Nonaqueous LiâAir Batteries: A Status Report - Chemical Reviews

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Nonaqueous Li−Air Batteries: A Status Report Alan C. Luntz*,† and Bryan D. McCloskey‡,§ †

SUNCAT, SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States Department of Chemical and Biomolecular Engineering, University of California, Berkeley, California 94720, United States § Environmental Energy Technologies Division, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States 1. INTRODUCTION ‡

It is well-recognized that electrification of transportation would do much to minimize our consumption of fossil fuels, and thereby reduce CO2 emissions and consequent effects on climate change, as well as provide better energy security and lessen the economic drain of fossil fuel imports. As a result, a transition to partial road electrification is now occurring. Hybrid electric vehicles (HEV) are now common, plug in hybrids (PHEV) are now beginning to sell appreciably (∼50 000 in the United States in 2013), affordable electric vehicles (EV) with a limited driving range of ∼50−100 miles are also selling at a modest pace (>40 000 in the United States in 2013), and even a luxury EV (Tesla) with longer range has sold ∼20 000 in 2013. The major issue confronting complete electrification of road transport is simply a battery problem, i.e., developing a safe, long-lived, and cost-effective battery with sufficient specific energy and energy density (along with sufficient power density) to extend the driving range to cover most daily use. Current EVs are based on Li−ion batteries, and massive research and development has been and is still being devoted to optimize current generations of Li−ion batteries, as well as in developing so-called advanced Li−ion batteries (Si or Li metal anodes, high capacity cathodes). However, over the past ∼5 years, an opinion has been evolving by some in the battery community that even advanced Li−ion may not be a sufficient battery for mass market adoption of EVs, defined here as affordable and safe driving of a midsize car with a range >300 miles. This had led to fundamental research activities into entirely different kinds of battery chemistries, often referred to as “beyond Li−ion” (BLI). While these represent several totally different chemistries, they are unified in that each has a theoretical specific energy and/or energy density surpassing that of current Li−ion. While both metrics are undoubtedly important, which of the two is the most important for EV applications is somewhat debated, even among the different EV manufacturers. Traditional car companies emphasize more the importance of energy density (fitting the battery into the trunk, etc.), while Tesla emphasizes more the specific energy since they tend to design a car around the battery pack. Of all BLI chemistries, Li−air has the highest theoretical specific energy, and hence, this battery chemistry has attracted enormous research attention in the past couple of years. For example, according to the Web of Science (WOS), only 11 papers and patents were published in 2009 on Li−air or Li−O2

CONTENTS 1. Introduction 2. Galvanostatic Discharge−Charge of a Li−O2 Battery 3. Fundamental (Ideal) Li−O2 Electrochemistry 3.1. Proposed Mechanisms 3.2. Morphology of Li2O2 Deposits 3.3. Experimental Studies of Li−O2 Electrochemistry and Overpotentials 3.3.1. Kinetic Overpotentials 3.4. Theoretical Studies of Li−O2 Electrochemistry and Overpotentials 4. Potential Increase During Charging 4.1. Electrocatalysis 5. Parasitic Chemistry 5.1. Quantitative Definitions of Li−O2 Rechargeability 5.2. Electrolyte Stability 5.3. Cathode Stability 5.3.1. Non-Carbon Cathodes 5.4. Stability to Air Contaminants 6. Capacity Limitations and Charge Transport through Li2O2 7. Li Anode 7.1. Polymer Electrolytes 7.2. Solid State Electrolytes 8. System Limitations for an EV Battery 9. Summary and Outlook Author Information Corresponding Author Notes Biographies Acknowledgments References

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Special Issue: 2014 Batteries Received: January 29, 2014 Published: November 7, 2014

© 2014 American Chemical Society

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their own discretion when any claims are made based solely on battery cycling (often with only very short depth of discharge cycles) without any quantitative chemical backup. It is now well-known by researchers in the field that there are important technical challenges to developing any practical nonaqueous Li−air battery;2 the high charging potential that limits discharge−charge cycle efficiency, the low discharge capacities at higher currents that limit power density, the poor electrolyte and cathode stability that dramatically limit cycle life, and the poor Coulombic efficiency/safety issues in using Li metal. In addition, parasitic reactions from components in air other than O2 (H2O, CO2) imply the necessity either to scrub the air or to utilize pure O2. This implies a significant system level cost to the specific energy and energy density in a practical Li−air battery. Because air scrubbing and/or using an O2 tank is a system related issue and most laboratory studies to date are based on laboratory Li−O2 batteries, we will most often use the term Li−O2 instead of Li−air. In this review, we try to evaluate the literature to date as it relates to the key technical and system related challenges in developing a successful Li−air battery for EV use. We discuss the basic behavior and descriptors of the galvanostatic discharge−charge of a Li−O2 battery in section 2, the fundamental (ideal) Li−O2 electrochemistry in section 3, the high charging potential in section 4, the parasitic chemistry and electrochemistry in section 5, charge transport through Li2O2 and its implications for capacity limitations in section 6, prospects for using Li metal in section 7, and system limitations in section 8, and we provide a summary and outlook in section 9. All potentials quoted in this review are relative to the Li/Li+ potential. While this review is not based on a historical perspective, it is important to acknowledge the pioneers that initiated the field and stimulated its current high interest. Jiang and Abraham discovered the nonaqueous Li−air battery in 1996 serendipitously as they were attempting to study Li intercalation into graphite with a gel polymer electrolyte and had some air leaks.3,4 Recognizing that the effects they were seeing were due to Li−O2 electrochemistry, they designed experiments to specifically probe this and identified Li2O2 as the dominant crystalline product via Raman spectroscopy. Shortly thereafter various authors studied discharge capacities of Li−O2 batteries with different electrolytes.5 Interest in Li−air batteries increased substantially in ∼2007 when Bruce and collaborators showed that a Li2O2 packed electrode evolved O2 during charging, proving that Li−O2 could in principle form a rechargeable battery couple.6 They subsequently claimed that electrocatalysis could reduce the significant charging overpotential and increase cyclability.7 As will be discussed in detail later, this last conclusion was incorrect because the experiments used carbonate solvents that are not compatible with Li−O2 electrochemistry.8 In fact, most experiments by all authors prior to 2010 used carbonate solvents since these were wellknown from Li−ion batteries and were compatible with Li metal. IBM and several Department of Energy Laboratories established exploratory research programs in Li−air in 2009− 2010, and the field has grown roughly exponentially since then to its present state of hundreds of research participants.

batteries. However, the WOS indicates >300 papers and patents were published in 2013 with Li−air or Li−O2 in the title. Even a book dedicated to Li−air appeared in 2014.1 There are two varieties of Li−air batteries, a nonaqueous version using aprotic electrolytes and one using an aqueous electrolyte. The ideal reactions in the nonaqueous Li−air battery are simply electrochemical 2Li + O2 ⇆ Li2O2, with the forward direction describing discharge and backward direction for charge. For the aqueous Li−air (in an alkaline electrolyte), the fundamental electrochemical reactions are 4Li + 6H2O + O2 ⇆ 4(LiOH· H2O), again with the forward direction describing discharge and the backward direction for charge. In this review, we report a summary of work done to date (December 2013) on nonaqueous Li−air and Li−O2, and our own subjective evaluation of this work and its implications for the prospects of nonaqueous Li−air to fulfill the promise of becoming the EV battery of the future. There is also research activity on the aqueous version of Li−air, although this is significantly less than that for nonaqueous Li−air because of its lower theoretical specific energy and energy density. This research is not covered in this review. The theoretical specific energy (Wh/kg) and energy density (Wh/L) of a battery are defined simply by the weight and volume of the active components. In the fully charged state of Li−air, these correspond to those of Li metal (assuming a metallic anode and that the O2 comes from air breathing and is not carried in the battery), 11 400 Wh/kg and 6080 Wh/L, respectively. In the fully discharged state these correspond to those appropriate for Li2O2, 3458 Wh/kg and 3445 Wh/L. Neither limit is really meaningful for a battery that continuously gains weight during discharge, but the fully discharged state is certainly more appropriate if the O2 is carried on board. No practical nonaqueous Li−air battery exists yet, so it is impossible to state specific energies or densities that are actually achievable. However, we discuss in section 8 that practical energy densities will be lowered substantially from the theoretical ones due to anticipated system losses, and this may ultimately temper some of the current hyperbole regarding Li− air. The history of rechargeable nonaqueous Li−air batteries at this stage is so short that the field must be considered a work in progress. In fact, even the basic mechanisms and rationale for many of the fundamental properties of Li−air are still in dispute among many of the researchers in the field. One particularly difficult issue is that the Li−O2 electrochemistry (and chemistry) is extremely sensitive to electrolyte and gas impurities, e.g., H2O, so that authors for similar experiments often report different results and conclusions. We will discuss many of these disputed issues, and give our own evaluation for support or disagreement of the various proposals. Aspects that we will not highlight in this review, however, are conclusions based solely on repeated discharge−charge cycles without any quantitative chemical evidence as to what electrochemistry is occurring during the cycles. Because considerable parasitic chemistry currently occurs during the Li−O2 electrochemistry, even qualitative measurements (e.g., X-ray diffraction) of the formation and dissolution of some unspecified amount of Li2O2 are not really sufficient to evaluate most of the conclusions suggested. Since there have been a great many of such publications in the past few years, especially related to cyclability or electrocatalysis, we do not feel it appropriate or even possible in this review to offer opinions on all of these claims. We just caution all interested in the field to exercise

2. GALVANOSTATIC DISCHARGE−CHARGE OF A LI−O2 BATTERY In this section, we outline the key parameters that define galvanostatic discharge and charge of a Li−O2 battery. A typical laboratory Swagelok type Li−O2 battery used at IBM is shown 11722


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current density j can be inferred from the current i. Active stirring minimizes Li+ and O2 transport limitations in this cell. However, this cell only measures the dependence of the total electrochemical current density j on potential U or vice versa, without identification of the part of the current density due to the desired Li−O2 electrochemistry. To distinguish the two types of electrochemistry experiments, we will refer to the Swagelok and similar cells (porous cathode, two electrodes) as a Li−O2 battery and the bulk electrolysis cell (small area cathode, three electrodes) as a Li−O2 electrolysis cell. Figure 2a shows a typical single galvanostatic discharge− charge cycle in the Swagelok battery at a current of 0.23 mA/

in Figure 1a. This consists simply of clamping together a Li metal foil anode, an electrolyte soaked separator, and a porous

Figure 2. (a) Galvanostatic Li−O2 discharge−charge cycle at 200 mA/ gC in a Swagelok battery with 1 M LiTFSI in DME as the electrolyte and a Vulcan XC72 carbon black cathode. (b) Galvanostatic Li−O2 discharge−charge at 100 μA/cm2 current density on GC in an electrolysis cell (same electrolyte as in part a). Reprinted with permission from ref 10. Copyright 2012 American Chemical Society.

Figure 1. (a) Typical Swagelok type battery. MS refers to differentially pumped mass spectrometer used for DEMS. (b) Typical bulk electrolysis cell to study Li−O2 electrochemistry. Reprinted with permission from ref 9. Copyright 2012 American Chemical Society.

C cathode of ∼0.1 m2 total surface area (in ∼1 cm2 projected area) and a stainless steel mesh or other current collector. Pure O2 gas is used instead of air to avoid deleterious contamination issues (see sections 3.2 and 5.4). Extreme care is taken to both clean and dry all cell components and to ensure high hermetic integrity throughout the experiments. The only aspect of this cell that is unique compared to many test cells in the field is that it is connected to a differentially pumped mass spectrometer (MS) for quantitative analysis of the O 2 consumed in discharge (ORR) and the O2 and other gases evolved during charge (OER) via the technique called differential electrochemical mass spectrometry (DEMS). Thus, in addition to coulometry measurements, the detailed chemistry can be simultaneously measured quantitatively by the DEMS. Figure 1b shows a typical bulk electrolysis cell used to follow Li−O2 electrochemistry on a glassy carbon (GC) working electrode of ∼1 cm2 total area. This cell has three electrodes instead of only two as in the Swagelok battery so that all potentials can be related to that of a reference electrode at minimal current and resistive losses. Since the active area of the working electrode is approximately known, the electrochemical

cm2 (projected area) or 200 mA/gCarbon, i.e., potential U versus capacity Q, with a typical “good” electrolyte (1 M LiN(CF3SO2)2 [LiTFSI] in dimethoxyethane [DME]). There are several features that characterize this cycle: the equilibrium potential for the Li−O2 electrochemistry, U0, a potential less than U0 during discharge, Udis, a sudden death to the cell characterized by a sudden drop in Udis at some maximum discharge capacity Qmax, a complicated potential dependence during charging, Uchg, with an initial small overpotential followed by a continuous or sometimes step increase in the charging potential. All features except U0 have significant dependences on current i. An equivalent galvanostatic discharge−charge cycle in the electrolysis cell is shown in Figure 2b. This allows a separation of the battery’s structurally dependent properties from those that are inherent to the electrochemistry. The qualitative features in Figure 2b are similar to those of Figure 2a, implying that all qualitative features in Figure 2a are inherent to the cathode electrochemistry rather than some battery specific features. There are important quantitative differences, however. Because of the much smaller surface area of the GC relative to the active 11723


