Article pubs.acs.org/JPCC
On the Mechanism of Nonaqueous Li−O2 Electrochemistry on C and Its Kinetic Overpotentials: Some Implications for Li−Air Batteries Bryan D. McCloskey,† Rouven Scheffler,‡ Angela Speidel,‡ Girish Girishkumar,† and Alan C. Luntz*,†,§ †
Almaden Research Center, IBM Research, 650 Harry Road, San Jose, California 95120, United States Volkswagen Group, Inc., Belmont, California 94002, United States § SUNCAT, SLAC National Accelerator Laboratory, Menlo Park, California 94025-7015, United States ‡
ABSTRACT: Quantitative differential electrochemical mass spectrometry and cyclic voltammetry have been combined to probe possible mechanisms and the kinetic overpotentials, responsible for discharge and charge in a Li−O2 battery, using C as the cathode and an electrolyte based on dimethoxyethane as the solvent. Previous spectroscopy experiments (X-ray diffraction, μRaman, IR, XPS) have shown that Li2O2 is the principle product formed during Li−O2 discharge using this electrolyte/cathode combination. At all discharge potentials and charge potentials 4.0 V, the electrochemistry requires significantly more than 2e−/O2, and we take this as evidence for electrolyte decomposition. We find that sequential concerted (Li+ + e−) ion transfers to/from adsorbed O2* and LiO2* to produce/consume Li2O2 is the mechanism that is most compatible with these experiments. The kinetic overpotentials are extremely low relative to aqueous O2 reduction and evolution, and this implies that in principle a discharge−charge Li−O2 cycle is possible with high voltaic efficiency (∼85%) if electrolyte and cathode stability issues are resolved.
I. INTRODUCTION Over the past several years, there has been much active research into nonaqueous Li−air (or Li−O2) batteries with the hope that successful Li−air battery development would give a safe and cost-effective secondary battery with a much higher specific energy than Li−ion batteries. The net electrochemical reaction in a nonaqueous Li−air battery is 2Li + O2⇄ Li2O2, with battery discharge described by the forward direction and charge described by the reverse direction. The possible large increase in specific energy arises from two potential sources: (1) one of the reactants, O2, is not stored in the battery but comes from breathing air as in a fuel cell and (2) the use of Li metal as the anode rather than intercalated graphite (LiC6). Despite their great promise, there are significant challenges to developing practical Li−air batteries.1,2 However, even if all these practical challenges can be successfully overcome, the Li−air battery will ultimately be limited by the fundamental electrochemistry itself, and this depends in large part upon the mechanism of the electrochemical reaction and its kinetic overpotentials, η. Therefore, it is essential to understand these aspects of the electrochemistry in detail. In this paper, we present some new experimental evidence on the mechanism of the reaction and the magnitude of the kinetic overpotentials. There have now been many galvanostatic discharge−charge studies of nonaqueous Li−O2 cells at low current densities typically using SwagelokTM or coin cell type designs. Recent experiments, which combined discharge product analysis and © 2012 American Chemical Society
differential electrochemical mass spectrometry (DEMS) to analyze gases evolved during galvanostatic discharge and charge, demonstrate that even the basic electrochemical reaction depends critically upon the choice of solvent.3−6 For example, when using a significant component of organic carbonates (typical Li−ion solvents), solid carbonate electrodeposits are formed as a result of solvent decomposition during discharge, and CO2 is then evolved from these deposits on charge. This irreversible chemistry led to erroneous suggestions of a ∼1.5 V charging overpotential, and this resulted in initial7 and continuing8 emphasis on electrocatalysis in Li-air battery research. On the other hand, when using organic ethers such as dimethoxyethane (DME) as the electrolyte solvent, Li2O2 is the predominant discharge product.3,4 However, comparing oxygen evolution during charge (OER) to oxygen consumption during discharge (ORR) shows that the electrochemical reversibility (OER/ORR) is only ∼80% for DME3 and less for many other solvents9 so that rechargeability is extremely limited. Thus, while these SwagelokTM or coin type cells are good model systems for practical batteries and for studying the electrochemical stability of the battery components, they do not allow many aspects of the fundamental electrochemistry to be determined since their behavior is a combination of electroReceived: July 5, 2012 Revised: October 16, 2012 Published: October 25, 2012 23897
dx.doi.org/10.1021/jp306680f | J. Phys. Chem. C 2012, 116, 23897−23905
The Journal of Physical Chemistry C
Article
Figure 1. Schematic diagrams of the hermetically sealed bulk electrolysis cell and the DEMS cell.
