3.0 3.2 3.4

2.0

2.5

3.0

3.5

4.0

4.5

+

E / V (vs Li/Li ) 4.5

b

4.0

E /V (vs Li/Li+)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


3.5 3.0 2.5 2.0 0

2

4

6

8

10

12

14

16

Capacity / mC

Figure 1: a) Voltammetric profile at a scan rate of 20 mV s-1 for (continuous line) Pt-nps and (dashed line) Pt foil electrodes. The inset figure shows a magnification of the onset potential region for both electrodes. The arrows indicate the potential scan direction. b) (Dis)-charge curves for Pt-np electrodes at 5 µA. The electrolytic solution consisted of O2-saturated LiClO4 0.1 mol L-1/ethylene glycol dimethyl ether.

Figure 1b displays the (dis-)charge curves recorded between 2.0 V and 4.2 V for the Pt-np electrodes at 5 µA; the electrode geometrical area and the electrocatalyst mass are 7.0 mm2 and 45 µg, respectively. On basis of these curves, the charge associated with the discharge process is much higher than the charge involved in the charging process. This coulombic irreversibility is attributed to


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formation of a significant amount of reaction products, which leads to electrode coating (i.e., passivation) and hinders the battery charging process17–20. To investigate the reaction products that contribute to the kinetic studies, we used a higher amount of electrode material. We investigated the Pt-np/GDL electrode, to find that the charge capacity rather than the cut-off potential limits the curves (Figure 2 and Figure S2). This happens because the products that originate at low overpotentials are more readily oxidized, to improve process reversibility and to provide better understanding of the reaction kinetics11, as will be discussed below. We discharged the battery to a specific capacity of 1000 mA h g-1 by application of a constant current of 50 µA, which corresponds to 50 mA g-1. We employed a similar measurement protocol to characterize the Super C65/GDL conductive carbon electrode (Figure 2) without catalysts, to gain better understanding of the catalytic effect of the Pt-nps. The first capacity-limited (dis-)charge galvanostatic cycle of the Ptnp electrode reveals discharge potential up to 50 mV higher than the discharge potential of the Super C65 electrode. Moreover, the two electrodes differ widely during the charging process: the Super C65 electrode faces a larger overpotential as compared to the Pt-np electrode. The fact that the Pt-np catalytic effect benefits the OER more than the ORR process agrees well with literature works reporting that Pt is more active toward the OER6,8,9. Furthermore, the overpotential increases as a function of the oxidation charge for the Super C65 electrode, suggesting Li2O2 exhaustion. In contrast, the charge potential of the Pt-np electrode reaches a plateau for the highest oxidation charge values, which suggests that a higher amount of oxidizable chemical species (for instance, Li2O2) is produced by non-electrochemical steps; i.e., their amount is not limited by the faradaic reactions involved in the discharge step.


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4.5 4.0

E / V (vs Li/Li+)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


3.5

2.7 Pt-np 1st cycle Super C65 1st cycle

2.6

0

200

400

600

800

Capacity / mA h g

1000

-1

Figure 2: Galvanostatic discharge/charge of coin cells at 50 mA g-1 and limiting the capacity to 1000 mA h g-1 for () Pt-np and (---) for Super C65. The electrode geometrical area is 2.0 cm2.

Figure 2 reveals a minimum at 55 mA h g-1 (3960 s) and 7 mA h g-1 (504 s) for the Pt-np elecrode and Super C65, respectively, and a maximum at 500 mA h g-1 (10 h) and 400 mA h g-1 (8 h) for the Pt-np electrode and Super C65, respectively. At non-steady states, these profiles with minimum and maximum values are probably associated with slow partial dissolution of the intermediate LiO2, which is adsorbed onto substrate sites facing the solution21 and contributes to increasing the electrode potential (see Eq. 2). Superoxide anions and lithium ions in solution form LiO2 on Li2O2 particles, which accumulate on the electrode (Li+(sol) + O2-(sol) LiO2(s)) and produce more Li2O2 by disproportionation (2LiO2(s) Li2O2(s) + O2(g)) via a solution-based mechanism22. The increase in particle size over time also contributes to the observed profiles because large Li2O2 particles tend to be non-conducting and to increase the consumption of lithium ions and superoxide anions in solution during the discharge process due to their higher surface area21. This increases the kinetic overpotential (because of the decrease in the electroactive area) and decreases the concentration overpotential


