Unraveling the Complex Role of Iodide Additives in ... - ACS Publications

Jul 27, 2017 - potential of LiOH (3.42 V-alkaline and 3.82 V-neutral vs Li/. Li+), the newly .... target product and RM redox couple would be an ideal...
0 downloads 0 Views 7MB Size
Download PDF
Unraveling the Complex Role of Iodide Additives in Li−O2 Batteries

Yu Qiao,†,§ Shichao Wu,† Yang Sun,† Shaohua Guo,†,‡ Jin Yi,† Ping He,‡ and Haoshen Zhou*,†,‡,§ †

Energy Technology Research Institute, National Institute of Advanced Industrial Science and Technology (AIST), 1-1-1, Umezono, Tsukuba 305-8568, Japan ‡ National Laboratory of Solid State Microstructures & Department of Energy Science and Engineering, Nanjing University, Nanjing 210093, People’s Republic of China § Graduate School of System and Information Engineering, University of Tsukuba, 1-1-1, Tennoudai, Tsukuba 305-8573, Japan S Supporting Information *

ABSTRACT: Lithium iodide (LiI) has garnered considerable attention in aprotic Li−O2 batteries. However, the reaction mechanism is still under hot debate and is attracting increasing controversy due to contrasting observations. Herein, on the basis of thorough evidence, a relevant mechanism has been systematically illustrated. LiI has been revealed to promote the superoxide-related nucleophilic attack toward electrolyte by catalyzing the decomposition of peroxide intermediate, resulting in the accumulation of LiOH and other parasitic products. Also, they refuse to be oxidized by not only triiodide (I3−) but also iodine (I2), resulting in inevitable degradation. However, as a proton-donor, water can buffer the superoxide-related nucleophilic attack by reducing it to moderate hydroperoxide (HO2−). More importantly, the catalysis of iodide toward speroxide is restrained with the increase of alkalinity in water-contained electrolyte, resulting in the formation of Li2O2. Turning LiOH into Li2O2, the newly proposed mechanism leads to revolutionary reunderstanding toward the role of iodide and water in Li−O2 battery systems.

A

the related discussion rises to another upsurge because incredibly low charge potential and remarkably long cycle performance were achieved by the addition of water into iodide-contained aprotic Li−O2 batteries, in which LiOH became the ultimate discharge product.23 As such, the reaction process became more complicated because the role of water is another disputable topic, which still remains to be elucidated.24−27 On the basis of the current consensus and our recent study, the addition of water in an aprotic Li−O2 battery would not influence the formation of Li2O2 as the dominant product; thus, the mechanism for the formation of LiOH in an iodide-contained system becomes more and more confusing.19,28−31 Besides, more intensive controversy focuses on the iodide/triiodide (I3−)- and iodide/iodine (I2)-mediated oxidization reactions. Because the redox potential of the I−/I3− couple (∼3.0 V vs Li/Li+) lies below the decomposition potential of LiOH (3.42 V-alkaline and 3.82 V-neutral vs Li/ Li+), the newly proposed mechanism becomes thermodynamically unfavorable.32,33 Although several plausible factors (water

protic lithium−oxygen (Li−O2) batteries have attracted considerable attention due to its potential to provide remarkably high specific gravimetric energy compared to the state-of-the-art Li-ion batteries.1−4 However, there are several challenges seriously inhibiting the achievement of the reversible formation/decomposition of Li2O2, the dominant discharge product in aprotic system. Actually, the low energy efficiency and poor cycle ability can be mainly attributed to the parasitic reactions and high overpotential, which are tightly correlated with each other.5−7 As a soluble charge catalyst, redox mediators (RMs) have been extensively investigated due to their remarkable effect on reducing charge potential.8−12 The introduction of a RM into a charging process can effectively relieve the charge transport limitation by ingeniously transferring the decomposition of Li2O2 particles from a solid−solid interface to a liquid−solid interface.8,13−18 Among the numerous RMs, lithium iodide (LiI) received the most attention owing to its ordinary and low redox potential.14,19−22 However, further investigations demonstrated that the iodide would probably lead to side reactions during discharging, which is seriously against the “red line” of its feasibility as a RM, and not participate in the original oxygen reduction process.19,20 However, regardless of clarifying its influence toward oxygen reduction in nonaqueous conditions, © 2017 American Chemical Society

Received: May 30, 2017 Accepted: July 27, 2017 Published: July 27, 2017 1869

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

http://pubs.acs.org/journal/aelccp

Letter

ACS Energy Letters

Figure 1. Characterizations of the discharge products in Li−O2 cells with varying LiI contamination. (a) XRD patterns, (b) SEM images, (c) Raman, and (d) IR spectra of the cathodes harvested at the end of discharge, with a fixed 1.5 mAh cutoff capacity. The spectra are distinguished by different colors and offset for clarity. (e) The iodometric titrations quantify the deposited Li2O2. Combined with the acid/ base titrations, accurate amounts of LiOH are also shown for comparison. To avoid interference from the iodide in electrolytes, the discharged cathodes are further rinsed and dried before titration. (f) 1H NMR of D2O-extracted discharged products from both the cathode and separator (without rinsing). The spectrum of pristine electrolyte is shown as a reference. Carboxylates (formate and acetate) are quantified based on the benzene internal standard. Note that the proposed reaction scheme for electrolyte decomposition is shown inset.

on different charging plateaus remains contrasting among the previous reports.14,23,28−30 More importantly, the redox shuttling (water and/or soluble peroxides, etc.) between the cathode and lithium anode was poorly understood, resulting in the “charge” process suffering from an irrelevant iodide cycle, which would inevitably bring unnecessary misleading into the related discussion.12,28 In this case, without interference of the shuttle effect, the most controversial issues can be summarized in three points: (I) the influence of iodide toward LiOH formation during typical aprotic oxygen reduction; (II) the reactivity of oxidized iodide species (I3− and I2) toward the obtained Li2O2/LiOH in practical cells; and (III) the influence of additional water in an iodide-contained Li−O2 battery system, especially during the oxygen reduction process.

