As it is an irreversible process, not considered for the CO2 capture and release. Overall, these chemical transformations also do not help to mitigate global warming, as the energy consumed in different forms is much higher than the amount of CO2 fixed by chemical reactions. Here we describe the development of re-generable chemical fixation of CO2 by spiroaziridine oxindole, where CO2 is captured  under catalyst-free condition at room temperature both in aqueous and non-aqueous medium even directly from the slow stream of flue gas producing regioselectively spirooxazolidinyl oxindoles, a potential drug. The CO2-adduct is reversed back to the spiroaziridine releasing CO2 under mild conditions. Further both the fixation-defixation of CO2 can be repeated under near ambient conditions for several cycles in a single loop using a recyclable reagent.Chemical fixation of CO Under this circumstances, melting of thousand years old glaciers, desertification of fertile land, rise in ocean water level and acidification of ocean water had caused enormous detriment to diverse ecological environment1. Scientific and technical advancements to curve atmospheric CO2 concentration via limiting industrial emission and use CO2 as an alternative fuel source in the renewable energy sector, had been a recurrent course of study for past few years5. The reduction of CO2 can be considered as a typical cohesive technology to rise artificial efficiency in producing various valuable hydrocarbons like formic acid, methanol, methane, and C2–C4 olefins12. Several fascinating integrated protocols have freshly been reported for hydrothermal and photochemical CO2 reduction, e.g., metal/metal oxide redox reaction based solar two-step water-splitting thermochemical cycle for CO2 reduction via hydrogen generation22. Alongside, chemical fixation of CO2 has gained substantial importance in synthetic chemistry because CO2 could be used as a benign, abundant, inexpensive, and renewable C1 reserve to yield a variety of value-added chemicals e.g. esters, amides, aldehydes, carboxylic acids, alcohols, organic carbonates and 2-oxazolidinones, etc26. In particular, synthesis of therapeutically cherished and synthetically convenient five-membered cyclic urethanes such as oxazolidinones via cycloaddition of CO2 with aziridines has become one of the most promising approaches in this area, because this process possess 100% atom efficiency, which exactly matches one of the most substantial criteria of green chemistry31. Despite being an admirable strategy to chemically capture and recycle CO2, most of these protocols suffer from high energy demand and utilize costly catalysts/ionic liquids to achieve ambient or near ambient condition for CO2 fixation, even from highly enriched CO2 source40. However, emissions from thermal power plants contain numerous gaseous components like SO2, NO2 along with CO241. In these context, post-combustion CO2 capture, release, and storage (CCS) had been the most abundantly used protocols for CO2 purification from industrial exhausts. Various strategies are being industrialized for capture, release and storage (CCS) of CO2 from gas streams, where gas–solid adsorption by metal–organic frameworks, gas–liquid chemi-absorption by amines and carbonation by quick/slacked lime are notable46. However, chemical fixation of CO2 from contaminated sources under mild conditions to produce industrially vibrant chemicals and products faces great defies because of two main reasons: (1) the high ionization potential (IP), and (2) the negative adiabatic electron affinity (EA) of carbon dioxide. Therefore, most of the reports use harsh reaction conditions to overcome the high thermodynamic stability and chemical inertness of carbon dioxide. Hence, the development of a cost-effective and robust protocol for CO2 capture, storage, and release in ambient conditions along with utility is highly desirable. Further, the chemical fixation is an irreversible process producing stable covalent compounds and thus, till now it could not be utilized for CO2 capture and release. It might be a potential CCS protocol as it would produce valuable chemicals, provided the chemical fixation and the defixation (release) done under ambient conditions, the latter is an unmet challenge. Herein, we report the first regenerable chemical fixation, where CO2 fixation by spiroaziridine oxindole under atmospheric pressure at rt (30 °C) without any catalyst producing stable spirooxazolidinone, a potential drug candidate49, further reversed back (defixation) to the spiroaziridine releasing CO2 under mild conditions. This fixation and defixation cycle can be repeated in a single loop for several times using a recyclable reagent.Means of viable development, typically relying on more sensible resource management, is a conceit challenge in front of modern human society. Sustainability level in recent economic growth requires a massive