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(because of the consumption of superoxide anions in solution and of the coverage degree associated with LiO2 adsorbed onto active sites according to the equilibrium LiO2(ads) Li+(sol) + O2-(sol)). We performed EDX mapping of the electrodes together with the SEM images (Figure 3). Figure 3a illustrates the SEM image used for the Pt (Mα1) mapping overlaid in (b). On the basis of this overlaid image, the brighter spots in the SEM micrograph in Figure 3a can be associated with the nonhomogeneous distribution of Pt-nps; i.e., agglomeration sites. This agrees with literature data on Pt-nps on carbon matrixes4. Figures 3c and 3d display SEM images of the discharged electrode (at the specific capacity of 1000 mA h g-1) with increased magnification of 100 kx. The red rectangle in Figure 3c marks deposits of reaction product(s) (Li2O2). Intentional Li2O2 decomposition by beam focalization reveals the presence of Pt-np sites under the decomposed product(s) in Figure 3d. These figures hint that the bright Pt-np sites are more catalytically active for LiO2 and Li2O2 formation during ORR than the carbon sites.


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Figure 3: SEM and EDX mapping of Pt-np/GDL electrodes. (a) SEM image and (b) elemental mapping of Pt (Mα1) overlaid in the SEM image of the pristine electrode. (c) and (d) show the SEM image of one discharged electrode (1000 mA h g-1) with 100 kx magnification; in (c), the red rectangle emphasizes the Li2O2 deposits. (d) shows the SEM image after Li2O2 degradation by electron beam focalization, marked by the red rectangle, and evidences Pt-np sites under the Li2O2 deposits.

To examine the formation of intermediates and products associated with the ORR on the Pt-np electrode surface, we accomplished ex-situ SEM and XRD measurements. Figure 4a, 4b, and 4c illustrate the SEM images of pristine, discharged, and charged Pt-np electrodes, respectively; Figure 4(d) shows the respective ex situ XRD patterns. The SEM image of the discharged electrode (by a galvanostatic step at 50 mA g-1 to a limited capacity of 1000 mA h g-1) reveals that the discharge products are embedded in the porous matrix. In particular, dispersed Li2O2 particles measuring between 100 and 300 nm are easily observed on the electrode (compare with the SEM of the pristine electrode (Figure 4a)). According to the XRD profile, the product is mainly Li2O2 along with LiOH traces, which ACS Paragon Plus Environment

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point to a small degree of side reactions. This is expected during the discharge process as a result of solvent degradation due to superoxide nucleophilic attack, for instance. Other possible LiOH sources are the superoxide ion arising during the first electron transfer, which reacts with traces of water molecules from air contamination or with water residues in the electrodes from the alcohol-based Li+exchanged Nafion® binder (i.e. Lithion®)23–26. On the other hand, the XRD pattern of the charged electrode does not show any feature ascribed to the discharged product Li2O2 or the secondary product LiOH. Additionally, the XRD pattern and the SEM micrographs of the charged electrode and the pristine electrode are rather similar, indicating that the dispersed Pt-nps are very active in the OER process, thereby improving system reversibility. The absence of secondary products in the XRD pattern of the charged electrode does not exclude side reactions27–29 because the XRD technique might not be able to detect small amounts and/or amorphous products. However, at high discharge depth (e.g., 4000 mA h g-1) in potentiostat conditions at 2.21 V (vs Li/Li+), the diffractogram indicates that electrolyte degradation products are formed (see Figure 5).


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Figure 4: (a-c) SEM images of the pristine, discharged, and charged electrodes (capacity limited at 1000 mAh g-1), respectively. (d) Respective XRD profiles. Total mass loading of 1.3 and 0.9 mg for the discharged and the charged electrode, respectively.


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Current / A g-1

2.0

a 1.5

1.0

0.5

0.0 0

1000

2000

3000

4000

Capacity / mAh g-1 Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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3

b

2

PtCl2

LiOH/Li2O2

Li2CO3

1

Potentiostatic

LiOH/PtCl2 Li2O2

GDL

Li2CO3

GDL

Platinum

0

30

35

40

45

50

55

60

2θ θ / degrees

Figure 5: (a) Pt-np coin cell discharging until 4000 mA h g-1 in a potentiostatic condition of 2.21 V (vs. Li/Li+). (b) shows the XRD profile of the respective discharged electrode. Mass loading of 1.1 mg.