concentration, hydration enthalpies of Li salts, etc.) in this complex system would, to some extent, influence the redox potential of LiOH decomposition,34 related discussion did not draw any unanimous conclusion, in turn resulting in a more confusing divergence. Representatively, Zhu et al. indicated that I3− cannot directly react with Li2O2 powder and LiOH can be oxidized by I2 species with reversible oxygen evolution.30 Meanwhile, the reaction between Li2O2 and I3− species was strongly supported by Bruke et al., but they did not present any evidence for the formation of Li2O2 nor the reversible oxygen evolution in iodide-contained nonaqueous conditions. However, the decomposition of LiOH was primarily attributed to an I2-related reaction, resulting in irreversible formation of IO3−.28 Furthermore, the assignment of I−/I3− and I3−/I2 redox couples 1870

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

spectroscopy are employed to further qualitatively assess the parasitic products (Figure 1d). With the amount of carbonate species remaining almost unchanged, the adsorption peaks of formate (1364 cm−1) and acetate (1586 cm−1) present an obvious increase with the rising LiI concentration.39,40 More quantitative information on both Li2O2 and LiOH were obtained by iodometric and acid/base titrations (Figure 1e).24,28 The decreasing trend of Li2O2 species indicates that the typical aprotic oxygen reduction pathway has been essentially changed with the increase of LiI concentration, while the increase of hydroxide presents a proton-induced reaction pathway as the substitution. Furthermore, quantification of carboxylates (formate and acetate) is also demonstrated by NMR results (Figure 1f),41,42 in which the increasing trend of related carboxylate parasitic products is quantitatively reproved. Note that the proton from both hydroxide and carboxylates obtained in LiI-contained conditions would not be attributed to the trace amount of residual water in the aprotic electrolyte (less than 75 ppm in merely 40 μL of electrolyte); thus, the dominant source of proton would be rationally assigned to the decomposition of ether-based electrolytes, the only available proton source. According to previous reports, the decomposition of ether-based electrolyte has been mainly attributed to nucleophilic attack from the superoxide radical, resulting in subsequent oxidation decomposition into carboxylates, water, CO2, and so forth (shown as the inset scheme in Figure 1f).42,43 In the LiI-free typical condition, the majority of superoxide radicals would be further reduced into Li2O2,44,45 while a very limited amount of them tend to deposit as LiO2 or suffer a proton-reduction process into hydroperoxide (HO2−).19 However, once LiI is introduced, nucleophilic attack is greatly promoted due to rapid consumption of HO2− intermediate with the help of iodide catalyst,46 and this catalytic promotion has been further confirmed by the direct observation of IO− (Figure S5), an intermediate produced during the iodide-related catalysis procedure.18,20,47 The related reaction steps are listed as follows:

Clarifying these issues becomes the essential step for rationally assessing the availability of iodide as a RM in this complex system. In this study, redox shuttling has been effectively restrained by the employment of a solid electrolyte separator, and the above-mentioned hot-debated issues have been systematically unified: (I) During oxygen reduction, nucleophilic parasitic reactions would be further exacerbated by the iodide-related catalytic reaction toward hydroperoxide intermediate, resulting in accumulation of LiOH and other carboxylates. (II) Instead of employing commercial Li2O2 or LiOH powder as simulate targets, more rational evidence has been exhibited to identify the reactivity of oxidized iodide species within practical cell systems. Li2O2 can be decomposed by I3−, while LiOH refuses to be oxidized by I2. (III) Surprisingly, the additional water can restore the discharge product from irreversible LiOH back to Li2O2, promoting reversible oxygen evolution. The interesting “LiOH to Li2O2” transformation above is rationally attributed to the increasing alkalinity and related inactivation of the iodide-catalyzed hydroperoxide decomposition reaction. Thorough analysis and solid conclusions present in this study would lead to a revolutionary reunderstanding of the current controversial issues and provide deep insights into the development of both stable solvents and practical additives in Li−O2 and Li−air open battery systems. Inf luence of Iodide Additive toward the Aprotic Oxygen Reduction Reaction. Initially, we investigate the influence of various LiI concentrations toward a typical aprotic oxygen reduction process (Figure S1) with LiN(CF3SO3)2 (LiTFSI) in tetra-ethylene glycol dimethyl ether (TEGDME) electrolyte. The total salt concentration (LiTFSI + LiI) is kept at 1.0 mol/ L. In Figure S2, the discharge plateau present a trace rising, while the charge profiles change substantially with increasing LiI concentration (from 0 to 1000 mM), which will be further discussed in the next section. More importantly, the discharge chemistry presents an essential change (Figure 1). On the basis of the XRD patterns of cathodes discharged to 1.5 mAh capacity with various LiI concentrations (Figure 1a), only the crystalline Li2O2 phase can be observed in LiI-free condition (red trace), which is also consistent with the formation of typical toroidal Li2O2 particles shown in the corresponding scanning electron microscopy (SEM) image (Figure 1b).35,36 However, new diffraction peaks appear with the presence of LiI (orange trace, 10 mM), which can be indexed to crystalline LiOH (Figure S3), with the visualized appearance of several pieces of sheet inserted between Li2O2 toroids.37 With the continuously concentrated LiI (green trace, 100 mM), LiOH became dominant. At the same time, Li2O2 toroids totally disappeared, remaining on the surface densely occupied by a layer of flake-like LiOH species.29 Finally, the diffraction peaks of LiOH became sharper when the concentration of LiI was increased up to 1 M, coinciding with the growing size of LiOH flakes. Beyond the crystalline limitation of XRD characterization, more comprehensive information on discharge products has been revealed by adsorption spectroscopy. As shown in Raman spectra (Figure 1c), regardless of the gradual substitution from Li2O2 to LiOH with the increasing LiI concentration, the disappearance of LiO2-related peaks (at 1132 and 1498 cm−1) becomes noteworthy.38 Combine with the increasing trend of related parasitic products (941 and 1042 cm−1, Figure S4), the decrease of the superoxide intermediate indicates that the LiI additive would exacerbate relevant nucleophilic side reactions during discharge. Moreover, IR

solvent molecule + O2− → solvent molecule radical + HO2− HO2− + I− + Li+ → IO− + LiOH IO− + HO2− + Li+ → O2 + I− + LiOH total: 2HO2− + 2Li+ → O2 + 2LiOH (iodide catalysis)