improvement as it is far from an adequate level. According to the data released by Intergovernmental Panel on Climate Change (IPPC 2018), global surface temperature has mounted by approximately 1.5 °C from 1880 to 2018, which is a phenomenon caused by anthropogenic activities, predominantly greenhouse gases like CO2 is an overall non-polar molecule, but the presence of net partial charges [O−δ–C+2δ–Oδ] makes its susceptibility to nucleophilic as well as electrophilic attack at carbon and oxygen, respectively. As a consequence, substrate such as epoxide and aziridine with both reactivity centres are suitable for the fixation of CO240. However, all these require high pressure, -temperature and/or catalyst/additive. Designing substrate with tuned reactivity may lower the pressure and temperature for the chemical fixation of CO2 and may further facilitate the CO2 release. We envision that NH-free spiroaziridine oxindole 1 could be a suitable substrate with desired reactivity as the presence of oxindole unit may enhance the nucleophilicity of aziridine-nitrogen via an electron-donating resonance effect of nitrogen of oxindole unit and/or its anchimeric assistance  in dioxane and aqueous KOH. The exclusive formation of NH-free spiroaziridine 1a was detected by MS and NMR analysis. With great delight, when a slow stream of CO2 was passed through an aqueous dioxane solution of in situ synthesized spiroaziridine 1a at rt, within 30 min it produced exclusively CO2 adduct, spiro oxazolidione 2a in excellent yield . The dioxane was the optimized solvent for both in situ spiroaziridine formation and the chemical fixation of CO2. NaOH instead of KOH is also equally effective for the synthesis of spiroaziridine and subsequent CO2 fixation (entries 3 and 4). Further, when in situ generated spiroaziridine was taken in ethyl acetate and treated with slow stream of CO2 in absence of any base, it also gave the CO2-adduct within 1.5 h in 69% isolated yield (entry 5). Thus it can be concluded that the chemical fixation of CO2 by spiroaziridine does not require base as a catalyst/promoter. Ultimately with our great delight, the auto-chemical fixation of CO2 was successful with 12.5% CO2 in N2 as well as a stimulated coal flue gas  without any appreciable loss in the yield of the adduct (entries 6–8). These took longer reaction time, may be due to low concentration and retention of CO2 in solution.According to the presumption, we started our studies initially on synthesis of NH-free spiroaziridne oxindole 2-chemical fixation. As per the presumption, the decarboxylation (CO2 release) was initiated with the reaction of spiroxazolidinone in the presence of different Brǿnsted acids and the subsequent treatment of base to regenerate the spiroaziridine and its regeneration was quantified with the further chemical fixation of CO2 leading to spirooxazolidione again. Both the CO2 defixation and the fixation were optimized in dioxane and briefly summarized in Table 1a was detected when a dioxane solution of spirooxazolidinone was heated with triflic acid at 100 °C. The extend of formation of spiroaziridine was confirmed by its chemical fixation of CO2 and it gave only 24% yield of the resynthesized spirooxazolidinone 2a . Further, to avoid the cumbersome procedure for the preparation of dioxane-HX, we developed an efficient and handy reagent, NaI-phosphoric acid for the regeneration of spiroaziridine 1a (entry 8)  “”.Table 22a (defixation of CO2), the formation of spiroaziridinium ion 1a′ was detected by MS analysis prior to the treatment with base. However, in case of HI or NaI-H3PO4, exclusive formation of intermediate compound 3-(aminomethyl)-3-iodooxindole 4a′ and no 1a′ was observed by MS analysis prior to the reaction with base. The intermediate iodo-amine 4a′ was isolated and identified as N-tosyl compound 5a by MS and NMR analysis  was repeated for five times through the isolation of regenerated spirooxazolidinone 2a using solid NaOH. All the cycles required equal CO2-defixation and fixation time-scale and exhibited quantitative regeneration of 2a  and the fixation (at rt) were successfully continued for consecutive five cycles in one-pot by treating with NaI-H3PO4 and solid NaOH. Excitingly the overall yield of spirooxazolidinone after five cycles was found to be excellent .The defixation of CO2 of the process, NaI supposed to regenerate after the treatment of 4a′/1a′ with NaOH. So, in principle, NaI may be reused for the subsequent cycles. For the purpose, the first regeneration cycle with release of CO2 was carried out as usual with the combination of NaI-H3PO4 and NaOH and subsequent chemical fixation of CO2 produced the spirooxazolidione. The subsequent cycles for the regeneration of spiroaziridine (CO2-release/defixation) and CO2-fixation were performed without further addition of NaI, only varying with the equivalent of H3PO4 and NaOH .Again, if we deeply look into the chemical reactions