Figure 6a presents the galvanostatic discharge conducted at a current density of 50 mA g-1 up to a capacity of 4000 mA h g-1. Figure 6b refers to the respective diffractogram of the discharged electrode. The XRD pattern does not evidence that carbonate is formed at this discharge depth, differently from the potentiostatic condition at 2.21 V (Figure 5). Carbonate formation is normally observed for discharged electrodes consisting of carbonaceous materials due to reaction between Li2O2 and carbon30. Given that the ORR preferentially occurs on Pt-nps than on the carbonaceous matrix at higher potentials, and on the basis of the SEM images and EDX data shown in Figure 3, the diffractogram depicted in Figure 5 suggests that graphitized carbon participates more effectively in the ORR at lower potentials only—the detected Li2CO3 corresponds to reaction involving Li2O2 adsorbed onto carbon sites. Figure 4d gives no evidence of Li2CO3 formation during the charging process, either, even though Li2CO3 is normally produced by carbon decomposition in the presence of Li2O2 or its ACS Paragon Plus Environment

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intermediates at potentials more positive than 3.5 V (vs. Li/Li+) 30. All these findings corroborate with the preferential Li2O2 formation on Pt sites at higher discharge potentials.

E / V (vs Li/Li+)

3.0

a 2.9

2.8

2.7

0

1000

2000

3000

4000

Capacity / mAh g-1 Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


b

Li2O2

Platinum GDL

GDL

Li2O2

Li2O2

LiOH

30

35

40

45

50

55

60

2θ θ / degrees

Figure 6: (a) Galvanostatic discharge of a Pt-np electrode in coin cell configuration up to 4000 mA h g1 at 50 mA g-1. (b) ex-situ XRD profile of the discharged electrode. Total mass loading deposited on GDL = 0.8 mg.

Figure 7a and 7b/7c respectively show the Nyquist and Bode diagrams obtained for the Pt-np electrodes at 2.65 V, 2.41 V, and 2.21 V dc potentials superimposed by an ac amplitude of 5 mV within 300 Hz to 13 mHz. We measured impedance after pre-treatment for at least 7200 s (after application of the dc potentials), which provided a pseudo-steady state. In Figure 7a, the charge transfer resistance (as determined from the intercept of the semicircle at low frequency) decreases at low potentials. However, the resistance rises from 15 kΩ.cm2 (at 2.41 V) to 18 kΩ.cm2 (at 2.21 V) possibly because Li2O2


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agglomerates emerge on the electrode surface, as shown above, partially blocking the electrocatalytic sites. To obtain the kinetic parameters from EIS, we employed a kinetic model that had been previously developed on the basis of the chemical mechanism described below14. The elementary steps are based on the ORR mechanism proposed in the literature 31–34,

K1  → S − O2( ads ) S + O2 ← 

(1)

k2 S − O2 ( ads ) + Li + + e−  → S − LiO2( ads )

(2)

k3  → S 2 − Li2O2( ads ) + O2 2 S − LiO2( ads ) ← '

(3)

k4 S − LiO2( ads ) + Li + + e−  → S − Li2O2( ads )

(4)

k5  → S + Li2O2 S − Li2O2( ads ) ← '

(5)

k3

k5

where the subscript “ads” means that the species is adsorbed onto the substrate (S) active sites. The corresponding rate laws are obtained on the basis of these elementary steps. To simplify the kinetic investigation, we have not considered the solution-based mechanism because this process is slow and can be disregarded at the high-frequency range 35, and the amount of LiO2 dissolved in DME is low 22. Due to the oscillating perturbation potential in the transient state, the rates associated with Equations 2, 3, and 4 are expressed as ν2, ν3, and ν4, respectively:


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 ∂v2 v2 = v20 +   ∂θ LiO 2 

  ∂v  ( jωt ) +  2  ∆Ee( jωt )  ∆θ LiO2 e   ∂E θLiO2 E

(6)

 ∂v3 v3 = v30 +   ∂θ LiO 2 

 ( jωt )  ∆θ LiO2 e  E

(7)

 ∂v4 v4 = v40 +   ∂θ LiO 2 

  ∂v  ( jωt ) +  4  ∆Ee( jωt )  ∆θ LiO2 e   ∂E θLiO2 E

(8)

where j is the imaginary unit, ω is the angular frequency, and vi0 is the steady-state rate for elementary step i. These equations involve partial derivatives of the elementary step rates with respect to θ LiO , 2

which is the coverage degree associated with the LiO2 intermediate, and these partial derivatives are resolved using the mass balance during the steady-state regime. The terms ∆θ LiO2 and ∆E correspond to 0 the θ LiO = θ LiO + ∆θ LiO exp( jω t) oscillation amplitude of and to the sinusoidal ac potential, 2