In summary, on the basis of direct experimental evidence, the formation of LiOH during iodide-contained aprotic oxygen reduction has been first ascribed to an iodide-related catalysis effect toward hydroperoxide intermediate. Furthermore, more solid evidence for the related catalysis in the aprotic Li−O2 cell has been demonstrated with MnO2 additive (Figure S6), another classic catalysis toward peroxides. The catalysis toward hydroperoxide intermediate tips the balance from Li2O2 formation to LiOH accumulation during oxygen reduction, meanwhile accelerating the parasitic decomposition of electrolyte. This essential variation on the reaction pathway has been built at the expense of consuming expensive electrolyte, which would directly lead to a serious query regarding the availability of LiI additive. Mediated Reactivity of Specific Oxidized Iodide Species toward Obtained Li2O2/LiOH. The center assessment criteria for RMs in Li−O2 batteries undoubtedly focus on the mediated reactivity of a related redox couple toward the oxidization of corresponding target discharge products. Generally, directly 1871

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

Figure 2. Assessments of the mediated activity for I−/I3− and I3−/I2 redox couples toward the oxidization of Li2O2/LiOH. (a) Schematic and photo of the modified Li−O2 pouch cell. After discharge (deposited with Li2O2 and/or LiOH), the cathode is divided in half and disconnectedly reassembled into the new pouch cell with 1 M LiI electrolyte. The subsequent charging process is performed on a single side (electro-side), with the other side remaining without electrochemical connection (free-side). Because both sides share a single separator, the oxidized iodide species would gradually diffuse to the free-side during charging. (b,c) The related cathodes are deposited with a mixture of LiOH/Li2O2 (discharged in 10 mM LiI) and LiOH (discharged in 1 M LiI), and the charging processes are conducted with a cutoff voltage at 3.3 and 3.7 V, respectively. The photos of the corresponding separator after charging are shown in the inset. Raman spectra of both the separator (electrolyte) and cathode are present below the related voltage profiles, and the spectra collected at different positions on both sides are distinguished by colors and offset for clarity. The related positions are marked in the scheme and photo.

saturated golden yellow on the electro-side separator (left half, points A−E). However, due to the diffusion of I3−, the yellow color gradually turns faint on the free-side (right half, points F− J). Moreover, concentration gradient distribution of I3− species has been obtained by Raman spectra collected from the specific positions on the separator, with a pair of isolated peaks (117.5 and 145.6 cm−1) assigned to I3− species.48 More importantly, on the corresponding cathode surface, regardless of the D-band and other peaks of parasitic products, the Raman peak of Li2O2 totally disappeared on the electro-half (points A−E), while the peak intensity assigned to LiOH species remained unchanged after charging. Besides, on the neighboring free-side (points F− J), Li2O2 only remains on the specific I3−-free area, with LiOH remaining almost unchanged compared to the discharge state (bottom black trace). In this case, combined with the soloLi2O2 comparison group (Figure S8), the mediated reactivity of I3− toward Li2O2 has been confirmed, while LiOH refuses to be decomposed on an I−/I3−-related plateau (3.05 V vs Li/Li+), which also coincides well with each of their thermodynamic equilibrium potentials.32,33 Another disputed issue focuses on the higher I3−/I2 plateau (3.5 V vs Li/Li+), on which the mediated reactivity of I2 toward LiOH has been investigated on the cathode discharged with 1 M LiI (deposited with LiOH and other carboxylates

mixing the target product (commercial Li2O2 or LiOH) into RM-contained electrolyte with its specific oxidized state (RMn+) is a simple and visualized method to estimate the reaction activity;30 however, both the accurate amount and concentration of corresponding reactants present a large difference from the environment of a practical Li−O2 cell, which would inevitably influence the rationality of related discussion. Thus, reverting the reaction back to a practical cell system and simultaneously collecting information on both the target product and RM redox couple would be an ideal way to confirm the mediated activity. Combining the employment of our newly designed Li−O2 pouch cell and Raman spectroscopy (Figure 2), the mediated reactivity of oxidized iodide (triiodide: I3−; iodine: I2) toward discharge products (Li2O2 and LiOH) has been investigated. After assigning the corresponding iodide-related redox couples to the relevant charging plateaus (Figure S7), we first focus on the mediated reactivity of the I−/I3− redox couple. As the target discharge products, both Li2O2 and LiOH deposit on the cathode after being discharged to 1.5 mAh capacity in the Li−O2 cell with 10 mM LiI additive (orange trace in Figure 1), and the subsequent charging process is performed with fresh 1 M LiI-based electrolyte (Figure 2b). After cutoff at the end of the I−/I3− plateau, I3− was homogeneously scattered, with the 1872

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

Figure 3. Characterizations of the discharge products in Li−O2 cells with 1 M LiI additive and varying water contamination. (a) Voltage profiles of Li−O2 cells with 1 M LiI concentration and varying additional water contamination. The cells are fixed at a 1.5 mAh cutoff discharge capacity and 3.7 V cutoff charge potential. (b) SEM images, (c) XRD patterns, and (d) Raman spectra of the cathodes harvested from corresponding cells at the end of the discharge process. (e) The acid/base and iodometric titrations conducted on the discharged cathodes quantitatively estimate the deposited LiOH and Li2O2, respectively. (f) IR spectra performed on the electrolytes collected from the corresponding cell at the end of each discharge process.