involved during the release and capture of CO2 is transported to a different vessel, where it releases the CO2 (regeneration) either after being heated or after a pressure decrease or after any other change of conditions around the sorbent. The sorbent resulting after the regeneration step is sent back to capture more CO2 in a cycle. This makes additional cost of the process. It will be desirable to conduct both CO2 capture and the release in a single vessel, this is possible when both are near similar conditions. In our case, 70 °C was found to be optimum temperature for the CO2 defixation. So, we further studied the temperature effect on CO2 fixation. Interestingly, it showed a near horizontal line for the fixation at 5 °C, 30 °C, 50 °C, 60 °C and 70 °C, respectively, with > 95% yield in each case  loaded with the captured CO2 fixation, we performed both CO2 defixation and fixation at 70 °C and continued for five cycles. With our great delight, it showed almost quantitative yield of spirooxazolidinone 2a in each cycle and an excellent overall yield after five cycles. This chemical fixation-defixation (five) cycles at 70 °C are repeated for three times with a standard deviation of 0.47–1.70 ] can work well in one-pot (single vessel) for several cycles with excellent recovery using recyclable reagent under near ambient conditions. More importantly, the process regioselectively produced bioactive spirooxazolidinoyl oxindole in quantitative yields under extremely mild conditions . The CO2-adducts are stable compounds with high melting points, these can be stored for months under ambient conditions and can be reversed back to the sorbent as and when it requires. So, these findings in the ongoing research can open up a new avenue of the chemical fixation for the development of smart innovative technology towards the energy and cost effective practical CCS and would find abundant applications in CO2 fixation and -defixation chemistry towards chemical utilization of CO2 in industry.In summary, the first regenerable chemical fixation by spiroaziridine oxindole proved to be an excellent protocol for spontaneous and reversible CO3a  was dissolved in dry dioxane (8 ml) and cooled to 0 ºC. Chlorosulfonic acid  was added drop wise and the reaction mixture was stirred for 2 h at room temperature (rt). 14 ml of 1 M aqueous NaOH solution was added dropwise to quench the acid at 0 ºC and stirred at 70 °C for 16 h. The complete conversion to spiroaziridine was detected by MS analysis. Next, a slow stream of CO2 was passed through the solution at rt for 30 min. After complete consumption of 1a  the dioxane was removed under reduced pressure and the residue was extracted with EtOAc (3 × 10 ml), washed with brine solution and dried over anhydrous Na2SO4. Combined organic layer was concentrated and purified by silica gel flash chromatography using EtOAc/hexanes (1:1) to afford the desired CO2-adduct 2a .Amino alcohol Note In case of stimulated flue gas  or 12.5% CO2 in N2, the stream of gas was passed through the solution for 18 h.2a  in dry dioxane (6 ml), sodium iodide  and o-phosphoric acid  were added. The mixture was stirred at 70 °C and the consumption of 2a was monitored by TLC and GC–MS. After 5 h, aqueous NaOH solution  was added and stirred for 30 min. The exclusive regeneration of spiroaziridine 1a was confirmed by MS analysis. No spirooxazolidinone 2a and iodoamine 4a were detected in MS analysis at this stage. The crude solution containing spiroaziridineoxindole 1a was further used for the chemical fixation of CO2. So, the slow stream of CO2 was passed through the solution for 30 min. The GC–MS analysis of the crude mixture with naphthalene as an internal standard showed quantitative formation of spirooxazolidinone 2a (98%). Usual work and flash column chromatographic purification as discussed in general procedure gave the compound 2a .To a solution of spirooxazolidinone 2a  in dry dioxane (6 ml), sodium iodide  and o-phosphoric acid  were successively added at 70 °C and the reaction (consumption of 2a) was monitored by TLC. After complete consumption of 2a (5 h), it was brought to 0 °C and solid NaOH powder  was added to the reaction mixture. After attaining rt, it was stirred for additional 1 h. The stream of 100% CO2 was passed through to the suspended mixture for 1 h at rt. The solid mass was filtered off and washed with dioxane (2 × 5 ml). The combined organic solvent was evaporated to dryness under reduced pressure. The crude compound was dissolved in dioxane (6 ml) and 150 μl of the solution was taken out for the GC–MS analysis with naphthalene (5 mg) as an internal standard. The analysis showed 97% yield of the spirooxazolidinone 2a. So the calculated amount of resynthesized 2a was found to be 148.5 mg and 150 μl of the solution contained 3.7 mg of 2a. The resynthesized compound 2a (148.5 − 3.7 = 144.8 mg) was used for the second cycle for the regeneration of