2

2

0 respectively, where θ LiO is the coverage degree of the adsorbed LiO2 intermediate at the steady state. 2

The terms ∆θ LiO2 and ∆E can be associated with the temporal variation of θ LiO : 2

Γ

dθ LiO2

dt

= ΓFjω∆θ LiO2 exp( jωt ) = v2 − v3 − v4

(9)

where Γ is the maximum surface coverage. Hence, the faradaic impedance can be calculated from the oscillatory faradaic current density (∆if) for each angular frequency, through Equations 6, 8, and 10.

Zf =

∆E ∆E = ∆i f F ν 2 − ν 20 + ν 4 − ν 40

(

(10)

) ACS Paragon Plus Environment

19


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After mathematical treatment, the final equation to model the impedance response is based on Equation 11. The rate constants are scanned; the selected ones are based on the smallest difference between the experimental modulus and the theoretical impedance modulus36. The methodology used to determine the errors for each rate constant is described in the Supporting Information (see also Figures S3a-f).

−bE   k2 K5K1[O2 ]e−bE   k2 K5 K1[O2 ](1−θLiO2 )be + k4e−bE   − − k4θLiO2 be−bE  − −bE   K1K5[O2 ] +1+ K5  K1K5[O2 ] +1+ K5   + k2 K5K1[O2 ](1−θLiO2 )be − k θ be−bE Z f −1 = − 4 LiO2 ' −bE k2 K5[O2 ]K1 e k3[O2 ] K1K5[O2 ] +1+ K5 + 2k3θLiO2 + + k4e−bE + IωΓF K1K5[O2 ] +1+ K5 K1K5[O2 ] +1+ K5

(11)

Considering that GC-MS measurements (Figure S4) did not detect degradation or reaction products in the electrolytic solution after the (dis)-charge cycle shown in Figure 1b (probably due to their low concentration resulting from the small ratio between faradaic charge and electrolytic solution volume), these chemical species should not influence the electrochemical response of the Pt-np electrode in solution during the EIS measurements. Even so, we only plotted the modeling results obtained from Equation 11 for the EIS data obtained at 2.65 V because the parallel reactions preferably take place at lower potentials 30,37, as observed in the XRD pattern at 2.21 V (Figure 5). We used a double layer capacitance (Cdl) value of 130 µF cm-2 in the model. However, Figure 7a evidences that the theoretical impedance diagram deviates from experimental results probably because the electrode surface is inhomogeneous. Consequently, the adjustment with the Constant Phase Element (CPE, 1/Cdl(jω)n with 0 < n < 1) is better than the adjustment with the pure capacitive element (double layer capacitance). Cyclic voltammetry in acid medium (Figure S5) helped to determine the


20

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value of Γ (7.29 x 10−10 mol of sites per cm2), where an atomic hydrogen-associated oxidation charge of 210 µA per square centimeter of the real area was considered for polycrystalline platinum38,39. Taking the values obtained for the rate constants listed in Table 1 into account, the ratedetermining step (RDS) is, as expected, the superoxide formation step involving k214,31. Electrodes containing Pt-nps increase the RDS by two orders of magnitude as compared to the Pt bulk electrode (around 10-6 mol cm-2 s-1 based on previous studies14). This specific modification in the RDS might result from the Pt particle size and could explain the shifted onset potential. Therefore, the Pt-np key role is specifically modifying the RDS, enhancing the activity toward the ORR6,14,40.


21


a

-Z'' / kΩ Ω.cm2

40

13 mHz

30

20

10

0

13 mHz

0

5

10 15 20 25 30 35 40 45

Z' / kΩ Ω.cm2 80

b Phase Angle

60

40

20

0 -2

-1

0

log f

1

2

50

|Z| / kΩ Ω cm2

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 31

3

c

40 30 20 10 0 -2

-1

0

1

2

3

log f Figure 7: (a) Nyquist diagrams for Ptnp at (●) 2.65 V, (▲) 2.41 V, and (□) 2.21 V dc potentials. (b) and (c) are Bode plots. (○) Theoretical Nyquist diagram obtained from the kinetic model. (─) Theoretical modeling of experimental results obtained by using CPE instead of double layer capacitance. High k3 and k4 values as compared to k2 indicate Li2O2 formation as confirmed by the XRD results. These results reflect in the appearance of two oxidative processes in Figure 1, which displays