Role of Water in an Iodide-Contained Li−O2 Battery: Turning LiOH Back into Li2O2. After comprehensively verifying the role of LiI additive in an aprotic Li−O2 battery system, our study was extended to investigate the influence of water on a LiIcontained Li−O2 battery system. The motivation was not only restricted to the widely controversial topic but also focused on the specific role of additional water, the essential factor for the evolution from a Li−O2 battery to a more practical Li−air open battery system. Therefore, due to the complexity and significance of the above-mentioned system, rigorously elucidating the reaction mechanism with sufficient solid evidence became remarkably necessary (Figures 3 and 4). As shown in Figure 3a, during charging, the I−/I3−-related plateau was stretched with the increasing additional water content, while the I3−/I2-related plateaus remaining unchanged. Besides, the morphology of discharge products suffer from a visualized change (Figure 3b). In detail, the disordered flakelike products gradually turn to gather together, resulting in the formation of flower-like big particles, which are well-organized by many thick plates layer by layer under 30% water conditions (v/v %). Surprisingly, on the basis of the XRD (Figure 3c) and Raman (Figure 3d) observations, the discharge products are gradually replaced by Li2O2 with water addition. This transformation has also been quantitatively confirmed by titration (Figure 4e), in which the amount of LiOH (blue bar) continuously decreases, while, on the contrary, Li2O2 (red bar) gradually increases with additional water content. It is worth noting that the transition from LiOH to Li2O2 via water

byproducts, Figure 2c). At the end of the iodine-related plateau, Raman peaks of I3− species totally disappear on the electro-side, and the color of the electrolyte (infiltrated into the separator) turns brown, which indicates further oxidization from I3− to I2. However, both obtained LiOH and other parasitic products refuse to be oxidized on the I3−/I2 plateau. Therefore, the entire charging process turns out to be a pure Li-iodide electrochemistry process without any oxygen evolution (as will be proved by DEMS later). Therefore, the I−/I3− redox couple acts as an eligible mediator toward the decomposition of Li2O2, while LiOH rejects being decomposed even at a higher I3−/I2 potential plateau. Moreover, taking the negative role of LiI toward aprotic oxygen reduction into consideration, during a typical cycle (Figure S2), both the LiOH and carboxylates formed during discharging still remain on the cathode after charging (Figures S9 and S10), while the superfluous oxidized iodide species (I3− or I2) are reduced in the subsequent discharge process. During cycling (Figure S11−S14), although the charge overpotential can be illusively controlled by LiI additive among the initial several cycles, the major active-site blocker merely turns from carbonate (iodide-free condition) to hydroxide (iodide-contained condition), which also leads to cell degradation during cycling. In another words, instead of the LiOH-related reversible aprotic Li−O2 battery system, the nature of the electrochemical cycling has been proved as the Liiodine redox reactions. 1873

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

Figure 4. Segmented study of different discharge−charge stages performed in a Li−O2 cell with 1 M LiI and 5% additional water. (a) Voltage profile with the selected stages marked by specific colors. The photos inset present the separator collected during charging, which reflects the color change of LiI-contained electrolyte at corresponding charge stages: I− (no color), I3− (golden yellow), and I2 (brown), respectively. (b,c) Raman spectra recorded on the dried cathodes and corresponding wet cathodes harvested at the selected points. The spectra of the wet cathode (without rinsing) demonstrate additional information from residual adsorption species (HO2− et al.), which are highlighted by a yellow frame. (d,e) IR and Raman spectra of electrolytes collected at related points. (f) Oxygen evolution rate during charging versus capacity. Besides the 5% water contained condition (green trace), O2 evolution in nonaqueous (red trace) and 30% water (purple trace) conditions is also presented for comparison.

efficiently controlled, restraining the negative LiOH pathway (dominant in LiI-contained nonaqueous conditions), meanwhile promoting the formation of Li2O2. In order to further prove this conclusion, segmented analysis (3 different stages, Figure 4a) has been performed in a Li−O2 cell with 1 M LiI and 5% (v/v %) additional water. At the end of each stage, deposited products (Figure 4b), surface adsorptions (Figure 4c), soluble peroxide/hydroxide (Figure 4d), and oxidized iodide in electrolyte (Figure 4e) have been systematically studied by comprehensive spectroscopic investigations. During the first discharge stage (point A to B), combined with water, O2 accepts two electrons and is reduced to both hydroperoxide (HO2−) and hydroxide via a conventional “oxygen−peroxide” redox reaction.49 Once produced, hydroperoxide suffers from a rapid decomposition by iodiderelated catalysis.47 The consumption of hydroperoxide further promotes the initial oxygen reduction step and the accumulation of LiOH, resulting in the hydroxide as the only observable product both on the cathode and in electrolyte during this stage. The related sets of reactions are illustrated as follows:

addition seems quite against common knowledge because the water would undoubtedly introduce an additional proton, which would tend to promote the formation of LiOH. Actually, on the contrary, introducing water into an aprotic Li−O2 battery system cannot lead to the transformation from Li2O2 to LiOH, although water itself is a good proton donor.24,25 Enlightened by the newly proposed role of water,31 the iodidecatalyzed reaction pathway here might be essentially changed by the introduction of water. On the basis of the IR spectra collected from discharged electrolytes (Figure 3f), the concentration of H2O2 gradually increases with increasing water content, which indicates that iodide-related catalysis toward peroxide species has been effectively restrained. Also, this variation can be attributed to the increase of hydroxide (alkalinity) dissolved into discharged electrolytes (Figure 3f). Besides, this catalysis inactivation tendency can also be proved by the reduction of related parasitic products observed in Raman spectra (Figure 3d), which has been assigned to an iodide-catalyzed electrolyte decomposition process during oxygen reduction. Thus, with the increasing concentration of OH− in electrolyte, catalysis of iodide species toward soluble peroxide species would be 1874

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

Figure 5. Proposed reaction mechanism in iodide-contained Li−O2 battery systems. (a) Aprotic Li−O2 cell with LiI additive. (b) Li−O2 cell with both iodide additive and additional water.