spiroaziridine and the fixation of CO2 using the same procedure as mentioned above i.e. the use of NaI-H3PO4, solid NaOH and the stream of CO2. The GC-yield of the second cycle was observed to be 96%. Similarly, another three cycles were carried and the GC-yields were found to be 99%, 97% and 95%, respectively.To a stirred solution of spirooxazolidinone 2-defixation and fixation cycles were carried out following the similar procedure as above without separating out the solid by-products and isolation of re-synthesized spirooxazolidinone in the intermediate cycles.The one-pot CO2a  in dry dioxane (6 ml), sodium iodide  and o-phosphoric acid  were successively added at 70 °C and the reaction (consumption of 2a) was monitored by TLC. After complete consumption of 2a (5 h), it was brought to rt and solid NaOH powder  was added to the reaction mixture at 0 °C. After attaining to rt, it was stirred for additional 1 h. The stream of CO2 was passed through to the suspended mixture for 1 h at rt. The complete consumption of in situ regenerated spiroaziridine and the formation of spirooxazolidinone 2a were monitored by TLC and MS analysis. Without separating out the solid mass and the isolation of spirooxazolidinone, another consecutive four cycles were repeated by adding the same amount of sodium iodide and o-phosphoric acid followed by solid NaOH and the stream of CO2 for each cycle in the same pot. The consumption of the intermediate substrate and regeneration of the product were monitored during each cycles by TLC and MS analysis. At the end of 5th cycles, the solid mass was filtered off and washed with dioxane (3 × 10 ml). The combined organic solvent was evaporated to dryness under reduced pressure. The crude compound was dissolved in dioxane (6 ml) and the GC–MS analysis with naphthalene as an internal standard showed 95% overall yield of the spirooxazolidinone 2a for the five cycles. The silica gel flash column chromatographic purification of the crude with hexanes-EtOAc (1:1) gave the spirooxazolidinone 2a  as a white solid.To a stirred solution of spirooxazolidinone 2a  in dry dioxane (6 ml), sodium iodide  and o-phosphoric acid  were successively added at rt and the reaction (consumption of 2a) was monitored by TLC. After complete consumption of 2a (5 h), solid NaOH powder  was added to the reaction mixture at 0 °C. After attaining to rt, it was stirred for additional 1 h. The stream of CO2 was passed through to the suspended mixture for 1 h at rt. The complete consumption of in situ regenerated spiroaziridine and the formation of spirooxazolidinone 2a (97% GC yield) were monitored by TLC and MS analysis. For the next cycle, the reaction mixture was acidified with o-phosphoric acid  and stirred at 70 °C without further addition of sodium iodide. After complete consumption of the spirooxazolidinone (monitored with TLC), solid NaOH powder was added  and the stream of CO2 was passed through for 1 h to reproduce the spirooxazolidinone 2a . This process was repeated for five consecutive cycles. GC–MS analysis showed almost quantitative yield of spirooxazolidinone in each stage and finally the spirooxazolidinone 2a  was isolated after fifth cycle by flash chromatography using hexanes-EtOAc (1:1).To a stirred solution of spirooxazolidinone Note (a) GC yield is determined by using naphthalene as internal standard; (b) the release of CO2 from spiroxazolidinone and its subsequent regeneration using CO2 fixation is considered as one complete cycle. (c) At constant temperature (70 °C) five consecutive cycles of CO2 fixation and defixation was accomplished using above method . GC yield in resynthesis of spirooxazolidione 2a was monitored at each stage (Fig. age Fig. .Supplementary Information.


We employed mostly X‐ray photoelectron spectroscopy (XPS) in combination with dispersion‐corrected density functional theory calculations (DFT‐D3). If the surface is pre‐covered by BMP‐TFSI species (model electrolyte), post‐deposition of Li (Li+ ion shuttle) reveals thermodynamically favorable TFSI decomposition products such as LiCN, Li2NSO2CF3, LiF, Li2S, Li2O2, Li2O, but also kinetic products like Li2NCH3C4H9 or LiNCH3C4H9 of BMP. Simultaneously, Li adsorption and/or lithiation of Co3O4(111) to LinCo3O4 takes place due to insertion via step edges or defects; a partial transformation to CoO cannot be excluded. Formation of Co0 could not be observed in the experiment indicating that surface reaction products and inserted/adsorbed Li at the step edges may inhibit or slow down further Li diffusion into the bulk. This study provides detailed insights of the SEI formation at the EEI, which might be crucial for the improvement of future batteries.In this work we aim towards the molecular understanding of the solid electrolyte interphase (SEI) formation at the electrode electrolyte interface (EEI). Herein, we investigated the interaction between the battery‐relevant ionic liquid (IL) 