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cathode area in the Swagelok cells, Qmax is ∼103 smaller, and the cathode current density j (defined as the current normalized to the total electrochemically active surface area in the cathode) is ∼103 higher. One key point is that, in both configurations, the galvanostatic discharge−charge cycle does not reveal any information about the identity of the electrochemistry; it simply measures the sum of all electrochemical processes. There is no guarantee that this cycle is dominated by the electrochemical reactions 2Li + O2 ⇆ Li2O2, although for the moment we assume this is true and discuss limitations in this assumption later. The thermodynamic standard potential for formation of bulk Li2O2 is 2.96 V, and this is often stated as U0. However, the open circuit voltage (OCV) in both the Swagelok and electrolysis cells is ∼3.1 V (vs Li/Li+) prior to discharge, but ∼2.85 V after discharge to sudden death. We take these OCV values to be U0 on these different surfaces since the equilibrium potential is for the surface reaction of 2Li +O2 = Li2O2 on whatever surface is present, and this need not be same as the standard potential for bulk Li2O2 formation. For example, the equilibrium potential for the reaction at a C surface (initial OCV) may be different than that on a Li2O2 surface (after sudden death). The latter is more appropriate for most discussions since most of the electrochemistry during discharge/charge involves growth/dissolution of Li2O2 on Li2O2 (see section 3). A high Qmax given the active component weights is of course the main reason for interest in Li−air batteries, so naturally there have been many studies of the dependence of Qmax on the type or design of the C cathode, especially at low i where Qmax is large. Many authors have attributed this dependence of Qmax to various microstructure of the C cathodes, e.g., porosity, mesoporosity, etc.11 However, the most convincing argument is simply that this dependence is proportional to the total surface area of the C cathode12 (and hence approximately on the carbon weight g2/3 assuming a wide distribution of carbon aggregate particle sizes). Given a C cathode, Qmax is dominated by electrical passivation of the cathode and is discussed extensively in section 6. In galvanostatic discharge, both Qmax and Udis are strong functions of i as shown in Figure 3a. This is particularly troublesome since this implies a poor energy-power trade-off for Li−O2 batteries, and may make it difficult to take advantage of the potentially high specific energy of Li−air batteries at reasonable power densities. Many other authors have reported similar current dependences in galvanostatic discharges.5a,13 Figure 3b shows related discharge experiments in the bulk electrolysis cell on a GC cathode. While there are some differences from Figure 3a, the similarities imply that the i dependence in both Qmax and Udis is inherent to the electrochemistry. We show in section 3 that the dependence of Udis on i is not quite the same in the two types of cells, but that the dependence of Qmax both in Li−O2 batteries and in the bulk electrolysis cell are well-described by a simple charge transport model discussed in section 6. The complicated dependence of Uchg on the charging capacity Qchg following galvanostatic discharge is one of the most important (and controversial) aspects of Li−O2 batteries. The high potential rise occurring during charge underlies many important limitations in current Li−O2 batteries: a poor discharge−charge cycle electrical efficiency, limitations in electrolyte electrochemical stability, and poor cyclability. Although Uchg originates from the cathode electrochemistry

Figure 3. (a) Galvanostatic discharges in a Swagelok battery as a function of current, with XC72 cathode and 1 M LiTFSI/DME electrolyte. Udis is the average plateau discharge potential at a given current, and Qmax is the capacity to a discharge voltage of 2 V. Q is normalized to the weight of C in the cathode since this is typical in the field. (b) Galvanostatic discharges on GC in the electrolysis cell using the same electrolyte as in part a at the current densities in the legend. Reprinted with permission from ref 16. Copyright 2013 American Chemical Society.

rather than some battery specific design aspects, the exact details of how fast Uchg rises during charging in Swagelok or equivalent batteries in the literature seems to depend on many factors: discharge current and morphology,14 the purity of all components in the battery,15 the type of C in the cathode,10 etc. We defer discussion of the details of the properties of Uchg and proposals for its rise during charging until section 4. Early experiments using carbonate solvents in the electrolytes observed that Uchg rose almost immediately to >4 V. This was incorrectly interpreted as a high Li−O2 kinetic overpotential for charging and has led to a strong focus on electrocatalysis in the literature. However, as will be pointed out in sections 4 and 5, this rapid rise was instead related to electrolyte decomposition that formed carbonate and carboxylate products rather than Li2O2 during discharge. Once formed, these do not electrochemically oxidize until ∼4 V. More stable (and pure) ether electrolytes that produce principally Li2O2 upon discharge usually give an initial low Uchg that rises with Qchg in a manner similar to Figure 2.

3. FUNDAMENTAL (IDEAL) LI−O2 ELECTROCHEMISTRY Before discussing the technical limitations in Li−O2 batteries, it is essential to first discuss the ideal Li−O2 electrochemistry since a Li−O2 battery can never overcome the inherent 11724


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transport limitations to Qmax, and this will be discussed in section 6. As eqs 1−4 are written above, they all involve inner sphere charge transfer processes for species adsorbed on the surface; i.e., there are no solution electrochemical or chemical processes occurring. There is certainly hard NMR and other experimental evidence that Li2O2 is completely insoluble in most nonaqueous electrolytes (Ksp < 10−10), but there is no evidence that we are aware of regarding the solubility of LiO2 in nonaqueous electrolytes. Some authors13c argue that a LiO2 solution process involving disproportionation in solution followed by precipitation (eq 4a) is necessary to explain the large toroids of discharge products observed by them and many others.

limitation imposed by the ideal Li−O2 electrochemistry itself, i.e., of the mechanism for the reaction and its kinetic overpotentials. 3.1. Proposed Mechanisms

Because of the wide range of experimental observations in Li− O2 batteries, details of the discharge and charge mechanisms are somewhat contentious. In this section, we present proposed mechanisms, and in the following two sections we discuss experimental and theoretical evidence related to these proposals. Most authors9,17 argue that the dominant electrochemical reactions are given by the surface electrochemistry: 2Li ⇆ 2(Li+ + e−) +

−

+

−

(anode)

(1)

Li + e + O2 * ⇆ LiO2 *(cathode) Li + e + LiO2 * ⇆ Li 2O2 *

(cathode)

LiO2 * + LiO2 * ⇆ Li 2O2 * + O2

(cathode)

LiO2 + LiO2 → Li 2O2 + O2 → Li 2O2 * + O2

(2)

We discuss in section 3.2 that the formation of toroids (in ether-based electrolytes) is likely related to H2O contamination in the battery and that without H2O the Li2O2 seems to only form in thin conformal layers on the surface. With H2O contamination, however, reaction 4a or another solution process is possible. Other authors20 suggest that reaction 2 is actually composed of two time separated solution steps (eqs 2a and 2b) instead of the concerted ion−electron transfer surface mechanism implied in eq 2.

(3) (4)

Here the * refers to a surface adsorbed species. The forward direction of all arrows describes discharge, and the reverse arrows describe charge. For galvanostatic conditions appropriate for battery discharge, steady state must be achieved for the intermediate LiO2* since Li2O2 is the ultimate product of the discharge. This implies that d[LiO2*]/dt = 0, and thus, the discharge current at constant overpotential i ∝ k2[Li+*][O2*]. [Li+*] and [O2*] are the surface concentrations of the two species. This same equation for i occurs when either eq 3 or 4 finally produces Li2O2 following eq 2 so that no kinetic (or current) measurements can ever determine whether eq 3 or 4 dominates discharge. Which dominates depends on whether the chemical barrier to disproportionation and O2 evolution of eq 4 is larger or smaller than the thermodynamic enthalpy change for eq 3 at a given U plus the Li+ desolvation barrier that defines the electrochemical rate of reaction 3. Although eqs 3 and 4 have been discussed as totally separate mechanisms for the final step of Li2O2 formation, they are in reality quite similar. The only difference is whether the final Li+ addition comes from a Li+ adsorbed on the surface as LiO2* or from solution, and these are similar since the O2 adsorbed on the surface is in equilibrium with that in the solution and continuously adsorbing/desorbing. Since eq 4 is a purely chemical process, if it is exothermic during discharge, it must be endothermic during charge and therefore unlikely. Therefore, it has been suggested17a,b,c that a simultaneous 2e− direct charge reaction occurs bypassing the LiO2* intermediate as in eq 5. 2(Li+ + e−) + O2 ← Li 2O2

(4a)

O2 + e− → O2−

(2a)

O2− + Li+ → LiO2 → LiO2 *

(2b)

In this case the formation of LiO2* is viewed as a precipitation of insoluble LiO2 formed in solution. However, to our knowledge, there is no experimental evidence that O2− exists even transiently in solution in the presence of high Li+ concentrations (it is of course well-known to be a long-lived species in many stable solvents in the absence of Li+). We outline in section 3.3 experimental evidence against reaction 2 being time separated. In order to rationalize the complicated Uchg behavior, Lu and Shao-Horn21 have proposed that charging consists of two different processes: one with low overpotentials at the surface followed by a bulk oxidation process at higher overpotentials. The low overpotential process is described as delithiation, reaction 3, followed by disproportionation, reaction 4. For the higher overpotentials, reaction 5 is suggested to dominate as a bulk oxidation reaction. In later work,14,22 it was suggested that the low overpotential process is simply due to delithiation, reaction 3, and that a higher overpotential region arises from a two-phase bulk oxidation as in reaction 6, presumably coupled with disproportionation (reaction 4) to evolve O2.

(5)

−

However, this 2e simultaneous direct charge reaction seems unlikely in electrochemistry since two electron processes have much higher entropic barriers than 1e− ones, and this is the basis for the sequential 1e− transfers in electrocatalysis.18 The mechanism represented by eqs 1−3 is the Li+ analogue of the conventional proton coupled charge transfer in the aqueous associative oxygen reduction reaction/oxygen evolution reaction (ORR/OER) involving H2O2.19 The fundamental difference to the aqueous case, however, is that the product species (Li2O2*) is formed on the surface by the ion transfers and remains there permanently since it has absolutely no solubility in most nonaqueous electrolytes, e.g., carbonates, ethers, etc. The buildup of this insoluble (and electrically insulating) Li2O2 has important implications for charge

2e− + 2Li+ + 2LiO2 ← 2Li 2O2

(6)

This suggestion is predicated on the assumption that regions in the deposit exist during charging with phases of pure or high concentrations of LiO2. These mechanisms are discussed in section 4. Still other authors13a,20b have suggested that net electrochemical reactions such as 4(Li+ + e−) + O2 → 2Li 2O*

(7)

2(Li+ + e−) + Li 2O2 * → 2Li 2O*

(8)

also contribute during deep discharge or with reactive metal catalysts. To our knowledge, Li2O has never been observed by 11725


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The apparent resolution of this experimental conflict is that the formation of toroids seems related to H2O content in the cell.15 No toroids were observed in the most anhydrous cell possible at any current, while the size of the toroids and their layering increased with the amount of H2O at a fixed i.15 In addition, at a fixed H2O content, the size of toroids/particles decreased with i.15 At sufficiently high i, no toroids/particles were observable even at high water contents. This current dependence of the toroid formation is consistent with that observed previously.13c,24a As the H2O content increases, the fall off in toroid size occurred at higher and higher i.15 Our experience has been that most of the residual H2O in the battery is not due to that introduced as an electrolyte impurity (which is easily checked by Karl Fischer titrations), but rather from imperfect drying of the porous C cathode, separator, or in the O2 gas. Unless specified, all experiments reported here that are attributed to the IBM group are based on nearly anhydrous electrolytes (10−30 ppm of H2O) and carful drying and gas handling techniques that minimize H2O contamination.10 We cannot comment on the H2O content of other work in the literature, but suggest that the formation of toroids in etherbased electrolytes at a given i is a reasonable indication of the H2O content present. The mechanism of how H2O induces toroids is likely some solution-mediated process as suggested by others,13c given the large dimensions of toroids formed and the charge transport limitations inherent in Li2O2 (section 6). Reference 15 suggests that H2O induces some solubility of LiO2* because of its high Gutmann acceptor number and that this enables a solution mechanism for Li2O2 formation in addition to the one present from the surface electrochemistry. However, evidence was also presented in ref 15 that, in a strictly anhydrous nonaqueous electrolyte where no toroids are formed, there is no apparent difference in morphology and electrochemistry between that occurring in the electrolysis cell on GC and that occurring on a porous C cathode in the Swagelok battery. There will of course be quantitative differences in some of the parameters describing the electrochemistry (Qmax, Udis, Uchg, etc.) due to the different microscopic current densities and total cathode surface area in the two configurations. However, when H2O is present in the battery, there must be a significant additional Li−O 2 mechanism to account for the toroid formation. As discussed in section 5.4, addition of H2O to the anhydrous electrolyte and formation of toroids causes the discharge capacity to increase with H2O content as larger and larger toroids are formed. However, it also causes a more rapid rise in Uchg, and this causes enhanced stability issues.15 Therefore, H2O contamination has a mixed effect in Li−O2 batteries. The crystallinity of Li2O2 deposited on porous cathodes is easily obtained from the Li2O2 XRD line widths after correcting for instrumental broadening (Debye−Scherer analysis). In agreement with other authors,13c,14 we also found average spherical crystal sizes of ∼15 nm at an i where toroids are not formed (or at all i in anhydrous electrolytes15). This crystallinity was independent of the discharge current (and discharge capacity at a fixed current) in the anhydrous electrolyte so that there was no evidence for a transition to an amorphous Li2O2 as suggested by Adams et al.13c It is possible that the absence of the Li2O2 XRD peaks by Adams et al. at higher currents was related to the lower discharge capacity Qmax at the higher currents (as discussed in section 6). When H2O was present in the electrolyte, the average XRD size of the

any spectroscopic technique as a product of the Li−O2 reaction. This is fortunate since its formation is entirely irreversible as packed electrodes with Li2O do not evolve O2 at charging potentials below the oxidation of the electrolyte.8b,d 3.2. Morphology of Li2O2 Deposits

As many authors have noted, the morphology of Li2O2 formation on the cathode must be intimately connected with the mechanism of discharge. AFM studies of galvanostatic discharges in the electrolysis cell on low surface area GC cathodes show that rough homogeneous films of ∼2−10 nm are produced at sudden death.23 The thickness depends on discharge current and is related to electrical passivation of the electrode at the given current density by the Li2O2 film. This passivation will be discussed in detail in section 6. On various types of porous carbons, layered structured toroids of several hundred nanometers in size have been reported by many authors, especially at low discharge currents.13c,22b,24 For fixed i, the toroid size increases with discharge capacity.22b TEM experiments show that the toroids are highly crystalline with the Li2O2 (0001) facet normal to the axis of the toroid24a as shown in Figure 4.