O2 + H 2O + 2e− → HO2− + OH−

O2 + H 2O + 2e− → HO2− + OH−

2Li+ + 2HO2− → 2LiOH + O2 (iodide catalysis)

HO2− + H 2O → H 2O2 + OH−

overall reaction: O2 + 2H 2O + 4Li+ + 4e− → 4LiOH

H 2O2 + 2OH− + 2Li+ → Li 2O2 + 2H 2O overall reaction: O2 + 2Li+ + 2e− → Li 2O2

Although the formation of LiOH presented here is very similar to the one observed in aprotic conditions, the essential difference focuses on the source of the proton. Instead of destructively taking the proton from decomposed electrolyte in an aprotic system, herein, water became the prior proton donor and buffered the nucleophilic attack from superoxide, which is definitely a positive revolution for the entire system. Subsequently, with the increase of OH− concentration in electrolyte, the catalysis of iodide (toward hydroperoxides) gradually became inactived (point B to D). On the basis of Raman evidence (point C, Figure 4b,c), HO2− became stable under alkaline circumstances as an adsorption product (Figure S16). Continuously, HO2− chemically reacts with water and turns to a more stable soluble peroxide species, H2O2, which has also been proved by the IR spectra of the electrolyte (point C, Figure 4d). With the accumulation of H2O2, combining with hydroxide in the electrolyte, Li2O2 is chemically produced (point D, Figure 4b,c), accompanied by the restoration of water. The related equilibrium would keep moving forward with the precipitation of Li2O2. Thus, the total reaction can be regarded as a 2e−/O2 reaction with Li2O2 as the final deposited product:

Note that the Li2O2-formation mechanism was further confirmed during a relevant iodide-free water-contained oxygen reduction process (Figure S17). Besides, the newly proposed discharge mechanism has also been proved by the employment of MnO2 as another conventional catalyst toward HO2−/H2O2 (Figure S18). However, the essential difference is that once the catalysis refuses to be restrained (MnO2-based solid catalyst), as performed in iodide conditions, the accumulation of LiOH would continue, without being altered to the Li2O2 pathway. During the initial stage of charging (point D to E), peroxiderelated characterization peaks totally disappear in corresponding adsorption spectra, which indicates the complete decomposition of peroxide with the help of the I−/I3− redox couple. On the basis of the results of quantitative differential electrochemical mass spectrometry (DEMS), the oxidization of Li2O2 has been confirmed as a typical 2e−/O2 process (Figure 4f).41 During the subsequent charging performed on the I−/I3− plateau (point E to F), without further oxygen evolution, I3− starts to accumulate in electrolyte with the appearance of related peaks in Raman spectra (point F, Figure 4e), which is also consistent with the golden yellow present on the infiltrated electrolyte (point F, inset Figure 4a). Finally, with the disappearance of I3−-related Raman peaks, the iodide-contained electrolyte turns brown at point G, which indicates the further oxidization of I3− into I2 species (Figure S7). However, a certain 1875

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

molecules, resulting in LiOH and carboxylates as the dominant discharge products. (2) During charging, Li2O2 can undergo mediated oxidization by the I−/I3− redox couple, while LiOH refuses to decompose even at the I3−/I2-related potential plateau. (3) As a proton donor, water transforms the oxygen reduction process into the formation of hydroperoxide and hydroxide. Also acting as a hydroxyl donor, water promotes the deactivation of iodide due to the increasing alkalinity, resulting in the survival of hydroperoxide, which directly leads to a transformation from the initial LiOH-deposition reaction pathway to a subsequent Li2O2-formation process. In summary, although the I−/I3− redox couple acts as a competent mediator toward the oxidization of Li2O2, the inevitable iodide-promoted electrolyte decomposition and the formation of irreversible parasitic products still remain serious challenges for its application in practical Li−O2 battery systems. However, instead of solvent molecules, water can provide an additional proton to buffer the severely nucleophilic aggressiveness of superoxide by transforming it into a relatively moderate hydroperoxide. More importantly, the hydroxyl from water inactivates the catalysis of iodide toward peroxide intermediates, transforming the formation of irreversible LiOH to reversible Li2O2. Consequently, the positive role of water has been demonstrated, which is essentially helpful to improve the reversibility and remedy the intrinsic defect of iodide additive. The mechanism demonstrated in this study would not only present a rational explanation for the current relevant controversy but also provide a series of key information for the design route of electrolyte solvents and additives that afford more practical Li−O2 batteries.