22

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voltammetric peaks at 3.30 V and at 3.85 V, related to LiO2 and Li2O2 electrooxidation, respectively. We can disregard the back reaction of Equation 3 because the k3' and oxygen solubility (9.57 mmol cm-3 41) values are low. On the basis of these kinetic data, we can still determine the k4/k3 ratio, which points to the preferential pathway for Li2O2 formation; Equations 3 and 4 correspond to the disproportionation reaction (chemical step) and to the second electron transfer (electrochemical step), respectively. We determined a value of 1.46 for the k4/k3 ratio, which suggests higher preference for the electrochemical step. This Li2O2 formation pathway helps to increase the battery energy density because more electrons can be transferred per oxygen molecule. On the basis of the k4/(k3 + k4) ratio and considering that one electron should be transferred in the first step, we estimate that about 1.6 electrons are transferred per reduced oxygen molecule, which is higher than the transfer observed for the carbon electrode without catalyst 37. This indicates the large energy conversion efficiency of the Ptnp sites for the ORR in the aprotic medium.

Table 1: Kinetic and equilibrium rate constants for the Pt-np electrode at 2.65 V. K1

0.18 ± 0.01

k2

k3

k3 '

k4

k5/k5’[Li2O2]

/mol cm-2 s-1

/mol cm-2 s-1

/cm s-1

/mol cm-2 s-1

/cm3 mol-1

(2.0 ± 0.1).10-4

1.77 ± 0.04

(9.10±0.28).10-2


2.610 ± 0.005

23.01 ± 0.39

23


Page 24 of 31

CONCLUSIONS Pt-nps influence the ORR and the OER in the aprotic medium in the case of lithium-air batteries. The onset potential of the ORR shifts to more positive values when Pt-nps dispersed in graphitized carbon are used instead of the Pt bulk electrode or carbonaceous materials only. Moreover, the overpotential associated with the OER is lower for the Pt-np electrodes as compared to the pure carbonaceous electrode. According to the XRD and electrochemical measurements, the Pt-np electrodes produce a higher amount of Li2O2 than the carbon-based electrochemical system and suitably oxidize Li2O2 deposited on the electrode surface. On the basis of an electrochemical method in the frequency domain, the rate-limiting step is the first electron transfer for LiO2 formation. Moreover, the values determined for each elementary step suggest that 1.6 electrons per reduced oxygen molecule are transferred for Li2O2 formation. The observed preference for the electrochemical pathway rather than the chemical disproportionation reaction also indicates higher energy density for the lithium-air batteries. Although lithium-air batteries made from positive electrodes composed of Pt-nps dispersed in the carbonaceous material are far from being used in practical energy storage devices, alternatives can be proposed for optimum performance to be achieved. The kinetic studies reported herein enable better understanding of the role that a nanocatalyst plays during the reactions occurring in lithium-air batteries, possibly contributing to the development of positive electrodes granting good performance of these power sources in the future.

ACKNOWLEDGMENTS We are grateful to FAPESP (Projects 2012/21629-1, 2011/12668-0, and 2011/21545-0), Helmholtz Association to Karlsruhe Institute of Technology (KIT)/Helmholtz Institute Ulm (HIU), and Deutschen Bundesstiftung Umwelt (project 31452/01) for financial support. The authors acknowledge


24

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SGL Carbon AG for kindly providing the SIGRACET® 35bc Gas Diffusion Layer and Imerys for supplying C-NERGY Super C65. The authors also acknowledge Prof. Dr. Luiz Alberto Beraldo de Moraes and Dr. Eduardo José Crevelin for the laboratory infrastructure for the GC-MS.

Supporting Information Available: Image of the airtight glass tube with coin cell holder; galvanostatic discharge/charge of coin cell; methodology to obtain the error from the kinetic model; GC-MS measurements for the electrolytic solution before and after the (dis)-charge process; voltammetric cycle of Pt-nps on glassy carbon in H2SO4 0.5 mol L-1. This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Author *E-mail: [email protected]. Telephone: +55 16 3315 4862 Notes The authors declare no competing for financial interest.


25


Page 26 of 31

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The Role Platinum Nanoparticles Play in the Kinetic Mechanism of

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