Figure 4. Top two panels, SEM of Li2O2 toroids under different growth conditions. Bottom right, TEM of individual toroid. Bottom left, electron diffraction from a single toroid. Red and blue dots are the calculated diffraction spots for Li2O2 along (0001). Inset, schematic of toroid. Reprinted with permission from ref 24a. Copyright 2013 American Chemical Society.

This observation of toroids has naturally caused much discussion about alternative mechanisms to reactions 1−4 for discharge/charge in Li−O2 batteries because it is impossible to rationalize these large dimensions with the small length scales associated with electrical passivation of the nanometer scale films produced in the electrolysis cell. The observation of large toroids has been taken by some authors13c as evidence for the solubility of LiO2 followed by solution reaction/precipitation, reaction 4a. At somewhat higher currents, smaller particles (∼20 nm) have also been observed.13c,14 However, we (IBM) never observed these toroids via SEM in our experiments at any current with optimally anhydrous and purified ether electrolytes that are moderately stable to Li2O2 (see section 5.2), i.e., DME or TEGDME. SEM of the cathode following discharge appeared to only give thin conformal coatings on the C aggregate particles. 11726


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crystallites increased significantly with the H2O content to ∼50 nm, paralleling the increase in the toroid size; Gallant et al.14 previously also observed increased crystallinity when toroids were formed. 3.3. Experimental Studies of Li−O2 Electrochemistry and Overpotentials

The simplest experimental technique to probe the Li−O2 electrochemistry is cyclic voltammetry (CV) in a bulk electrolysis cell such as that in Figure 1b, and there have now been many such studies varying the electrolyte, working electrode surface, concentration of electrolyte, O2 pressure, etc.9,17a−c,25 In this experiment, Li2O2 is presumably produced during the O2 reduction part of the cycle, equivalent to discharging a battery, and the Li2O2 then oxidized during the oxidation part of the cycle, equivalent to charging the battery. However, several caveats must be kept in mind in accepting these conclusions based solely on CVs. Most important is that a CV only measures the total electrochemistry occurring in the cell, and it is only an assumption that this current is related directly to the assumed Li−O2 electrochemistry. Therefore, some of the peaks observed may be simply reduction/oxidation of impurities, e.g., an H2O impurity induces new reduction and oxidation peaks. Also, because of the large volume of electrolyte, even minimal concentrations of impurities in the electrolyte can distort the results by chemically reacting with Li2O2 at slow scan speeds. Thus, asymmetry in the oxidation peak relative to that of the reduction peak could be due to impurities rather than reaction of the Li2O2/LiO2 with the electrolyte. Laoire et. al17a,b made the original studies of the Li−O2 electrochemistry in various solvents using CVs on a presumed inert GC working electrode. The authors discussed the results in terms of Pearson’s hard−soft acid−bases, a concept that seems only relevant if the Li−O2 electrochemistry is solutionbased. However, arguments are presented against solutionbased Li−O2 electrochemistry below for strictly anhydrous electrolytes. Laoire et al. observed several peaks in the CVs in the various solvents and assigned these to many of the different electrochemical steps of section 3.1. They argue that one electron redox processes are solely responsible for the current at the smallest |U − U0|, e.g., a solution analogue of reaction 1 in discharge. However, this is incompatible with quantitative DEMS experiments which always give (e−/O2)dis ∼ 2 during discharge (see section 5) . McCloskey et al.9 in an ultraclean and dry electrolyte (by careful drying of all cell components and repurifying electrolyte components) observe only one reduction peak and one oxidation peak on GC in 1 N LiTFSI/ DME (with a long tail extending to 4.5 V). This suggests that some of the extra peaks observed by Laoire et al. may be due to contaminants, e.g., H2O induced redox peaks. Peng et al.17c also observe a single reduction/oxidation peak in CVs on a roughened Au electrode as shown in Figure 5, where it is also compared to O2 reduction and subsequent oxidation in an electrolyte without [Li+]. CVs with higher [Li+] at slower scan speeds show similar CV, but with a slightly sharper oxidation peak and with less separation relative to U0 than in Figure 5.9 Peng et al. also presented surface enhanced Raman spectroscopy (SERS) evidence that on an electrochemically roughened Au surface LiO2 is formed on the surface prior to Li2O2 formation, and use this as evidence for reaction 4 over 3 in discharge. Since they do not see LiO2 formation during charge, they cite this as evidence for reaction 5 during

Figure 5. Cyclic voltammetry at a roughened Au electrode in O2 saturated 0.1 M nBu 4NClO4 in CH 3CN containing various concentrations of LiCLO4 as indicated. The scan rate was 1 V s−1. Reprinted with permission from ref 17c. Copyright 2011 Wiley-VCH Verbg GmbH & Co. KGaA, Weiheim.

charge. While the spectroscopy seems convincing, some caution is appropriate since it is never obvious that SERS measures the dominant species in the bulk or even on the surface, nor that the detailed mechanism is completely identical on a carbon surface as on roughened Au. Figure 5 also illustrates the fact that the onset potential for reduction with even small concentrations of Li+ present is ∼0.25 V higher than the onset potential for reduction in an electrolyte without Li+ that produces O2− in solution. This has been observed in other electrolytes on GC as well.9,17a Thus, it is easier to reduce O2 in the presence of Li+ and the thermodynamic stability of the LiO2* intermediate formed must be greater than that of an O2− solution intermediate since ΔG = −eU. McCloskey et al.9 argue that the comparison in Figure 5 and similar figures in refs 9 and 17a imply that the O2 reduction in Li+ electrolytes must be the concerted process of eqs 2−4 rather than time separated processes as implied by eqs 2a and 2b. If it was the latter, then the O2 reduction in the Li+ electrolyte should occur at the same potentials as without it since eq 2b is purely chemical and cannot contribute to the potential generated by the electrochemistry. The free energy stabilization for the chemical step Li+ + O2− → LiO2* (ΔG2b) must appear as an overpotential. McCloskey et al.9 also studied the dependence of the CVs for a LiTFSI/DME electrolyte and GC on the O2 partial pressure in the headspace above the cell and the Li+ concentration. The O2 pressure dependence of the discharge current i at a fixed potential was well-described when [O2*] is related to the partial pressure of O2 gas P through a conventional Langmuir isotherm, i.e., [O2*] = ((bP)/(1 + bP)), with b = 0.35. Also, i is nearly linear in [Li+] for [Li+] < 0.1 N, but roughly independent of [Li+] above this concentration. These results are in good agreement with the quasi-steady-state kinetics appropriate for the CV as suggested previously for eqs 1−4, i.e., i ∝ k2 [Li+*][O2*]. The most serious issue with inferring mechanistic details from only CVs is that they do not chemically resolve the electrochemistry. However, linear sweep voltammetry combined with DEMS in cells such as those of Figure 1a can resolve the chemistry as shown in Figure 6. 11727


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One important conclusion from Figure 6b is that there is absolutely no potential for which a one electron process dominates the electrochemistry of either discharge or charge within the time resolution of the linear voltammetry experiment (∼10 s). This means that once LiO2* is initially formed in these experiments, it transforms to Li2O2 by either reaction 3 or 4 and therefore cannot be thermodynamically stable or metastable at any potential as suggested by Laoire et al.17a,b and Peng et al.17c However, at very low [Li+], LiO2* could in principle be kinetically metastable if eq 4 dominates over eq 3 since eq 4 is quadratic in LiO2*. 3.3.1. Kinetic Overpotentials. There can be many sources of current dependent polarization losses in Li−O2 batteries: fundamental kinetic overpotentials arising from the surface electrochemistry (ηdis and ηchg), resistive or iR losses, and concentration polarizations at higher currents. The latter could be due to Li+ ion and/or O2 transport to the active cathode site. It is straightforward to show that in typical laboratory Li−O2 batteries with thin cathodes and separators (∼0.1−0.2 mm) these concentration polarizations are minimal at currents ∼10 mA/cm2 (projected area) by relaxation experiments, i.e., stopping the discharge or charge for some period of time at open circuit and then restarting. Since restarting discharge or charge occurs at the same Udis or Uchg as before, this argues that the contributions from these concentration polarizations are minimal. Charge transport limitations through the Li2O2 deposit on the cathode can lead to a concentration polarization for the charge, and this is discussed in detail in section 6. Although Udis is not completely constant during the initial part of discharge (prior to the onset of sudden death) in Li−O2 batteries (e.g., see Figure 3a), a plot of the average Udis versus i is approximately logarithmic for i < 0.1 mA per cm2 (projected area), but linear at i > 0.1 mA per cm2 above16 as shown in Figure 7. This implies that Li−O2 batteries are dominated by

Figure 6. Linear scan voltammetry at 0.5 mV/s using the DEMS cell (XC72 carbon cathode, 1 M LiTFSI in DME). (a) Current is the solid line, and the points are the O2 (and CO2) measured by the DEMS. There was very little gas evolution ( 3.5 V, the e−/O2 ratio exhibits a continuously increasing deviation from ∼2.2. This implies that some electrochemically initiated reaction between the electrolyte and Li2O2 (or other products formed during ORR) occurs from the outset, but increases dramatically at higher charging potentials. O2 evolution continues until ∼4.5 V in the DEMS cell, indicating that some Li2O2 persists on the surface up to these high potentials. This result is similar to that observed in the bulk electrolysis cell CV, where an oxidative scan to 4.55 V is necessary to produce a fully cyclical CV.26 Above ∼3.5 V CO2 is evolved, also indicating some electrolyte decomposition, and at ∼4.8 V, electrolyte oxidation dominates the electrochemistry. The overriding conclusion from this result is that there is considerable parasitic chemistry and electrochemistry occurring, and this will be discussed in detail in section 5, where it is shown to be quite dependent upon the exact electrolyte and cathode surface material. Thus, one must use considerable caution in trying to make conclusions based solely on electrochemistry experiments that only measure i versus U or vice versa.

Figure 7. Measurement of Li−O2 approximate discharge plateaus (Udiss) in Swagelok batteries as a function of galvanostatic current i during discharge in the Swagelok battery with 1 M LiTFSI/DME electrolyte. Cathodes are either P50 C paper or XC72 C with PTFE binder on SS mesh. The projected area of both cathodes is ∼1 cm2. Reprinted with permission from ref 16. Copyright 2013 American Chemical Society.

the fundamental kinetic overpotentials (e.g., Tafel expression) only at very low currents, but dominated by iR losses due to the cell impedance at i > 0.1 mA per cm2, i.e., at nearly all practical battery currents. While similar experiments have not also been done for charging in Li−O2 batteries, one anticipates similar iR losses for charging as well. Therefore, engineering minimal cell impedance is a critical issue, and only experiments in the bulk 11728


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electrolysis cell with a low current reference electrode can measure ηdis and ηchg originating from the fundamental electrochemistry. Figure 3 b (and Figure 23a) shows Li−O2 galvanostatic discharges at various current densities j in the electrolysis cell with LiTFSI/DME as the electrolyte. As discussed in section 3.2, Li2O2 forms a continuous thickening film on the GC surface as a result of discharge. At a given j, the discharge plot shows an initial drop in potential, a linear decrease in potential with Qdis, and a “sudden death” characterized by a rapid decrease in U. The latter two phenomena are discussed in section 6 to result from charge transport limitations through the thickening Li2O2 film. We therefore assign the initial drop in U at Qdis ≈ 0 to the kinetic overpotential for discharge ηdis(j), without distortions due to charge transport limitations in the regime for Li2O2 growth on Li2O2 that dominates the electrochemistry during battery discharge. The complicated rising potential of Uchg as shown in Figure 2 is one of the major issues today in Li−O2 batteries, and its strong variation in different experiments has led to considerable controversy as to its mechanistic origin. The fact that a rising Uchg is seen both in the battery and the electrolysis cell on GC cathodes implies that its origin is in the electrochemistry rather than some battery specific property. Several possibilities for this behavior and their mechanistic assumptions will be deferred until section 4. Galvanostatic charge profiles as a function of j obtained in the bulk electrolysis cell are similar to that in Figure 2b) except that the initial potential depends on j.16 A likely mechanism for the overall potential rise during charging is that it is simply a result of electrolyte decomposition products building up at the Li2O2−electrolyte interface during the charging, as discussed in section 4. Therefore, the kinetic overpotentials for charging are defined in a similar manner to that for discharge; i.e., ηchg(j) is the potential shift relative to U0 at the outset of charging. Combining these measurements and Figure 3b (and other similar measurements) with the definition above for ηdis(j) and ηchg(j) gives the Tafel plot for fundamental kinetic overpotentials in Figure 8a. Note that the Tafel plot is curved, with a slope of ∼120 mV/decade at low currents and substantially higher slopes of 200−300 mV/decade at higher currents. Laoire et al.17b previously obtained Tafel slopes by analysis of the CV that are similar to these, i.e., differing significantly in the low and high current regimes. If only these fundamental kinetic overpotentials were limiting a Li−O2 battery, the discharge− charge cycle electrical efficiency would be 1 − [(ηdis(j) + ηchg(j))/(U0)]. At a current density j ∼ 10 μA/cm2 on the GC electrode, corresponding to a i ∼ 10 mA/cm2(projected area) Li−O2 battery, this corresponds to a cycle efficiency >85%. Thus, the fundamental electrochemistry is excellent at practical Li−O2 battery currents, and Li−O2 can in principle form an excellent quasireversible battery couple. Unfortunately, cell impedances and iR losses reduce the cycle electrical efficiency (Figure 7), and because Uchg rises during the charging to >4 V, the cycle efficiency is in actuality much less than this fundamental limit and unsatisfactory at present.