amount of LiOH remains undecomposed at the end of charging due to its inert nature toward both I−/I3− and I3−/I2 redox couples. As a comparison, in nonaqueous conditions, no oxygen evolution is present during charging (red trace, Figure 4f) because LiOH refuses to be oxidized as the dominant discharge product. However, in 30% water conditions, the oxygen evolution plateau also becomes longer on the corresponding extended I−/I3− redox plateau (purple trace, Figure 4f), which is consistent with the increasing proportion of obtained Li2O2. Furthermore, characterizations toward charging products are also fit well with the newly proposed reaction mechanism (Figures S19−S22). Consequently, related reaction processes have been comprehensively summarized (Figure 5). In nonaqueous systems, as an intermediate of superoxide-related nucleophilic attack toward electrolyte, hydroperoxide (HO2−) suffers from rapid decomposition by the catalysis of iodide. This catalysis effect, in turn, promotes further nucleophilic attack, resulting in accumulation of LiOH and carboxylates. Subsequently, the irreversible nature of LiOH toward oxidized iodide species drives the charging process to turn in to a pure iodide oxidization process, without any oxygen evolution. The LiOH pathway conducted in an aprotic Li−O2 cell can be essentially regarded as an iodide-promoted electrolyte decomposition process with irreversible oxygen reduction. However, in the iodide-contained Li−O2 cell with additional water, at the initial stage of oxygen reduction, O2 is reduced to hydroperoxide species, with water acting as a proton donor and a buffer, which effectively transfers the target of superoxide-related nucleophilic attack from the electrolyte solvent to water molecule. Meanwhile, with the catalysis of iodide, hydroperoxide suffers from rapid decomposition, promoting the formation and deposition of LiOH. Subsequently, the catalytic activity of iodide is gradually restrained with the increasing hydroxide concentration in electrolyte, resulting in the presence of hydroperoxide as a stable adsorption. Herein, water turns to act as a hydroxyl donor and inactivate the iodide catalyst. Following with a series of relevant chemical reactions, Li2O2 precipitates on the cathode surface with the release of the recycled water. Note that this LiOH (4e−/O2) to Li2O2 (2e−/ O2) transformation step seems to coincide well with Burke et al.;28 however, the essential difference focuses on the reason for this amazing “LiOH to Li2O2” transition, which should not be attributed to exhaustion of water but be assigned to the waterpromoted Li2O2 formation pathway. Although the Li-iodide electrochemistry is continuously conducted during the end of the charge process with the successive production of I3− and I2 species, peroxide species, as the dominant discharge products, have undergone mediated oxidization with the help of the I−/ I3− redox couple on the 3.0 V potential plateau with the reversible evolution of oxygen. Herein, beyond the traditional acknowledgment of the role of water during oxygen reduction, the function of water has been largely developed into both a reducing parasitic reaction and promoting reversible oxygen reduction/evolution. In summary, as one of the most controversial issues in Li−O2 batteries, the role of LiI additive and the influence of additional water have been systematically investigated by the employment of a set of comprehensive spectroscopic analyses. The obtained conclusions can be summarized as follows: (1) In nonaqueous conditions, by catalyzing the decomposition of hydroperoxide (HO2−) intermediate, iodide additive exacerbates the nucleophilic attack of superoxide toward ether-based solvent

■ ■

EXPERIMENTAL METHODS See the Supporting Information for details on the experimental procedures. ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsenergylett.7b00462. Experimental details and additional results and discussions (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] or [email protected]. ORCID

Yu Qiao: 0000-0002-2191-3875 Jin Yi: 0000-0001-6203-1281 Ping He: 0000-0002-1498-8203 Haoshen Zhou: 0000-0001-8112-3739 Author Contributions

Y.Q. and H.Z. contributed to the design of the research. Y.Q. conducted the spectroscopy, titration, DEMS, and NMR characterizations and performed the experimental data analysis. Y.Q. and S.W. performed the XRD, SEM, and measurements. All authors cowrote the manuscript. H.Z. supervised the work. All authors discussed the results and commented on the manuscript. Notes

The authors declare no competing financial interest. 1876

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters

■

Reduction of Superoxide by Viologen. Chem. Commun. 2015, 51, 1705−1708. (18) Kwak, W.-J.; Hirshberg, D.; Sharon, D.; Afri, M.; Frimer, A. A.; Jung, H.-G.; Aurbach, D.; Sun, Y.-K. Li-O2 Cells with Libr as an Electrolyte and a Redox Mediator. Energy Environ. Sci. 2016, 9, 2334− 2345. (19) Kwak, W. J.; Hirshberg, D.; Sharon, D.; Shin, H. J.; Afri, M.; Park, J. B.; Garsuch, A.; Chesneau, F. F.; Frimer, A. A.; Aurbach, D.; et al. Understanding the Behavior of Li-Oxygen Cells Containing Lii. J. Mater. Chem. A 2015, 3, 8855−8864. (20) Landa-Medrano, I.; Olivares-Marín, M.; Pinedo, R.; Ruiz de Larramendi, I.; Rojo, T.; Tonti, D. Operando Uv-Visible Spectroscopy Evidence of the Reactions of Iodide as Redox Mediator in Li−O2 Batteries. Electrochem. Commun. 2015, 59, 24−27. (21) Yu, M.; McCulloch, W. D.; Beauchamp, D. R.; Huang, Z.; Ren, X.; Wu, Y. Aqueous Lithium−Iodine Solar Flow Battery for the Simultaneous Conversion and Storage of Solar Energy. J. Am. Chem. Soc. 2015, 137, 8332−8335. (22) Zhang, T.; Liao, K.; He, P.; Zhou, H. A Self-Defense Redox Mediator for Efficient Lithium-O2 Batteries. Energy Environ. Sci. 2016, 9, 1024−1030. (23) Liu, T.; Leskes, M.; Yu, W. J.; Moore, A. J.; Zhou, L. N.; Bayley, P. M.; Kim, G.; Grey, C. P. Cycling Li-O2 Batteries Via Lioh Formation and Decomposition. Science 2015, 350, 530−533. (24) Aetukuri, N. B.; McCloskey, B. D.; Garcia, J. M.; Krupp, L. E.; Viswanathan, V.; Luntz, A. C. Solvating Additives Drive SolutionMediated Electrochemistry and Enhance Toroid Growth in NonAqueous Li-O2 Batteries. Nat. Chem. 2014, 7, 50−56. (25) Schwenke, K. U.; Metzger, M.; Restle, T.; Piana, M.; Gasteiger, H. A. The Influence of Water and Protons on Li2o2 Crystal Growth in Aprotic Li-O2 Cells. J. Electrochem. Soc. 2015, 162, A573−A584. (26) Kwabi, D. G.; Batcho, T. P.; Feng, S.; Giordano, L.; Thompson, C. V.; Shao-Horn, Y. The Effect of Water on Discharge Product Growth and Chemistry in Li-O2 Batteries. Phys. Chem. Chem. Phys. 2016, 18, 24944−24953. (27) Li, F. J.; Wu, S. C.; Li, D.; Zhang, T.; He, P.; Yamada, A.; Zhou, H. S. The Water Catalysis at Oxygen Cathodes of Lithium-Oxygen Cells. Nat. Commun. 2015, 6, 7843. (28) Burke, C. M.; Black, R.; Kochetkov, I. R.; Giordani, V.; Addison, D.; Nazar, L. F.; McCloskey, B. D. Implications of 4 E− Oxygen Reduction Via Iodide Redox Mediation in Li−O2 Batteries. ACS Energy Letters 2016, 1, 747−756. (29) Li, Z.; Ganapathy, S.; Xu, Y.; Heringa, J. R.; Zhu, Q.; Chen, W.; Wagemaker, M. Understanding the Electrochemical Formation and Decomposition of Li2o2 and Lioh with Operando X-Ray Diffraction. Chem. Mater. 2017, 29, 1577−1586. (30) Zhu, Y. G.; Liu, Q.; Rong, Y.; Chen, H.; Yang, J.; Jia, C.; Yu, L.J.; Karton, A.; Ren, Y.; Xu, X.; et al. Proton Enhanced Dynamic Battery Chemistry for Aprotic Lithium−Oxygen Batteries. Nat. Commun. 2017, 8, 14308. (31) Qiao, Y.; Wu, S.; Yi, J.; Sun, Y.; Guo, S.; Yang, S.; He, P.; Zhou, H. From O2− to Ho2−: Reducing by-Products and Overpotential in LiO2 Batteries by Water Addition. Angew. Chem., Int. Ed. 2017, 56, 4960−4964. (32) Shen, Y.; Zhang, W.; Chou, S.-L.; Dou, S.-X. Comment on “Cycling Li-O2 Batteries via LiOH Formation and Decomposition. Science 2016, 352, 667. (33) Viswanathan, V.; Pande, V.; Abraham, K. M.; Luntz, A. C.; McCloskey, B. D.; Addison, D. Comment on “Cycling Li-O2 Batteries Via LiOH Formation and Decomposition. Science 2016, 352, 667. (34) Liu, T.; Kim, G.; Carretero-González, J.; Castillo-Martínez, E.; Bayley, P. M.; Liu, Z.; Grey, C. P. Response to Comment on “Cycling Li-O2 Batteries Via LiOH Formation and Decomposition. Science 2016, 352, 667. (35) Lu, Y. C.; Kwabi, D. G.; Yao, K. P. C.; Harding, J. R.; Zhou, J. G.; Zuin, L.; Shao-Horn, Y. The Discharge Rate Capability of Rechargeable Li-O2 Batteries. Energy Environ. Sci. 2011, 4, 2999−3007. (36) Adams, B. D.; Radtke, C.; Black, R.; Trudeau, M. L.; Zaghib, K.; Nazar, L. F. Current Density Dependence of Peroxide Formation in