(1.1 cm2, FMC Corp.) as the counter and reference electrodes. A large volume (∼15 mL) of 1 M LiTFSI or 0.5 M N(C4H9)4N(CF3SO2)2 [NBu4TFSI] in DME is used as the electrolyte and is vigorously stirred to eliminate O2 and ion transport effects. iR compensation is used for all CV. Two gas ports are built into the cell lid, and ∼1 bar O2 (research purity, Matheson Gas) is bubbled through the electrolyte. The cell is assembled in an argon glovebox, and special care is taken to ensure that the cell contents are never exposed to ambient atmosphere. For measurements of the O2 pressure dependence of the CV, the gas flow is composed of a calibrated O2/Ar mixture using MKS flow controllers to achieve O2 partial pressures less than 1 bar. To achieve O2 pressures greater than 1 bar, the cell is simply pressurized with pure O2 instead of using continuous flow (i.e., static O2 atmosphere). The DEMS cell, which resembles a SwagelokTM type cell, is built to provide high hermetic integrity. The active cell components include a 1.1 cm diameter Li metal anode, 2 × 1.25 cm diameter Celgard 2500 separators, and a 1.2 cm diameter carbon cathode. Either AvCarb P50 carbon paper or a stainless steel mesh coated with XC72 carbon particles using PTFE as a binder is the cathode. No significant differences in Li−O2 discharge or charge chemistry (i.e., e−/O2 during discharge and charge, respectively) have been observed for different types of carbon cathodes.11Approximately 60 μL of 1 M LiBF4 or LiTFSI in DME is used as the electrolyte. These components are stacked between a SS alloy 20 anode and cathode tip that have been hermetically sealed against a quartz tube using compressed Markez O-rings (Marco Rubber). Gases are fed and swept away from the cathode head space via two capillaries that have been silver soldered into the cathode tip. Evolved gases are sent to a calibrated differentially pumped Residual Gas Analyzer (Stanford Research Systems) for quantitative identification using a purge of the cell head space (∼1.5 mL) into a much larger evacuated volume using a high pressure pulsed gas valve (HPGC), with O2 gas as the purge for discharge and Ar for charge. During charge, the quantitative composition of the head space gas swept out of the cell is analyzed using the calibrated residual gas analyzer, with the
chemistry, chemistry related to electrolyte (and cathode) stability and reactant transport issues, and the latter two are dependent upon detailed cathode and cell design. In this paper, we present detailed studies of Li−O 2 electrochemistry on carbon using a single solvent: a lithium salt, LiN(CF3SO2)2 (LiTFSI) or LiBF4, dissolved in DME. This electrolyte system is the one that we have found by DEMS to be the most stable so far for Li−O2 discharge−charge chemistry and that is also compatible with a Li metal anode.3,9 We use a hermetically sealed conventional bulk electrolysis cell for cyclic voltammetry (CV) experiments and compare these results to similar experiments in a modified SwagelokTM type cell in order to separate the fundamental electrochemistry from electrolyte stability and cell dependent properties. In the bulk electrolysis cell, we use small surface area glassy carbon (GC) cathodes as the working electrode. In the modified SwagelokTM type cell, quantitative measurements of pressure changes (ΔP) and differential electrochemical mass spectrometry (DEMS) allow a quantitative probe of gases consumed/evolved during cell discharge/charge and their relationship to the coulometry, e.g., if/when species other than O2 are produced and the number of O2 molecules consumed or produced per electron. This information is essential in order to properly relate the currents observed in the bulk electrolysis cell to true electrochemical processes involving the Li−O2 couple. We emphasize here only aspects of the electrochemistry that relate to the mechanism of the reaction in this electrolyte on C and the related kinetic overpotentials.
II. EXPERIMENTAL SECTION Basic experimental procedures have been described in detail in previous studies3,10 and are only briefly described here. Figure 1 presents schematics of both the bulk electrolysis and DEMS cells employed in this study. A Biologic VMP3 Workstation is used for all electrochemical characterization, and all potentials quoted are relative to Li/Li+. The hermetically sealed bulk electrolysis cell uses a flat nonporous polished glassy carbon (GC, 1.1 cm2, Tokai Carbon, U.S.) as the working electrode and a high purity Li metal foil 23898
dx.doi.org/10.1021/jp306680f | J. Phys. Chem. C 2012, 116, 23897−23905
The Journal of Physical Chemistry C
Article
absolute quantities of the different gas components determined by comparing the calibrated mass spectrometer intensity for the various masses to the 36Ar peak of known Ar head space pressure and volume. Alternatively, the carefully calibrated ∼1.5 mL cell head space volume can be isolated, and pressure decay/ rise (ΔP) during cell discharge/charge can be monitored using an in-line pressure transducer (Omega PX419, 30 psia range, 0.08% accuracy). These ΔP measurements allow accurate quantification of the number of electrons per O2 consumed in discharge or produced during charging (e−/O2) when corrections are made for the small pressure rises due to gas evolution from parasitic reactions. During charging, the e−/O2 ratio obtained from the calibrated mass spectrometer measurement agrees within 3% of that from the pressure rise. Further technical details of the DEMS setup are available offline by contacting the authors. Although O2 also comes into contact with the Li metal in both the DEMS and electrolysis cells, we have not found this to contribute in any way to the observed electrochemistry, i.e., reference potential or gas consumption/evolution. For example, the electrochemical impedance in the DEMS cells increased modestly over a 24 h period under open circuit conditions, but this increase was nearly independent of the presence of ∼1 bar O2 pressure in the cell head space relative to a 1 bar Ar head space. In addition, quantitative gas evolution during cell charge is nearly identical under both Ar and O2 atmospheres, indicating that O2 evolved at the cathode does not react with the Li electrode. Both results suggest that Li metal electrodes are “protected” from extensive chemical reaction with O2 by a solid electrolyte interface, plus perhaps an extra LiOx layer. Thus, the Li remains a reasonable reference potential in the presence of O2, and the electrochemistry probed in both types of experiments is dominated by only the cathode Li−O2 chemistry. All solvents and salts used in this study were purchased from Novolyte (Purolyte electrolyte grade), stored in an argon glovebox (