Figure 8. (a) Experimental Tafel plots for Li−O2 discharge (ORR, blue triangles) and for charging following discharge (OER, red squares). (b) First-principles theoretical prediction of the Tafel plots for Li2O2 crystal growth, the dominant process in discharge (blue triangles), and Li2O2 dissolution, the dominant process in charge (red squares). Adapted with permission from ref 16, Copyright 2013 American Chemical Society, and ref 27, Copyright 2014 Springer Science + Business Media.

= 2.64 eV in ref 17d instead of the experimental value 2.96 eV. In general, DFT calculations of oxides require a correction for the well-known problem of the calculated overbinding of O2. Other authors adjust the O2 electronic energy to fit the calculated value of ΔG to the experimental one.28 Several authors have used first-principles thermodynamics based on DFT to predict the surface energies of the different Li2O2 facets so that the equilibrium structure of the Li2O2 electrochemical deposit (Wulff construction) is obtained.28,29 An example is given in Figure 9a from ref 28b. There are slight differences between the different calculations in the details of the Wulff construction, but most agree that the dominant facets are the O-rich (0001) and the (1−100). Figure 9b shows how the stability of the different facets and their terminations changes as U varies away from U0 during discharge and charge.17d This suggests that the O-rich (0001) and (1−100) are the equilibrium facets during discharge and that the O-rich (0001) and O-rich (1−100) are the equilibrium facets during charge. There is, however, no guarantee that the equilibrium facets are the ones dominant in the electrochemical deposition of Li2O2 or its electrochemical dissolution. Kinetic factors such as anisotropy in charge transport could ultimately determine the facets at the electrolyte interface.30 Some authors have calculated the stability of the different bulk lithium oxides and facets as a function of O2 pressure or O atom chemical potential as a way to probe the differences between discharging and charging conditions.28a,31 However, studying facet/ termination stability as a function of U provides a more direct connection to experiments. There is now substantial experience describing electrochemical mechanisms that involve coupled electron and ion transfer reactions by first-principles thermodynamics.18 For example, this approach has been successful in describing trends in electrocatalysis for the aqueous oxygen reduction (ORR) and

3.4. Theoretical Studies of Li−O2 Electrochemistry and Overpotentials

The (bulk) equilibrium potential U0 is defined by the Nernst equation as U0 = ΔG/2e, with ΔG the formation energy of bulk Li2O2 at standard temperature and pressure (STP). ΔG can be calculated using density functional theory (DFT), and gives U0 11729


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aqueous ORR/OER. This calculation gave modest overpotentials, ηdis ≈ 0.4 V for discharge and ηchg ≈ 0.6 V for charge, and suggested that the ηchg ≈ 1.5 V originally observed when using organic carbonate solvents was incompatible with the proposed Li−O2 electrochemistry.33 Of course, experiments have now definitively shown that carbonate solvents do not produce Li2O2 as the ultimate product during the electrochemistry, and this was the reason for the very high charging overpotential.8b−d This mechanism of Li2O2 crystal growth on Li2O2 has been investigated theoretically in much more detail on the basis of the electrochemistry occurring on the most stable facets and terminations (Figure 9b) and many of the steps and kinks on these facets using the periodic supercell surface model.17d Figure 10 shows the calculated electrochemical thermody-

Figure 9. (a) Calculated Wulff shape for Li2O2. Reprinted with permission from ref 28b. Copyright 2012 American Chemical Society. (b) Variation of different facet energies with potential U away from the calculated U0. The nonstoichiometric O-rich surfaces depend on U while the stoichiometric ones are independent of U. Reprinted with permission from ref 17d. Copyright 2013 AIP Publishing, LLC. Figure 10. Initial deposition of Li2O2 on 0.5 O-rich (0001) and dissolution on 0.5 O-rich (0001) facet of Li2O2. The images show the intermediate states of the cathode surface, where colors highlight atoms that have adsorbed during discharge or atoms that have yet to desorb during charge. Colors: O is dark gray and red; Li is light gray and yellow. Energies are plotted for the total system, where Li and O2 that are not adsorbed are included in solution (Li+ +e−) and O2 as the labels show. In the upper part, the potential-limiting step in forming an island of Li2O2 is the adsorption of Li after two LiO2 have been adsorbed. From there the energy almost does not change when adsorbing the third LiO2 that induces all the adsorbed O2 to “stand up”. In the lower part, after the desorption of one subsurface Li desorbing another is the limiting step. The LiO2 between them then readily desorbs while two surface Li move into the created Li vacancies below. Another LiO2 can desorb, and from here, the charging overpotential is expected to converge to the kink/step analogues as the pit is growing. Reprinted with permission from ref 17d. Copyright 2013 AIP Publishing, LLC.

oxygen evolution (OER).19a,32 This is based on calculating the limiting potentials for all thermodynamic steps to occur downhill in free energy. There is, however, a significant difference in describing Li−O2 electrochemistry since the Li2O2 formed builds up as a deposit on the cathode surface rather than dissolves into solution. Thus, after initial nucleation of Li2O2 on the cathode surface, the discharge is dominated by Li2O2 growth on Li2O2 rather than on the original cathode surface. In a similar manner, most of the charging occurs from the Li2O2 dissolution from Li2O2 surface itself rather than from the cathode surface. The application of first-principles thermodynamics to Li−O2 was originally based on calculating limiting potentials for a mechanism described by eqs 1−3on a stoichiometric step on the (1−100) facet of Li2O2 using a periodic surface supercell model of the Li2O2 surface.33 In addition to the assumption of Li2O2 growth/dissolution from the Li2O2 surface, it was assumed that Li+ + e− is in equilibrium with bulk Li metal, that O2 in the gas headspace is in equilibrium with O2 in solution which is also in equilibrium with O2 adsorbed on the surface, that all charge transfer involves the coupled adsorption/desorption of Li+ + e− at the cathode, and finally that interactions of the surface species with the electrolyte are significantly smaller than interactions with the surface and so are ignored. All except the first are similar to those used for theoretically describing proton coupled electron transfers in

namics occurring on the terrace of the O-rich (0001) facet, i.e., nucleating a new island during discharge and creating a pit during charge based on reactions 1−3. The electrochemical cycle consists of two Li2O2 since there are two Li2O2 per unit cell in this surface facet. As in aqueous studies, the thermodynamic overpotentials are defined as the limiting potentials at which all electrochemical steps in the reaction are downhill. For Figure 10, these correspond to ηdis = 0.33 V and ηchg = 0.36 V. It is likely that there are also additional 11730


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The nucleation regime has been treated by the growth of small Li2O2 clusters on graphite (and defects) or metals such as Au and Pt(111) using theoretical techniques similar to those described previously.34 These demonstrated several significant results relating to this nucleation regime: (1) that Pt likely dissociates Li2O2 during nucleation, (2) that the overpotentials for nucleating growth of Li2O2 on the basal plane of graphite or nanotubes is ∼1.5 V, and (3) that the overpotential for nucleating growth is considerably lower, ∼ 0.4 V, at COx species formed by oxidation at divacancies or ledges of the graphite. These calculations rationalize experimental observations that defects in nanotubes/C(0001) nucleate growth of Li2O2 and that the growth of Li2O2 on Li2O2 seems preferred to nucleating new growth sites.35 The intermediate regime considers Li−O2 electrochemistry on larger clusters/nanoparticles of different sizes.36 These authors argue that the surface of the clusters/nanoparticles for n > 2 are most stable in a high spin state, with the structure rearranging so that a LiO2 is at the outer surface of the cluster. n is the number of formula units in the cluster. Although some caution in this statement is necessary strictly on the basis of DFT results because of inherent DFT limitations to describe triplet states, it was also justified in high level quantum chemical calculations (MP3) for clusters with n < 4. The conclusion and impact that some LiO2 forms at the surface of the cluster/ nanoparticle is similar to the result of the surface science model, i.e., that the 0.5 O-rich (0001) facet (with half a monolayer of LiO2 at the surface) is the most stable surface at electrochemical equilibrium. Presumably, in a fluxional cluster with many low energy conformations one does not need to break stoichiometry for the LiO2 to appear at the surface of the cluster. Mo et al.28a suggest that Uchg is due to a mechanism of delithiation followed by disproportionation, reactions 3 and 4. These authors used the first-principles surface supercell model to investigate this mechanism on a wide variety of facets and terminations, albeit with a very small unit cell so that the charge energetics investigated was for removal of the entire surface layer. However, they defined the overpotential from the potential at which the chemical barrier for O2 evolution is also downhill. This led these authors to suggest that the minimum overpotential for charging is ∼0.7 V on the O-rich (0001). We do not agree with this definition of overpotential since it is only the electrochemical steps that are affected by potential. Any purely chemical steps, e.g., O2 evolution from disproportionation, maintain the same barrier at all potentials and simply affect the pre-exponential factors for the overall electrochemical path described. Using the conventional definition of the thermodynamic overpotential suggests that they would have obtained ηchg = U − U0 ≈ 0.3 V for the delithiation + disproportionation mechanism, and this is comparable to the ηchg on this same facet and termination for the mechanism defined by eqs 1−3. Because Uchg starts at low potentials and then rises to higher values (see Figure 2), Kang et al.22a have revised this mechanism to suggest that topotactic delithiation alone occurs in the low potential region to form LiO2/Li2O2 phase separated nanoparticles, followed by some other higher overpotential process. The motivation for this suggestion was the experimental claim that significant concentrations of LiO2 are present in the cathode deposits.14,37 However, this evidence is evaluated in section 4 and is at best ambiguous. This mechanism also predicts that, in the low potential region, no O2 is evolved, and this is inconsistent with

desolvation barriers between each electrochemical step. Accurate calculation of these is presently beyond current theoretical capability for liquid electrolytes because of the complexities of describing the complex solvation rearrangements. For solid electrolytes, such calculations may be feasible and could provide an important benchmark in the theoretical description of the process. However, extensive experience in aqueous ORR/OER suggests that the thermodynamic limiting potentials dominate the kinetic overpotentials since desolvation barriers are modest.19a,32 The rate limiting desolvation barrier V*desolv in Li−O2 during discharge at U = U0 − ηdis is estimated to be ∼0.7 eV from temperature dependent studies of the overpotentials.9 This is likely too small to inhibit significant reaction on the time scale of typical galvanostatic discharge or charge experiments (≫1 s) at 300 K, so that the onset currents are principally determined by the thermodynamic limiting potentials rather than the desolvation barriers. In any case, these thermodynamic limiting potentials must represent the minimum overpotential observed. Similar first-principles thermodynamic calculations on the (1−100) facet gave ηdis = 0.68 V and ηchg = 0.20 V. However, calculations on steps and kinks of all facets gave very low overpotentials for both discharge and charge of 4 V (see section 5.2). This process is schematically shown in Figure 11. A simple kinetic model of parasitic electrochemical kinetics gave excellent agreement with the charging profile of Figure 2b.26 This origin of the rise in Uchg is the basic motivation for identifying the initial charging potential in the electrolysis cell as the kinetic overpotential. The unfortunate implication from this is that deposition of a small amount of solid electrolyte decomposition products at the Li2O2−electrolyte interface causes the rise in Uchg, and this is difficult to avoid given the stability issues of electrolytes with 11732


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particles of Li2O2−LiO2.17d,28,37 However, DFT calculations of multiphase LiO2 regions suggest they can only become stable at charging potentials.22a Thus, when the cathode is removed from the battery (U = 0), a large concentration due to multiphase regions of LiO2 does not seem likely. There are two experiments that seem to suggest significant fractions of LiO2 in the cathode deposit after it is removed from the battery (U = 0), although both have ambiguities. Gallant et. al reported X-ray absorption near edge structure (XANES) O K edge spectra of discharged cathodes and assign a single peak at 534.9 eV as evidence for LiO2 at the surface (total electron yield detection) by comparison to earlier studies by Ruckman et. al40 on alkali superoxides. However, both Ruckman et al. and Puglia et al.41 observe that another dominant spectral feature of alkali superoxides is the O2 π* resonance at 529 eV. The resonance at 529 eV was entirely absent in the XANES spectra of Gallant et al., implying that the presence of LiO2 is at best unclear from this spectroscopy. Yang et al.37 cite Raman evidence for a significant concentration of LiO2 in their discharged samples, i.e., a peak at ∼1125 cm−1 that is close to that calculated for the O−O stretching frequency in LiO2 and observed in other superoxides. We also sometimes observe this peak, but always coupled with a second peak at 1525 cm−1. However, these two peaks were only observed when using PVDF as a binder and never when using PTFE as the binder. These two Raman peaks have been assigned previously to surface modification of PVDF by alkaline treatment to produce (CHCF) species, with a CC band at 1129 cm−1 and a CC band at 1525 cm−1.42 Since Yang et al. used PVDF as a binder in their experiments, it is possible that the peak assigned by them to LiO2 is in fact due to decomposition of the PVDF binder by either LiOH (formed from residual H2O in the cell) or Li2O2. It is a common assumption that the rising Uchg results from charge transport issues in the Li2O2 deposit, essentially from an iR drop in the Li2O2 itself. However, this is not compatible with the fact that Uchg is smallest at the outset of charging when the Li2O2 deposits are largest (and Figure 24). Charge transport issues and their effects on Li−air are discussed in section 6. Because of the rising Uchg and the assumption that this is some type of charge transport issue from the cathode to the electrochemically active Li2O2 surface, some have suggested using a redox mediator in the electrolyte to carry the charge and induce the oxidation of Li2O2.43 Chen et al.43b have added the mediator tetrathiafulvalene (TTF) to a LiCLO4/DMSO electrolyte. They report highly rechargeable galvanostatic discharge−charge cycles with a low Uchg on a battery with a nanoporous Au cathode and LiFePO4 anode. However, the total discharge capacity in terms of mAh for these experiments was exceedingly small because of the very low weight and capacity of the Au cathode. We have also performed related experiments using TTF (and other redox mediators) in a Swagelok battery with LiTFSI/DME as the electrolyte, P50 C paper as the cathode, and Li metal as the anode. With 1 mAh discharges, DEMS measurements showed that the amount of O2 evolved relative to that consumed in the cycle (OER/ORR) was only ∼0.35, significantly less than without the mediator (see section 5), even though the rate of rise in Uchg was modest. This indicates that there is substantial parasitic chemistry/ electrochemistry involving the TTF with either Li2O2 or the Li anode.