ACKNOWLEDGMENTS This work was partially supported by SPRING-ALCA from the Japan Science and Technology Agency (JST). Financial support from the National Basic Research Program of China (2014CB932300) and NSF of China (21373111 and 21633003) are acknowledged. Y.Q. acknowledges a scholarship from the China Scholarship Council (CSC). The authors thank Ms. Yang Liu (AIST) and Mr. Sixie Yang (Nanjing Univ.) for their help on UV−vis and DEMS characterization.

■

REFERENCES

(1) Luntz, A. C.; McCloskey, B. D. Nonaqueous Li-Air Batteries: A Status Report. Chem. Rev. 2014, 114, 11721−11750. (2) Lu, J.; Li, L.; Park, J. B.; Sun, Y. K.; Wu, F.; Amine, K. Aprotic and Aqueous Li-O2 Batteries. Chem. Rev. 2014, 114, 5611−5640. (3) Grande, L.; Paillard, E.; Hassoun, J.; Park, J. B.; Lee, Y. J.; Sun, Y. K.; Passerini, S.; Scrosati, B. The Lithium/Air Battery: Still an Emerging System or a Practical Reality? Adv. Mater. 2015, 27, 784− 800. (4) Aurbach, D.; McCloskey, B. D.; Nazar, L. F.; Bruce, P. G. Advances in Understanding Mechanisms Underpinning Lithium−Air Batteries. Nat. Energy 2016, 1, 16128. (5) McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Mori, T.; Scheffler, R.; Speidel, A.; Sherwood, M.; Luntz, A. C. Limitations in Rechargeability of Li-O2 Batteries and Possible Origins. J. Phys. Chem. Lett. 2012, 3, 3043−3047. (6) McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshoj, J. S.; Norskov, J. K.; Luntz, A. C. Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li-O2 Batteries. J. Phys. Chem. Lett. 2012, 3, 997−1001. (7) Adams, B. D.; Black, R.; Williams, Z.; Fernandes, R.; Cuisinier, M.; Berg, E. J.; Novak, P.; Murphy, G. K.; Nazar, L. F. Towards a Stable Organic Electrolyte for the Lithium Oxygen Battery. Adv. Energy Mater. 2015, 5, 1400867. (8) Chen, Y. H.; Freunberger, S. A.; Peng, Z. Q.; Fontaine, O.; Bruce, P. G. Charging a Li-O2 Battery Using a Redox Mediator. Nat. Chem. 2013, 5, 489−494. (9) McCloskey, B. D.; Addison, D. A Viewpoint on Heterogeneous Electrocatalysis and Redox Mediation in Nonaqueous Li-O2 Batteries. ACS Catal. 2017, 7, 772−778. (10) Pande, V.; Viswanathan, V. Criteria and Considerations for the Selection of Redox Mediatiors in Nonaqueous Li−O2 Batteries. ACS Energy Letters 2017, 2, 60−63. (11) Liang, Z.; Lu, Y.-C. Critical Role of Redox Mediator in Suppressing Charging Instabilities of Lithium−Oxygen Batteries. J. Am. Chem. Soc. 2016, 138, 7574−7583. (12) Qiao, Y.; Ye, S. Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions with Tetrathiafulvalene as a Redox Mediator in Li−O2 Battery. J. Phys. Chem. C 2016, 120, 15830− 15845. (13) Bergner, B. J.; Schurmann, A.; Peppler, K.; Garsuch, A.; Janek, J. Tempo: A Mobile Catalyst for Rechargeable Li-O2 Batteries. J. Am. Chem. Soc. 2014, 136, 15054−15064. (14) Lim, H. D.; Song, H.; Kim, J.; Gwon, H.; Bae, Y.; Park, K. Y.; Hong, J.; Kim, H.; Kim, T.; Kim, Y. H.; et al. Superior Rechargeability and Efficiency of Lithium-Oxygen Batteries: Hierarchical Air Electrode Architecture Combined with a Soluble Catalyst. Angew. Chem., Int. Ed. 2014, 53, 3926−3931. (15) Sun, D.; Shen, Y.; Zhang, W.; Yu, L.; Yi, Z. Q.; Yin, W.; Wang, D.; Huang, Y. H.; Wang, J.; Wang, D. L.; et al. A Solution-Phase Bifunctional Catalyst for Lithium-Oxygen Batteries. J. Am. Chem. Soc. 2014, 136, 8941−8946. (16) Kundu, D.; Black, R.; Adams, B.; Nazar, L. F. A Highly Active Low Voltage Redox Mediator for Enhanced Rechargeability of Lithium−Oxygen Batteries. ACS Cent. Sci. 2015, 1, 510−515. (17) Yang, L.; Frith, J. T.; Garcia-Araez, N.; Owen, J. R. A New Method to Prevent Degradation of Lithium-Oxygen Batteries: 1877