Figure 12. (a) Pictorial representation of the electrochemical mechanism proposed by Lu and Shao-Horn to explain the origin of the rise in Uchg. Reprinted with permission from ref 21. Copyright 2012 American Chemical Society. (b) Pictorial representation of the electrochemical mechanism proposed by Gallant et al. to explain the origin of the rise in Uchg and its dependence on deposit morphology. Reprinted with permission from ref 35. Copyright 2012 American Chemical Society.

delithiation process with no O2 evolution is not consistent with DEMS observations (section 5) that sees (e−/O2)chg ∼ 2 at the outset of charging. Presumably, the oxidation of the bulk twophase region is accompanied by disproportionation (reaction 4) to account for O2 evolution. Zhai et al.22b suggest that the low potential region in charging is from LiO2 regions formed by incomplete disproportionation during discharge and that the high potential regions are for charging Li2O2. They argue that the current dependence of the low potential region implies a disproportionation mechanism for discharge, but for reasons that are unclear assume a kinetic model for disproportionation that is first order rather than second order in LiO2. However, the suggestion that incomplete disproportionation occurs in discharge making regions of LiO2 is incompatible with pressure decay results that (e−/O2)dis ≈ 2.0 (section 5). Also, assigning low potential charge regions to LiO2* → Li+ + e− + O2 (a 1e−/O2 process) is also incompatible with the DEMS results that (e−/O2)chg ≥ 2 (section 5). The basis for any mechanism involving a region involving only LiO2 is the assumption that multiphase regions of LiO2 in Li2O2 are stable, even though crystalline LiO2 does not appear to be stable to disproportionation above −35 C. When LiO2 was synthesized by the reaction of Li2O2 + O3 (in Freon) at T = −65 C, the resulting crystalline material with an EPR signal assigned to LiO2 disappeared as T was slowly raised above T = −35 C.39 Many DFT calculations suggest that the dominant surface of Li2O2 is O-rich, i.e., covered by a monolayer or half monolayer of LiO2 or that nanoparticles should be core−shell 11733


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4.1. Electrocatalysis

Because of the rising or high Uchg, there has been much activity into electrocatalysis as a way to lower Uchg. However, most of these studies only measure galvanostatic discharge−charge curves, with no quantitative measurements of what electrochemistry is occurring, so that it is impossible to judge whether the effects are genuine electrocatalysis effects related to the Li− O2 electrochemistry. Of course the possibility of electrocatalysis to lower Uchg or its rise during charging depends upon the origin of Uchg. If the inherent kinetic overpotentials are small as argued by Viswanathan et al.16 and the rising Uchg is due to solid decomposition products depositing at the Li2O2− electrolyte interface,26 then electrocatalysis is both unnecessary and impossible. On the other hand, if the mechanism for the rise in Uchg is in the fundamental charging electrochemistry,21 then electrocatalysis could in principle lower Uchg in some circumstances. However, even then the kinds of electrocatalysts that would be effective are likely to be quite different from those that are appropriate for aqueous ORR/OER where their role is to optimize the splitting/recombination of the O−O bond. For a reversible Li−O2 battery, 18O2 isotopic experiments clearly show that OER following ORR does not break the O−O bond.9,10 True electrocatalysis occurs when an active catalytic site lowers the barrier (or overpotential) to the kinetic limiting step in the mechanism, but the catalytic site itself is not changed during the reaction; i.e., it has a large turn over number. For Li−O2 ORR and OER, this implies that the reactant and product have mobility to and away from the catalytically active site. In a completely anhydrous electrolyte, since Li2O2 (and most likely LiO2) is insoluble in the electrolyte, it is difficult to see how these species could have enough mobility for true electrocatalysis to occur. For example, if some particle is a catalyst for ORR it would rapidly be covered up by solid Li2O2 and the dominant electrochemistry then becomes simply Li− O2 ORR on Li2O2. On the other hand, in OER, it is unclear how the Li2O2 crystal (or LiO2 on the surface of the crystal) gets to the catalyst particle since they are presumed insoluble. Of course during the initial nucleation/final dissolution phase, when the electrochemistry is dominated by Li−O2 electrochemistry on the cathode surface itself, electrocatalysis is inherent in the nature of the surface. However, this represents a minimal part of the total discharge/charge. If there are soluble intermediates during OER and Li2O2 does not cover the catalyst nanoparticle, then true electrocatalysis may be possible. For example, this can occur when the dominant electrochemistry during charge is from parasitic reactions with the electrolyte that generate soluble intermediates. McCloskey et al.44 suggested that this is the reason that electrocatalysis is observed when using carbonates as electrolyte solvents for Li−O2, but none observed when using relatively stable anhydrous ethers20 (DME) as the electrolyte solvent. This is demonstrated in Figure 13. The apparent electrocatalysis with carbonate solvents dominated interest in the early Li−air days before it became clear that this electrochemistry is dominated by parasitic reactions with the solvent. Another possibility for Li−O2 electrocatalysis is when some H2O is present in the nonaqueous electrolyte since the formation of toroids suggests that there could be some soluble intermediate in the electrochemistry. At the time of writing, we know of no careful electrocatalysis studies with added H2O that combine coulometry with quantitative chemical analysis of the electrochemistry, so this possibility remains open.

Figure 13. (a) Galvanostatic discharge−charge of 1 mAh discharge when using 1 M LiTFSI in organic carbonates as the electrolyte and with various nanoparticle catalyst particles on XC72 C cathode. The apparent electrocatalysis is related to solvent decomposition. (b) Similar experiments but using 1 M LiTFSI in dimethoxyethane as the solvent. Reprinted with permission from ref 44. Copyright 2011 American Chemical Society.

Before leaving this section, we do want to acknowledge a few of the many electrocatalysis studies based only on coulometry. There are also many reviews of this.45 As stated repeatedly, however, it is impossible to judge whether any of these studies actually involve true electrocatalysis of the Li−O2 electrochemistry since there is no quantitative resolution of the chemistry occurring at the lowest currents, i.e., whether it involves simply ORR or OER or some other parasitic chemistry/electrochemistry (as in the case of carbonate solvents). One interesting example is the apparent electrocatalysis for Li−O2 ORR via different metals that form a nice volcano relationship arising from the Sabatier principle in catalysis as shown in Figure 14.25 However, many of the metals are known from the surface science literature to readily dissociate O2 (Pt, Pd, Ru), so that it is unclear from the coulometry alone whether the enhanced electrochemistry corresponds to formation of Li2O2. In fact, Barile and Gewirth46 showed that Pt, Pd, and Cu(II) oxide catalysts increased the amount of CO2 evolved and decreased O2 evolution. McCloskey et al.44 also showed that Pt nanoparticles enhanced electrolyte decomposition in the presence of O2 even at open circuit. We should also point out that, in the strong oxidizing conditions for OER, it is not apparent whether the reactive metal surfaces themselves are oxidized so that the apparent catalysis may be for the oxides.

5. PARASITIC CHEMISTRY There are numerous studies in the literature claiming extended Li−O2 cycling. However, most of these are based solely on coulometric measurements, perhaps combined with some qualitative spectroscopic measurement (XRD, etc.) that some unspecified amount of Li2O2 is formed during discharge and disappears during charge. In this section we discuss the parasitic 11734


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= Li2O2. Therefore, a 2 e−/O2 process should ideally be observed on both discharge and charge and correspond to the stoichiometric equivalent Li2O2 formation and oxidation. Several quantitative chemical techniques need to be combined to identify components formed by the parasitic chemistry and to assess the rechargeability. A combination of pressure decay/rise measurements in the cell headspace and quantitative differential electrochemical mass spectrometry (DEMS) fully quantifies gas consumption and evolution during a Li−O2 cycle, i.e., it allows a quantitative measurement of (e−/ O2)dis, (e−/O2)dis, OER/ORR, and parasitic gas evolution.10,46,50 Chemical titrations for Li2O2 have been developed and allow measurements of YLi2O2.51 Various spectroscopic techniques (XRD, IR, Raman, XPS, XANES, NMR, etc.) have been used to identify parasitic components formed during discharge or discharge−charge, but are generally used only as a qualitative indicator of the parasitic chemistry. A few quantitative applications of spectroscopy to Li−O2 will be discussed later. Obviously, all spectroscopic studies on cathodes must be accomplished without any contact to air from battery through spectroscopic measurement. Otherwise, impurities such as Li2CO3, LiOH, etc. could be observed that are not characteristic of the discharge. For spectroscopic and topographic techniques that require vacuum (XPS, XANES, SEM, TEM), this implies a load-lock system (or glovebag) for the transfer. Comparing results from two previous studies of Li−O2 with a LiTFSI/DME electrolyte and C cathode highlights the limitations of two of the most common qualitative spectroscopic techniques to identify product formation. The only product observed using both XRD and Raman spectra of discharged cathodes was Li2O2.10 However, a quantitative peroxide titration protocol finds only 91% Li2O2 yield during discharge.51c This implies that other compounds are produced during discharge, and this was confirmed using other spectroscopic techniques (discussed below).12,20a,51c,52 XRD is invisible to amorphous products, and both techniques are too limited in sensitivity to observe small concentration components. Since Li2O2 is very sensitive to X-ray, electron beam, and ion beam induced damage, spectroscopic techniques based on these probes require special care and dose dependent studies to separate native products from those produced by beam-induced chemistry. Our experience is that both XPS and SEM are very dose dependent. XPS studies by Younesi et al.53 of Li2O2 parasitic chemistry with electrolytes showed composition changes in products upon X-ray exposure. Two of the spectroscopic techniques that are sufficiently sensitive to detect minor impurities and can potentially be quantified are NMR and IR. Proton and fluorine NMR of both the electrolyte and D2O-extracted cathode products, pioneered by Bruce’s laboratory20a and Nasybulin et al.,54 is an excellent way to determine impurities in either the electrolyte or the cathode. Inclusion of an internal standard allows quantification of the amount of parasitic product produced. This showed that the products LiHCO2 and LiF accounted for most of the Li2O2 yield loss observed during discharge.51c Solid state NMR of the discharge product is also possible, although chemical shifts for 6 Li and 7Li are so small as to not be particularly chemically diagnostic.55 17O NMR is considerably more chemically diagnostic (but more expensive).55c The latter did identify Li2O2 as the dominant product formed with a LiTFSI/DME

Figure 14. Apparent nonaqueous Li+−O2 reduction potentials for different metal surface measured at 2 μA/cm2 current density as a function of the calculated oxygen adsorption energy ΔE0 relative to that of Pt. Reprinted with permission from ref 47. Copyright 2011 American Chemical Society.

chemistry and electrochemistry of Li2O2 with the electrolyte, cathode, and components of air other than O2. Quantitative studies highlighted in this section show that considerable parasitic chemistry occurs with all battery components, so that claims of high rechargeability should be accepted cautiously. 5.1. Quantitative Definitions of Li−O2 Rechargeability

In Li−ion batteries, rechargeability is essentially determined by Coulombic efficiency since parasitic chemistry/electrochemistry is minimal. However, completely different definitions are necessary for Li−O2 because of the more extensive parasitic chemistry/electrochemistry. This was underscored by the early claims of high Li−O2 rechargeability using carbonate solvents in the electrolyte.7,48 In that case, although many Coulombic cycles were possible, each cycle was dominated by solvent decomposition.8b−d,a Key quantitative chemical signatures of Li−O2 rechargeability are the following: (1) the yield of Li2O2 relative to that anticipated from the current and ideal cathode reaction 2(Li+ + e−) + O2→ Li2O2 during discharge is YLi2O2 = 1.00, i.e., no other products are formed during discharge either on the cathode or in the electrolyte, e.g., no LiOH, Li2CO3, LiF, carboxylates, etc.; (2) during discharge, the only electrochemical current consumes O2, (e−/O2)dis = 2.00, and during charge, all electrochemical current evolves O2, (e−/O2)dis = 2.00; (3) no parasitic gas evolution (H2, CO2, etc.) occurs during the discharge−charge cycle; (4) all O2 consumed during discharge (ORR) is released during charge (OER) so that OER/ORR = 1.00 If all requirements are met, then the Li−O2 battery is perfectly rechargeable. In addition, a long calendar life is necessary, and this requires that all thermal parasitic chemical reactions between components of the battery are minimal or at least self-limiting, e.g., between Li metal or Li2O2 and the electrolyte.49 The motivation for these criteria is obvious. For a Li−O2 battery to be considered truly rechargeable, all oxygen consumed during a galvanostatic discharge (ORR) needs to be evolved during a galvanostatic charge (OER) of equal capacity, and this OER/ORR ratio should remain unity for many cycles. This requirement also implies that no other gases are evolved besides O2 during charge. As we discuss in this section, the only reversible active electrochemical cathode reaction identified (via isotopic O2 labeling) is 2(Li+ + e−) + O2 11735


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solutions,58 and this produces a false negative Li2O2 yield with those cathodes. Hase et al.51b developed a nonaqueous titration that relies on the oxidation of an oxoammonium salt. However, this technique is sensitive to parasitic side products (Li2CO3, LiOH, etc.) that can also be oxidized.