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878

Letter

ACS Energy Letters the Li-O2 Battery and Its Effect on Charge. Energy Environ. Sci. 2013, 6, 1772−1778. (37) Wu, S.; Tang, J.; Li, F.; Liu, X.; Zhou, H. Low Charge Overpotentials in Lithium-Oxygen Batteries Based on Tetraglyme Electrolytes with a Limited Amount of Water. Chem. Commun. 2015, 51, 16860−16863. (38) Zhai, D. Y.; Wang, H. H.; Lau, K. C.; Gao, J.; Redfern, P. C.; Kang, F. Y.; Li, B. H.; Indacochea, E.; Das, U.; Sun, H. H.; et al. Raman Evidence for Late Stage Disproportionation in a Li-O2 Battery. J. Phys. Chem. Lett. 2014, 5, 2705−2710. (39) Ottakam Thotiyl, M. M.; Freunberger, S. A.; Peng, Z. Q.; Bruce, P. G. The Carbon Electrode in Nonaqueous Li-O2 Cells. J. Am. Chem. Soc. 2013, 135, 494−500. (40) Qiao, Y.; Ye, S. Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions on Carbon Electrodes in Li−O2 Battery. J. Phys. Chem. C 2016, 120, 8033−8047. (41) McCloskey, B. D.; Valery, A.; Luntz, A. C.; Gowda, S. R.; Wallraff, G. M.; Garcia, J. M.; Mori, T.; Krupp, L. E. Combining Accurate O2 and Li2o2 Assays to Separate Discharge and Charge Stability Limitations in Nonaqueous Li-O2 Batteries. J. Phys. Chem. Lett. 2013, 4, 2989−2993. (42) Freunberger, S. A.; Chen, Y. H.; Drewett, N. E.; Hardwick, L. J.; Barde, F.; Bruce, P. G. The Lithium-Oxygen Battery with Ether-Based Electrolytes. Angew. Chem., Int. Ed. 2011, 50, 8609−8613. (43) Meini, S.; Tsiouvaras, N.; Schwenke, K. U.; Piana, M.; Beyer, H.; Lange, L.; Gasteiger, H. A. Rechargeability of Li-Air Cathodes PreFilled with Discharge Products Using an Ether-Based Electrolyte Solution: Implications for Cycle-Life of Li-Air Cells. Phys. Chem. Chem. Phys. 2013, 15, 11478−11493. (44) Johnson, L.; Li, C. M.; Liu, Z.; Chen, Y. H.; Freunberger, S. A.; Ashok, P. C.; Praveen, B. B.; Dholakia, K.; Tarascon, J. M.; Bruce, P. G. The Role of LiO2 Solubility in O2 Reduction in Aprotic Solvents and Its Consequences for Li-O2 Batteries. Nat. Chem. 2014, 6, 1091− 1099. (45) Yu, Q.; Ye, S. In Situ Study of Oxygen Reduction in Dimethyl Sulfoxide (DMSO) Solution: A Fundamental Study for Development of the Lithium−Oxygen Battery. J. Phys. Chem. C 2015, 119, 12236− 12250. (46) Bray, W. C.; Liebhafsky, H. A. Reactions Involving Hydrogen Peroxide, Iodine and Iodate Ion. I. Introduction. J. Am. Chem. Soc. 1931, 53, 38−44. (47) Liebhafsky, H. A.; Wu, L. S. Reactions Involving Hydrogen Peroxide, Iodine, and Iodate Ion. V. Introduction to the Oscillatory Decomposition of Hydrogen Peroxide. J. Am. Chem. Soc. 1974, 96, 7180−7187. (48) Kaya, K.; Mikami, N.; Udagawa, Y.; Ito, M. Resonance Raman Effect of I3− Ion by Ultraviolet Laser Excitation. Chem. Phys. Lett. 1972, 16, 151−153. (49) Yeager, E.; Krouse, P.; Rao, K. V. The Kinetics of the Oxygen Peroxide Couple on Carbon. Electrochim. Acta 1964, 9, 1057−1070.

1878

DOI: 10.1021/acsenergylett.7b00462 ACS Energy Lett. 2017, 2, 1869−1878