electrolyte and showed that Li2CO3 was the dominant product formed with LiTFSI/carbonate as the electrolyte. FTIR can also identify minor parasitic side products and Li2O2 in the cathode.8c,20a,38,52 Peng et al.56 also calibrated their FTIR results by analyzing known composition mixtures of Li2O2 and Li2CO3 or LiHCO2, thereby allowing a quantitative comparison of the amount of these parasitic side products to Li2O2 formed during discharge. In addition, Lim et al.57 devised an in situ XRD technique to analyze Li2O2 diffraction intensities during discharge and charge. They quantified by comparing to Li2O2 scattering intensities from (presumably) pure Li2O2. However, they obtained much lower YLi2O2 than analogous measurements in our laboratory using ex situ titrations. The origin of their low YLi2O2 is unknown, but may be related to the unusual cell architecture and/or X-ray damage. Various titration techniques for determining the total moles of Li2O2 produced during discharge and therefore YLi2O2 have been devised. An example is given in Figure 15 and compared

5.2. Electrolyte Stability

Despite the many claims of Li−O2 rechargeability in the literature, no electrolyte satisfies the quantitative chemical criteria given in the previous section for rechargeability in a single discharge−charge cycle, let alone many such cycles. Therefore, identifying a stable electrolyte is the most pressing scientific challenge currently facing nonaqueous Li−air batteries. Once the issue of carbonate-based electrolyte stability became identified, many studies have explored electrolyte stability with various solvent classes and salts. McCloskey et al.10 explored numerous anhydrous solvent/salt combinations using DEMS and pressure decay/rise measurements, where ORR and OER can be measured directly along with other gas evolution. Various measures of rechargeability for a variety of electrolytes are summarized in Table 1. The most important Table 1. Summary of DEMS Results for Various Salt and Solvent Combinationsa cathode

solventb

XC72

DME

P50

MPPTFSI DMSO 1NM3 NMP THF DME CH3CNc TGE

P50 XC72 XC72 XC72 XC72 P50 XC72

Li+ saltd

OER/ ORR

(e−/ O2)dis

(e−/ O2)chg

CO2/ ORR

H2/ ORR

TFSI Trif ClO4 BF4 BOB TFSI

0.78 0.74 0.77 0.78 0.36 0.33

2.01 2.00 2.00 2.06 2.33 2.30

2.59 2.71 2.59 2.65 6.41 7.04

0.07 0.05 0.05 0.04 1.26 0.01

0.03 0.08 0.08 0.08 0.01 0.28

0.51 0.48 0.58 0.72 0.78 0.88 0.75

2.05 2.14 1.96 2.01 2.01 2.05 2.04

4.05 4.44 3.35 2.80 2.59 2.33 2.71

0.03 0.11 0.03 0.03 0.06 0.04 0.03

0.02 0.04 0.02 0.09 0.01 0.01 0.08

BF4

a

Reprinted with permission from ref 10. Copyright 2012 American Chemical Society. bDME = dimethoxyethane, THF = tetrahydrofuran, TGE = triglyme, CH3CN = acetonitrile, DMSO = dimethyl sulfoxide, NMP = N-methyl pyrrolidone, 1NM3 = tri(ethylene glycol)substituted trimethylsilane, MPP-TFSI = N-methyl-N-propylpiperidinium bis(trifluoromethylsulfonyl) imide. cExperiment performed using lithium iron phosphate (LiFePO4) as the anode. Otherwise, Li metal was used as the anode for all experiments. dLiTFSI = lithium bis(trifluoromethane sulfonyl) imide, LiBF4 = lithium tetrafluoroborate, LiTrif = lithium triflate, LiBOB = lithium bis(oxalato) borate, LiClO4 = lithium perchlorate.

Figure 15. (a) Galvanostatic discharge−charge of Li−O2 batteries employing 1 N LiTFSI in DME as an electrolyte and P50 carbon paper as the cathode. (b) O2 consumption, nO2,d, and Li2O2 formation, nLi2O2,d, during discharge. (c) O2 evolution, nO2,c, and Li2O2 oxidation, nLi2O2,c, during charge. Reprinted with permission from ref 51c. Copyright 2013 American Chemical Society.

conclusion from Table 1 is that although some electrolytes have an (e−/O2)dis that is close to an ideal value of 2 (e.g., DME and DMSO), no electrolytes have an (e−/O2)chg value close to 2, and OER/ORR is less than 0.9 for all electrolytes after a single galvanostatic cycle. Except for CH3CN that is not stable to Li metal, the best solvents were ethers, especially DME. This is the reason that this solvent was used in most of the examples described in this review. However, it must be emphasized that this solvent is by no means stable enough for a practical commercial battery.

to the O2 consumed and produced, but will be discussed in detail later. McCloskey et al.51c used an iodometric titration for H2O2 after cathodes were dissolved in H2O. Extensive calibration showed that the reaction Li2O2 + 2H2O → 2LiOH + H2O2 was quantitative for Li2O2 formed electrochemically on a C cathode. However, some cathode materials currently being explored catalyze peroxide oxidation in aqueous 11736


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Xu, Zhang, and co-workers used a variety of spectroscopic techniques to study the discharge product and in situ gas chromatography mass spectrometry (GC/MS) to study gas evolution during charge in order to probe solvent and salt stability.8d,54,55,59 The GC/MS characterization for multiple electrolyte solvents indicated considerable gas evolution other than O2. They showed that electrolytes composed of an ethereal solvent formed the most Li2O2 during discharge and evolved the most oxygen during charge. Given the DEMS results, electrolyte stability limitations were initially thought to be due to electrochemical parasitic reactions occurring during charge, because (e−/O2)dis is close to its desired value, but not (e−/O2)chg.10 However, these DEMS results are also compatible with a mechanism whereby thermal chemical reactions between Li2O2 and the electrolyte occur after Li2O2 has been electrochemically formed (via a 2 e−/O2 process) on discharge. This would cause a reduction in the total amount of Li2O2 produced during discharge and less O2 evolution during charge. Measurement of YLi2O2 via a titration resolves this uncertainty concerning the loss mechanisms in rechargeability. Figure 15 presents O2 consumption (measured using pressure decay) and Li2O2 formation (measured using iodometric titration) during galvanostatic discharge of a battery with a LiTFSI/DME electrolyte and carbon cathode.51c This combination produced among the highest OER/ORR characterized in previous studies (see Table 1).10 Although a ∼2e−/O2 process is observed during discharge, more O2 is consumed than Li2O2 produced, YLi2O2 < 1.0, indicating that O2 is being lost to parasitic reactions during discharge. McCloskey et al.51c also quantified YLi2O2 in batteries with carbonate, DMSO, and TEGDME-based electrolytes (Table 2) and find that the highest YLi2O2 is only 91%, the yield of the battery shown in Figure 15.

Figure 16. (a) Yield of Li2O2, YLi2O2, as a function of the Li−O2 battery discharge rate. (b) O2 consumption during the first 10 h of cell discharge at various discharge rates and corresponding e−/O2 values. Reprinted with permission from ref 51c. Copyright 2013 American Chemical Society.

allowing a slow parasitic chemical reaction to occur between the Li2O2 and electrolyte. A small amount of parasitic electrochemistry may also occur given the slightly higher than 2e−/O2 process, but this accounts for only a small fraction of the total decomposition observed. Thus, Figure 16 provides indirect support to the conclusion that thermal chemistry of Li2O2 with the electrolyte is the dominant parasitic process in discharge. Because electrolyte stability is such an important and wellrecognized challenge for Li−O2 batteries, there have been numerous studies of product formation during discharge with various electrolytes, including those used in Figure 15 and Table 2. They have identified many parasitic products originating from the electrolyte. Carbonates,8a,c,d amides,50a,60 ethers,20a,52 nitriles,60b,61 sulfoxides,49,60b,62 sulfones,61 and various Li salts54,63 have all qualitatively been shown to partly decompose during discharge (and extended cycling, which will be discussed later). This is entirely consistent with the DEMS and titration studies with similar electrolytes.8b,10,51c For example, the most prominent parasitic products identified in batteries using LiTFSI/DME as electrolyte and C cathode are lithium formate (LiHCO2), lithium acetate (LiCH3CO2), lithium carbonate (Li2CO3), and lithium fluoride (LiF), all of which are attributed to decomposition of DME and LiTFSI.20a,51c,54 NMR and carbonate titrations indicate that more LiHCO2 and LiF are formed during discharge than LiCH3CO2 and Li2CO3.51c The low YLi2O2 arises from the highly reactive nature of Li2O2 as it is electrochemically forming during discharge. Li2O2 is an extraordinarily strong nucleophile, hydrogen abstractor, and oxidizer and can therefore react with organic molecules in a variety of ways. It is commonly assumed that LiO2, the weaker nucleophilic intermediate for Li2O2 formation, induces decomposition of electrolytes. However, this would lead to a 1e−/O2 process on discharge (i.e., LiO2 formation via a 1e−/O2 process, followed by chemical decomposition by the LiO2). This reaction pathway does not seem to contribute significantly to the chemical decomposition processes since (e−/O2)dis ∼ 2 is always observed, as shown in Figure 15 and Table 1. Given the extraordinary reactivity of Li2O2, numerous chemical electrolyte decomposition pathways are possible during discharge. Understanding these pathways could help in the design of a stable electrolyte through molecular engineering. Theoretical calculations by Bryantsev et al.61 using solvated O2− as a proxy to study solvent stability in the presence of a reduced oxygen nucleophile show that numerous solvent classes, including carbonates, sulfonates, lactones, sulfones, and phosphonates, are unstable to nucleophilic substitution.

Table 2. Li2O2 Titrations (YLi2O2)a electrolyte 1 N LiTFSI/DME

1 N LiTFSI/DMSO 1 N Li triflate/tetraglyme

cathode P50 PTFE-coated P50 KB (PTFE-bound) super P (PTFE-bound) XC72 (PTFE-bound) 13 C (PTFE-bound) XC72 (Nafion-bound) P50 P50

YLi2O2 (%) 90.9 90.3 85.5 86.3 83.3 77.2 84.4 80.9 81.9

± ± ± ± ± ± ± ± ±

0.8 1.1 1.0 1.0 0.5 0.7 0.8 1.3 0.8

a

Reprinted with permission from ref 51c. Copyright 2013 American Chemical Society.

The conclusion from Figure 15 and Table 2 is that parasitic reactions occurring in the battery during discharge are primarily chemical rather than electrochemical in nature. Figure 16 presents YLi2O2 as a function of discharge current i, and corresponding O2 consumption for a battery similar to that in Figure 15. YLi2O2 decreases at low i although (e−/O2)dis remains constant and is independent of i. Solid Li2O2 growth on the cathode is the dominant discharge mechanism in this battery, and this implies that only the surface layer of Li2O2 can react chemically with the electrolyte. At low discharge rates, the surface remains exposed to the electrolyte for a longer time, 11737


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Recognizing that the surface of solid Li2O2 is the dominant reactive species, Laino and Curioni64 studied the decomposition mechanism of propylene carbonate at this surface. They subsequently studied nucleophilic attack of carbonates, amides, and sulfoxides, and hydrogen abstraction from ethers on Li2O2 surfaces, all of which were also reported to have low to nonexistent activation barriers.65 Assary et al. studied DME decomposition mechanisms at the surface of Li2O2 clusters (representing the Li2O2 surface) and showed that α/βhydrogen abstraction possesses a relatively low activation barrier (∼1 eV).66 There have been several experimental attempts to develop screening methods for stability using KO2, either as a solid or combining it with a crown ether that solubilizes it to form O2− or reacting it with Li+ that leads to an ion exchange, forming LiO2, which then quickly dismutates into Li2O2 and O2.42b,61,67 There have even been attempts to measure electrolyte stability by simply adding Li2O2 to the electrolyte.60b However, when there are parasitic chemical reactions that produce insoluble species, a thin decomposition layer may form on the surface of the solid KO2/Li2O2, and this thin layer passivates further reaction with the electrolyte. Thus, even electrolytes with known stability issues in Li−O2 batteries can appear stable using these techniques.60b Therefore, the best screening method for electrolyte stability seems simply to discharge a battery and characterize quantitative rechargeability using the techniques outlined in section 5.1. In addition to the dominant parasitic chemical processes occurring during discharge, there are also electrochemical parasitic processes that occur during charge. The linear scan voltammetry of Figure 6 demonstrates this emphatically. In fact, linear scan voltammetry of all the electrolytes in Table 1 show that (e−/O2)chg always rises significantly from ∼2 at potentials ∼0.5 V below the solvents’ characteristic electrochemical oxidation potential.10 Although several attempts to rationalize a lowering of the solvent electrochemical window by Li2O2 were initially proposed,10 it now seems likely that oxidation of the parasitic products formed in discharge are in part responsible for the lowering of the overall electrochemical stability in discharged batteries. Figure 15 demonstrates the lower charge stability issue in galvanostatic discharge−charge since the O2 evolution is always less than the Li2O2 oxidation during charge. Therefore, a small fraction of the Li2O2 oxidation is contributing to a parasitic processes rather than O2 evolution. The net result in Figure 15 is that of the loss of Li2O2 to parasitic reactions is ∼9% for discharge, but ∼13% for charge. The same dual parasitic processes are also prevalent for DMSO and TEGDME-based electrolytes.51c As was discussed in section 4, one suggestion for the rise in Uchg is that the parasitic processes discussed here deposit solids (carbonates, carboxylates) at the electrolyte−Li2O2 interface during charging, and these have higher oxidation potentials than Li2O2. This causes the mixed oxidation potential to rise during charging to those appropriate for carbonate/carboxylate oxidation. The discharge and charge parasitic reactions outlined above naturally result in severe cycling limitations. Figure 17 presents many of the quantitative chemical measures of rechargeability over the first 5 cycles of a Li−O2 battery with LiTFSI/DME as the electrolyte.10 This represents one of the more stable electrolytes studied to date. All measures of rechargeability continuously deviate from the ideal limits with cycling. This implies that the accumulation of decomposition products in the

Figure 17. (a) Five galvanostatic discharge−charge cycles for a battery with 1 N LiTFSI/DME electrolyte and a carbon cathode. (b) Oxygen consumption during discharge and evolution during charge for the five cycles. (c) OER/ORR and CO2 evolved during charge (mCO2) divided by ORR. (d) e−/O2 for discharge and charge of each cycle. Reprinted with permission from ref 10. Copyright 2012 American Chemical Society.

battery with cycling leads to even more parasitic chemical and electrochemical reactions in the higher cycles. This accumulation of decomposition products both on the cathode and in the electrolyte ultimately leads to a reduction in cell capacity. Many authors use repeated galvanostatic cycles, often based on only limited discharge−charge capacities, to discuss rechargeability. Figure 17 emphasizes that this can be quite misleading since several galvanostatic cycles show an increase in capacity before decreasing, while the quantitative measures of rechargeability continuously decrease. Some solvents may be more stable than the ethers now commonly used for electrolytes, e.g., dimethylacetamide (DMA) or nitriles, and therefore may sustain more cycles. However, these do not form good SEI with Li metal and so require an artificial SEI. Walker et al.60a reported extended cycling using a LiNO3/DMA electrolyte and carbon cathode. The LiNO3 formed a good SEI on the Li metal and therefore protected the DMA from continuous reduction by the Li metal. Using a nonquantitative DEMS, they showed that O2 is the dominant gas evolved in the 80th charge, with very little CO2 or H2 evolved. 1H NMR of their solvent after 65 cycles also confirmed that only minor solvent decomposition occurred. However, the authors of this study pointed out that the capacity of each cycle was limited to ∼15% of the full initial capacity and if the full initial capacity was used, a significant decline in discharge capacity was observed after only 5 cycles. In summary, electrolyte stability issues arise from parasitic chemical reactions of the electrolyte components with Li2O2 during discharge and electrochemical reactions during charge. Both the solvent and salt seem to have stability issues. We are currently unaware of any electrolyte that is sufficiently stable to give a truly rechargeable Li−O2 battery. Several authors have reported extended cyclability using cathodes other than C, but the same electrolytes as discussed here. Some of these studies also measure some of the quantitative measures of rechargeability and are discussed in the following section. However, it is unclear how cathode composition can dramatically affect the electrolyte chemical stability to Li2O2. 11738


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5.3. Cathode Stability

All results discussed so far used C cathodes for obvious reasons. Some of these use C nanoparticles and a binder, while others are binder free (P50 Avcarb, VACNT). However, it is wellknown in the fuel cell literature that C corrosion is a problem at high oxidizing potentials during fuel cell startup in the presence of O268 so that it is likely an issue as well in Li−O2 batteries, especially during charge. McCloskey et al. 26 originally demonstrated that C cathode oxidation in Li−O2 batteries is a serious concern. When they used a 13C- carbon black cathode, ∼ 40% of the total CO2 evolved during charge was 13CO2, confirming that some of the parasitic decomposition in the full discharge−charge cycle involved the cathode (see Figure 18).

Figure 19. (a, b) Galvanostatic discharge−charge coulometry on the first cycle for DMSO and tetraglyme electrolytes, respectively, at a 13C carbon cathode. (c, d) Moles of CO2 evolved from the carbon cathodes, removed from the cells at the states of discharge and charge indicated by the green dots in parts a and b, and then treated with acid to decompose Li212CO3 and Li213CO3 and Fenton’s reagent to decompose the lithium carboxylates. The values (166 etc.) on the xaxis do not represent a scale but indicate the states-of-charge at which the cathodes were sampled. The numbers 2, 3, and 5 (c) and 2 and 5 (d) correspond to the electrodes analyzed at the end of discharge on those cycles, and the number 2C in part c corresponds to the analysis of the carbon electrode at the end of the second charge. Reprinted with permission from ref 69. Copyright 2013 American Chemical Society. 13

C cathode oxidation to form Li2CO3, but throughout the charge cycle, not just at potentials >3.5 V. Thotiyl et al. also reported that the carbon cathode’s relative hydrophilicity/ hydrophobicity affects both the carbon and electrolyte stability, with more hydrophilic carbon leading to increased carbon and electrolyte decomposition.69 The mechanism of cathode corrosion is not fully understood. McCloskey et al.26 used both 13C cathode and 18O2 isotope labeling to show that a significant fraction of the O in the 13 CO2 evolved at the end of charge originated from the electrolyte (Figure 18). This suggests that a mechanism for C corrosion is the electrochemical formation of highly reactive mobile species from the electrolyte parasitic reactions, which then diffuse to the cathode surface and react with the C. This is consistent with observations by Gallant et al. that Li2CO3 builds up at the C−electrolyte interface with cycling.35 For cathodes using a binder, its stability is also an issue. Stability issues using PVDF/Kynar have been demonstrated by Xu et. al and Black et al., and as discussed previously in section 4.8d,42b This decomposition may be the origin of the LiO2-like peaks observed in Raman using this binder. It is unclear, however, whether PVDF is chemically unstable in a completely anhydrous electrolyte. Nasybulin et al. presented an exhaustive study of binder stability in which they ball mill various candidate polymers with KO2 and Li2O2.70 Of 11 different polymers studied, polyethylene appeared to be the most stable, with PTFE being slightly less stable, and PVDF being considerably less stable. McCloskey et al. showed that the Li2O2 yield for discharge of a C paper without binder, with a PTFE binder, and with a lithiated Nafion were all independent of the binder.51c Therefore, although binder stability may be a secondary problem (especially with PVDF), electrolyte and carbon stability issues now dominate the loss in rechargeability.

Figure 18. (a) Galvanostatic discharge−charge profiles of a Li−O2 battery with LiTFSI/DME as the electrolyte and a 99% 13C cathode. (b) Molar evolution rates (m′) for O2, 13CO2, and 12CO2 + 13CO2 during the galvanostatic charge measured by DEMS. (c) m′ of various CO2 isotopes during a linear oxidative scanning voltamogram (0.1 mV/s) after a 1 mAh discharge under 18O2. Reprinted with permission from ref 26. Copyright 2012 American Chemical Society.

They suggested that the 13CO2 was evolved from Li213CO2 produced by a thermal chemical reaction of Li2O2 with the C during discharge and that the 12CO2 was formed by parasitic reactions of Li2O2 with the electrolyte. However, several experiments have now shown that most of the Li213CO2 is produced during charge rather than discharge. Xu et al. used 13 C magic angle spinning solid state NMR with 13C-enriched electrodes to show that there was no enhancement of Li213CO2 after discharge in ethereal-based electrolytes.59 Thotiyl et al.69 quantified Li2CO3 and Li carboxylate formation during discharge and charge of a 13C cathode battery by monitoring CO2 evolution during an acid wash (to oxidize the carbonate) and Fenton’s reagent wash (to oxidize the carboxylates) of extracted cathodes. They find little carbonate formation from cathode decomposition during discharge, but that C oxidizes at charge potentials >3.5 V in both DMSO- and TEGDME-based electrolytes (see Figure 19). McCloskey et al.51c also observe 11739


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5.3.1. Non-Carbon Cathodes. Because of the problem of C corrosion when used as the Li−O2 cathode, many authors have attempted to discover new cathode materials that do not react with Li2O2 or other parasitic products. However, in most of these studies quantitative measures of rechargeability have not been determined so it is impossible to accurately judge claims of rechargeability. There are, however, a few very visible and highly encouraging studies that suggest extended rechargeability by using alternate cathode materials. Peng et al.56 report that batteries containing nanoporous gold (NPG) cathodes and a LiClO4/DMSO electrolyte give 100 cycles with nearly perfect ∼2.0e−/O2 for the 100th discharge and charge as shown in Figure 20 (it should be noted

Figure 21. Galvanostatic discharge−charge curves and capacity retention of TiC cathodes. (a) Galvanostatic discharge−charge cycles recorded in 0.5 M LiClO4 in DMSO at a geometric current density of 1 mA cm−2. (b) Capacity retention for the same battery as in part a. Inset to part b: ratios of the number of electrons to oxygen molecules observed at a TiC cathode on discharge and charge in 0.5 M LiClO4 in DMSO. Reprinted with permission from ref 71. Copyright 2013 Macmillan Publishers Ltd. Figure 20. Charge/discharge curves (a) and cycling profile (b) for a Li−O2 cell with a 0.1 M LiClO4−DMSO electrolyte and a NPG cathode, at a current density of 500 mAg−1 (based on the small mass of Au in the cathode). Right inset, ratios of the number of electrons to oxygen molecules upon reduction (discharge) and oxidation (charge). Reprinted with permission from ref 56. Copyright 2012 AAAS.

increased cyclability is possible with a more stable carbon-free cathode, DMSO is well-known to oxidize to dimethylsulfone via a chemical reaction with Li2O251c,62 so that parasitic chemical reactions with the electrolyte during discharge should occur unaffected by the change in cathode material. Inspired by Peng et al. and Thotiyl et al.’s results, we have tried to duplicate these experiments as closely as possible, i.e., with a Li metal anode, a thick 15 μm NPG cathode, or a TiC cathode similar to the one by Thotiyl et al., and LiTFSI/DMSO as the electrolyte. Figure 22 presents our results for oxygen reduction and evolution during the first galvanostatic cycle on the NPG and TiC cathodes. Unfortunately, these studies were unable to reproduce the spectacular results obtained by Peng et al. and Thotiyl et al. for reasons that remain unclear. Using OER/ORR to quantify rechargeability of a single cycle, both the NPG (OER/ORR = 0.44) and TiC (OER/ORR = 0.45) cathodes actually exhibit worse rechargeability on the first cycle than carbon cathodes (OER/ORR = 0.51) using the same electrolyte. Dimethylsulfone formation is readily observed using 1 H NMR of the extracted electrolyte after a modest capacity discharge.51c Thus, these results are not consistent with extended cycling by simply changing cathode material. However, the experiments of Figure 22 are not completely identical to those of Peng et al. and Thotiyl et al. Figure 22 used thicker nanoporous gold electrodes than those used in Peng et al. and a Li metal anode instead of LiFePO4 as the anode as in Thotiyl et al. However, it is difficult to see how these differences could radically change the rechargeability observed. Our results highlight the fact that parasitic chemical reactions between electrodeposited Li2O2 and the electrolyte can still

that graphical integration of the DEMS curves presented in this article do not support that OER/ORR ∼ 1.0). The coulometric cycles (U vs Q) also closely overlapped for cycle 1−100. Quantitative FTIR studies of the discharge product after many cycles showed no evidence of LiCO2H or Li2CO3 contamination. Because of the very thin NPG cathodes (100 nm) and consequently limited weight of the cathode, the total discharge capacity (in terms of mAh) in the cycles was very small. In an attempt to find more practical inexpensive and lightweight cathode materials, Thotiyl et al.71 screened cathodes of nanoparticulate TiC, SiC, and TiN bound to a stainless steel mesh using PTFE with LiCLO4/DMSO and LiPF6/TEGDME as electrolytes and LiFePO4 as the anode. An acid wash to evolve CO2 from carbonates that were produced showed that the TiC cathode and LiCLO4/DMSO electrolyte formed 0.1 mA/cm2 projected area), hole tunneling is the dominant charge transport mechanism in anhydrous Li−O2 batteries. Since tunneling length scales d will always be limited to 5−10 nm before sudden death, another charge transport mechanism is necessary to have high capacities at high currents. Certainly, hole polaron charge transport gives longer length scales, but is presently limited to supporting only low currents. One possibility is to raise the temperature of the Li−O2 battery so that the hole polaron migration supports higher currents, although this clearly has adverse consequences to other aspects of Li−O2 batteries, e.g., parasitic reactions, etc. Another possibility is to admix some impurity into the Li−O2 deposit that either induces states at εF or donates charge and shifts the

Figure 24. Galvanostatic discharge followed by galvanostatic charge, both at 100 μA/cm2. The different curves show discharge to different “sudden death” potentials (Ubias) before initiating recharge. The plots are given as a function of relative capacity Qdis/Qmax to emphasize that the initial charging potential is independent of the final discharge potential. Reprinted with permission from ref 81. Copyright 2013 American Chemical Society.

that charge transport issues enhance overpotentials during charging, these observations imply that charge transport is a less significant problem in charge than discharge. This is rationalized by noting that both the hole tunneling barrier and polaron populations are strongly (exponentially) dependent on εF, and εF depends in an opposite manner with overpotential during discharge and charge. Varley et al.80 suggest that the dominant effect on hole tunneling of the charged vacancies and hole polarons is to serve as repulsive scattering sites, and in aggregate these act as a resistance for the 11743


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VBM closer to εF. Several possibilities have been discussed,83 but none demonstrated yet to be of practical value. It will be very difficult to find some dopant that reversibly deposits with the Li2O2 during discharge and evolves with the Li2O2 during charge.

7. LI ANODE In all discussions of Li−air, it is tacitly assumed that Li metal is the anode because use of todays Li−ion anodes (Li intercalated graphite) reduces the theoretical fully charged specific energy to ∼1000 Wh/kg. This means that the practical specific energy gains relative to current Li−ion will undoubtedly be much more modest than when using Li metal, although no system analysis has yet been done using a Li/C anode. In addition, future Sibased anodes may not be practical for Li−air because of unavoidable oxidation of Si to form SiO2. A Si anode-based Li− O2 battery has been reported;84 however, chemical aspects of its rechargeability are still unknown. Unfortunately, attempts to use Li metal as an anode in a secondary battery has a long and to date unsuccessful history. In fact, it was only when Li intercalated graphite, with its self-limiting and stable SEI, replaced Li metal as the anode that Li−ion batteries became practical. However, there continues to be extensive research into trying to overcome the problems of cycling Li metal, in part because this is recognized as a necessary ingredient in many of the BLI chemistries (Li−O2, Li−S, Li−polymer, advanced Li−ion). An excellent review of past and current research into using Li metal is in ref 85. The chief problem with using Li metal is associated with its morphological changes in cycling, a common issue in cycling many different metals. This has two important consequences: (1) formation of dendrites, which can cause internal shorts in the cell and results in a serious safety issue, and (2) the need to reform the solid electrolyte interface (SEI) with cycling, which consumes Li and electrolyte. This results in a low Coulombic efficiency and therefore limited cycle life, as well as requiring excess Li in the anode. Strategies that minimize dendrite formation generally do little to optimize Coulombic efficiency and vice versa. Both problems become significantly worse with increasing current density and with repeated cycling as the morphological changes become amplified. Obvious mitigation strategies include reducing the current density by using large surface area metal anodes and suppressing dendrites using pulse charging. The in situ SEI is typically nonuniform in chemical composition through its depth and by being spatially heterogeneous laterally so that it is difficult to control morphological changes caused by the resulting nonuniform current distributions. Traditional approaches to minimizing Coulombic efficiency loss focus on optimizing the in situ SEI in terms of its stability, uniformity, flexibility, i.e., by varying solvent, salt, or additivies (e.g., tributylamine, CO2, SO2, vinylene carbonate). It has, however, proven difficult to achieve a completely satisfactory solution to the dendrite problem when using liquid electrolytes, and most of the common functional additives are likely incompatible with Li−O2 electrochemistry. An ingenious scheme to prevent dendrite formation in liquid electrolytes has been proposed by Ding et al.86 They add small concentrations of Rb+ or Cs+ to the electrolyte such that their reduction potentials are lower than that of the Li+. The Rb+ or Cs+ ions in the electrolyte are attracted to the high field regions of any developing Li asperities and therefore function as a selfhealing electrostatic shield as shown schematically in Figure 26.

Figure 26. Illustration of Li deposition in the presence of the selfhealing electrostatic shield mechanism. Reprinted with permission from ref 86. Copyright 2013 American Chemical Society.

At low current densities, this approach did effectively reduce dendrites, but there is no evidence that it will work as well at higher current densities. This approach also does not appear to solve the low Coulombic efficiency problem with Li metal. In addition, it is unlikely to be applicable in Li−air batteries since the reactions Rb−O2 or Cs−O2 at the cathode will remove these ionic species from solution. 7.1. Polymer Electrolytes

There has been much research into using polymer electrolytes for Li metal since many of these are stable with respect to interfacial reaction with Li metal.87 As a result the Coulombic efficiency remains high with cycling. If dendrites could be prevented with these polymers as well, they could provide a practical solution to Li metal cycling. Furthermore, polymers would inhibit, although not enitrely block, O2 crossover to the Li metal compared to liquid electrolytes. Unfortunately, typical polymer/salt electrolyte formulations, such as the common PEO/LiTFSI system, do not block dendrite formation.88 However, research on polymer electrolyte design has identified two interesting avenues to potentially reduce or eliminate dendrite formation. Monroe and Newman89 have proposed that Li dendrite formation can be prevented if the electrolyte or an artificial SEI has a shear modulus approximately twice that of the Li metal. The mechanical properties of PEO can be enhanced by using cross-linked polymer networks,90 although cross-linked PEO still has an insufficient modulus for dendrite suppression. In an elegant attempt to improve cycling life of Li metal batteries, segmented block copolymers, with a soft PEO block imparting Li+ conductivity and a hard polystyrene block imparting mechanical stiffness to suppress dendrite formation, have 11744


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been employed by Balsara and co-workers as electrolytes.91 Another polymer-based route to eliminate dendrites was also highlighted by Monroe and Newman,92 and predicts higher cycles to failure (and hence diminished dendrite growth) as the electrolyte’s Li+ transference number approaches unity, i.e., t+Li = 1. Therefore, research into developing polymer single ion conductors is important for Li metal batteries.93 Known polymer Li+ conductors with t+Li approaching unity are currently limited by low room temperature conductivity, even though significant efforts have vastly improved polymer electrolyte conductivity over the past few decades. 94 A further complication in the pursuit of polymers with t+Li = 1 is that materials design strategies to improve conductivity typically reduce t+Li. Nearly all research on Li+ polymer electrolytes has focused on Li−ion batteries employing either graphite or Li metal as an anode. Therefore, the effect of O2 on the Li metal/ polymer interface is unknown, but likely increases the propensity for dendrite formation as a result of increased SEI inhomogeneity.

is possible that regenerative braking requirements for fast charging may require higher C rate capability, and this would then have to be accommodated in a separate smaller battery/ capacitor with fast charging capability. We have discussed in the previous sections that although the fundamental electrochemistry seems excellent as the basis for a rechargeable Li−O2 battery, there are several important science/technical challenges hindering development. In this section, we first briefly mention some of the engineering challenges associated with developing a practical Li−air battery, and then discuss what specific energy and energy density are likely to be achievable in a practical Li−O2 or Li−air EV system (assuming that all but one of the science and engineering challenges can be overcome in an efficient way). The technical challenge that we do not think is feasible to ever overcome in the battery itself is the parasitic chemistry occurring with components in air other than O2 (CO2, H2O, etc.). Therefore, a Li−O2 EV battery will require either on board air scrubbing or closed O2 cycling, and either approach implies a significant lowering of specific energy and energy density from the theoretical ones. Several recent publications2b,100,101have addressed the issue of practical energy densities in far more depth than considered here. Although they differ in some assumptions and details, the basic conclusions are similar. Since Li−air batteries have not yet reached a practical design threshold, the engineering challenges at this stage are largely speculative. Nevertheless, some seem quite likely to become issues that will have to be overcome in designing a practical Li− air battery. Conversion of Li metal to Li2O2 implies a volume change of ∼30% in the battery between the fully charged state and the discharged state. This will be extremely difficult to manage, requiring very compliant seals on the Li anode and a cathode geometry that can accept the large Li2O2 deposit without adversely affecting the O2 and electrolyte flow patterns. Because the current/aerial density will be ∼100 times less than that in a PEMFC, ∼100 times higher cathode surface areas and gas flow fields relative to those of typical PEMFC will be required. Large gas flow fields that give uniform filling of the cathode (constant current densities within the cathode) and that minimize weight and volume need be designed. Electrolyte retention in an open system using air must be guaranteed, either by finding suitable electrolytes with low enough vapor pressure or using some kind of vapor barrier/trap that does not impede O2 access. In a closed system, the O2 + electrolyte mixture must not reach a flammable/explosion limit. Thermal management will be difficult given the large surface area of the cathode. In fact, the T limits of a Li−O2 battery are at present unknown and may be heavily constrained. At the low T end, both the surface kinetics (overpotentials) and charge transport (capacity) will be adversely affected, while at the high T end parasitic chemistry may become an even more significant issue. Therefore, it is unknown at this time whether Li−O2 batteries will ever be able to meet the T range in Table 3. The overall cell impedance must be ≤30 Ω to satisfy the needed power requirements. This means that if a Li metal protection membrane is used, it must be very thin (∼25 μm) and pinhole free. Finally, overall cost must be kept at the target level and safety ensured. To consider what system level constraints impose on practical Li−air densities, we assume the following: (1) Uchg does not rise significantly during the charge cycle, and cell impedances are ≤30 Ω so that the discharge−charge cycle is ∼85−90% efficient at practical current densities, (2) electrical

7.2. Solid State Electrolytes

An alternative approach to polymer electrolytes is to use inorganic solid state Li−ion conductors which have t+Li = 1 and a high mechanical strength to suppress dendrites. Currently, the most common are LiPON,95 LATP,96 garnet-like oxide glasses,97 and most recently Li10GeP2S2 superionic conductor.98 However, some are not stable to Li metal (LATP, Li10GeP2S2) and therefore require a buffer layer. To our knowledge, only LiPON can be deposited as a continuous thin film at this time, while the others are discs made from sintered powders. They therefore are relatively thick (>100 μm) to prevent pin holes and hence have only modest net Li+ ion conductance. This approach has been pioneered by Polyplus as their protected lithium electrode (PLE) using LATP.99 At present, we are not aware of any Li metal battery using an inorganic conductor with sustained cycling at >1 mA/cm2 current density. Therefore, significant future developments will be necessary for inorganic Li−ion conductors to satisfy Li−air needs.

8. SYSTEM LIMITATIONS FOR AN EV BATTERY Requirements for a Li−O2 battery system to roughly satisfy USCAR long-term commercialization goals for an EV battery pack are outlined in Table 3 from ref 100. Meeting usable energy and power goals would give >300 mile range at 250 C). However, barring these events, Li−air and metal−air batteries in general should have high safety (except for the Li metal) since they are relatively immune to the thermal runaway issues that are a concern for Li−ion batteries. Even internal shorts in a Li−air battery cannot cause currents higher than that allowed by the ORR of the cathode, and this is limited to the rate of O2 ingress.

9. SUMMARY AND OUTLOOK This review has tried to summarize the current status of research on nonaqueous Li−air batteries, but of course filtered through the authors’ biases. Because of the very high theoretical 11746


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In addition to the technical issues regarding the cathode electrochemistry, there are still well-known safety and Coulombic efficiency issues related to morphology changes in cycling Li metal as the anode. If more traditional anodes (lithiated graphite) are used instead, the Li−air battery loses much of its theoretical specific energy and is therefore much less attractive. To avoid contamination issues related to using air, either an air purification system or high pressure O2 tank must be carried on-board for use in an EV. These both have significant system penalties of weight and volume, so that the likely practical specific energy and energy density in use will be substantially less than theoretical ones. Although it is too early to specify a complete design, preliminary estimates suggest practical specific energies of ∼450−600 Wh/kg and energy densities of 450−550 Wh/L. While these are much smaller than the theoretical ones, they are still close enough to what is needed for a mass market EV battery if the technical challenges are overcome. Therefore, active research activities into overcoming the technical challenges of Li−O2 are definitely warranted, and that is why this review and the field in general are focusing principally on these challenges. We know from experience that some daunting technical challenges can be only one good idea away from a practical solution. Other battery chemistries in various initial stages of research and development will undoubtedly compete with Li−air for the EV battery of the future, and each of these now have their own set of technical challenges. For example, an advanced Li−ion battery using Li metal anodes and advanced (Li-rich) cathode material might give a practical specific energy of ∼300 Wh/kg and an energy density of 600−800 Wh/L. However, this also requires solving the technical challenges related to using Li metal as an anode and the capacity fade in Li-rich cathodes. Other potential high specific energy beyond Li−ion technologies, e.g., Li-sulfur or two electron Mg−ion intercalation batteries, are also being considered as future alternatives to Li− ion batteries (although with a possibly more limited energy density). Li−S batteries are currently at a more advanced stage of development than Li−O2 batteries, but two technical issues still plague them: (1) the elemental sulfur, like Li2O2, is also an electronic insulator, and (2) the Li−S cathode electrochemistry produces polysulfides soluble in most common organic electrolytes which hence migrate to the Li anode where they react.103 For Mg−ion intercalation batteries, current Mg battery electrolytes are corrosive and unstable at modest anodic potentials. In addition, identification of high potential reversible Mg intercalation compounds has proven difficult.104 Thus, all possible beyond Li−ion battery chemistries, including the socalled advanced Li−ion batteries, have technical challenges as well. It is therefore important to periodically assess the status of each battery chemistry’s research and development as we feel it is impossible to pick a winner at this stage given the uncertainty of ultimate success of the research and development for any of the beyond Li−ion battery chemistries. We have tried in this review to provide such a critical assessment of Li−air, albeit from our own perspective.

specific energy, there has been enormous activity and optimism in the past couple of years concerning the future development of Li−air for EV applications. Most of this early activity in the field is directed at trying to understand the electrochemistry and to understand/solve the fundamental science/technical issues that have become apparent in Li−O2 batteries rather than in the design of practical batteries. Because this heightened activity is only a few years old, there are still many unsolved mysteries and considerable debate as to interpretations of the work in this field. This is heightened by the fact that typical coulometry cycles seem very sensitive to impurities. In particular, Li−O2 batteries seem very sensitive to H2O impurities, and it appears that the overall mechanism of discharge and possibly charge changes when H2O is present. In an anhydrous Li−O2 battery, thin films of Li2O2 are deposited. However, with H2O present, large toroids of Li2O2 can be formed at low currents. The fundamental anhydrous Li−O2 electrochemistry appears excellent and could in principle form the basis of an excellent battery couple. Kinetic overpotentials for both discharge and charge are modest, although the battery impedance produces larger potential losses. However, there are still serious scientific/technical challenges to making even a laboratory Li−O2 battery that is rechargeable, combining high capacity with high power and with high cycle voltaic efficiency. Probably the most immediate concern is the stability of all battery components (electrolyte, cathode, and anode) that affects rechargeability. There are many claims of rechargeability over many discharge−charge cycles based simply on coulometry, but it is impossible to judge the truth of these claims from only the coulometry. Quantitative chemical criteria for rechargeability (YLi2O2, (e−/O2)dis, (e−/O2)chg, OER/ORR, etc.) indicate that electrolyte and cathode stability are not yet sufficient to give a single cycle with >90% rechargeability. There are a few reports in the recent literature where some of the quantitative chemical criteria for rechargeability have been combined with cycling and seem to indicate electrolyte/ cathode combinations where high rechargeability has been obtained. Unfortunately, we have not been able to reproduce these results in very similar but not quite identical cathode/ electrolyte combinations. The reasons for these differences are not yet known. The rise in the charging potential Uchg in galvanostatic discharge−charge cycles implies a low cycle voltaic efficiency. We have suggested that this arises from solid deposition at the electrolyte−Li2O2 interface, as the rate of rise is sensitive to CO2 and H2O contaminants in the headspace or electrolyte. Others suggest this is a change in the charging electrochemical mechanism during the charge part of the cycle that occurs with much higher overpotentials. Although quite high discharge capacities are possible at low discharge currents of

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