1‐butyl‐1‐methylpyrrolidinium bis(trifluoromethylsulfonyl)imide (BMP‐TFSI), Li and a Co The formation of the solid electrolyte interphase at the electrode‐electrolyte interface is studied by X‐ray photoelectron spectroscopy and dispersion‐corrected density functional theory calculations. The interaction between an adsorbed ionic liquid, Li, and a Co3O4(111) thin‐film model anode reveals bis(trifluoromethylsulfonyl)imide and 1‐butyl‐1‐methylpyrrolidinium decomposition and lithiation of the cobalt oxide. Transition‐metal oxides such as cobalt oxides are of significant interest as battery conversion anodes in LIBs). The STM image in Figure a→|=|b→|=0.57±0.03 nm and an angle of α=120±5°, in excellent agreement with results of previous reports.).In the O 1s range the OBMP and NTFSI peaks simultaneously decrease in intensity up to 0.35 MLE of Li (−46 %). Now, a new peak (Nnew) evolves at ∼398 eV at the equal cost of both NBMP and NTFSI peaks. Though a rigorous assignment of this new peak is difficult, we tentatively propose that LiCN might be formed. Related compounds such as the lithium tetracyanoquinodimethane charge‐transfer complex (Li−TCQN) shows a BE of 398.6 eVBMP peak stepwise decreases in intensity and is almost completely absent after the last Li dose. Interestingly, at this point, the NTFSI peak does not decrease any further, even if for all other TFSI‐related peaks, we monitored a continuous decrease upon increasing the amount of Li (see description of FTFSI and OTFSI above and CTFSI and STFSI below in the text). Hence, we assume that at least part of the NTFSI peak (cyan) is related to a decomposition product either of TFSI or of BMP. A possible first TFSI decomposition product might be LixNSO2CF3, which could have been formed by dissociation of the S−N bond and subsequent bonding with Li. Such a species would be expected to exhibit a rather similar BE with the former NTFSI peak and hence cannot be resolved in our experiment.In the N 1s range the NxNCyHz, which requires C−N bond breaking. Studying reduction reactions of BMP by DFT calculations, Haskins et al.xNHy) show binding energies in the N 1s range of about 401.7 eV . The total peak area decreases again moderately by 12 % after deposition of 0.15 MLE of Li and by 37 % after the last deposition step.A possible BMP decomposition product would be Lis, R3N2NH76, 77TFSI peak (−CF3 bond) gradually decreases in intensity upon increasing the Li dose (0.15–4.0 MLE of Li). After post‐deposition of 0.15 MLE of Li the CTFSI peak decreased by 42 % and almost completely after the last Li dose. Going to the CBMP and Calkyl peaks, a quantitative analysis during Li exposure is hardly possible because of the broad feature between 285 and 290 eV induced by these species, which can principally be fitted by two peaks of varying intensity. Hence, we correlated the intensity changes of the CBMP peak with that of the NBMP peak, while the Calkyl peak was free floating. After the last Li dose only a small amount of Calkyl remains. The total peak area decreases by 11 % after deposition of 0.15 MLE of Li and by 69 % after the last deposition step.In the C 1s range it is clearly visible that also the anion‐related CTFSI doublet gradually decreases in intensity upon increasing the Li dose, while two new peaks (Snew1 and Snew2) emerge. After the first Li dose (0.15 MLE), the Snew1 doublet arises with a BE of the S 2p3/2 peak at ∼168 eV, the second new doublet (Snew2) sets in at 161.7 eV after deposition of 0.7 MLE of Li. After post‐deposition of 0.15 MLE of Li the STFSI peak decreased by roughly 20 % and by ∼70 % after the last Li deposition step . The total peak area did almost not change after deposition of 0.15 MLE. The total intensity loss after the last Li exposure is ∼20 %.Moving to the S 2p range, the SBMP peak also reveals a pronounced intensity loss. The loss of intensity of the anion‐ and cation‐related peaks in the XP spectra is accompanied by the formation of new peaks (filled yellow), reflecting the formation of various decomposition products. Based on their BEs, possible products are LiCN , LiF , Li2O2  and Li2S  as well as other Li bonded fragments of the anion such as LixNSO2CF3 or the cation such as LixNCyHz. The formation of these Li‐bonded fragments is associated with the initial stages of the chemical SEI formation. In general, in our previous experiments we observed similar new peaks (decomposition products) on different model electrode surfaces such as HOPG,[Overall, during stepwise postdeposition of approximately 0.15–4.0 MLE of Li all TFSI‐related peaks gradually decrease in intensity, reflecting the decomposition of the anions. The cation‐related N as HOPG,13, 14To shed some light on possible first reaction products of BMP‐TFSI and Li, we employed DFT‐D3 calculations of the BMP‐TFSI crystal structure in which we inserted up to four Li atoms. By looking first at the electronic properties of the neat BMP‐TFSI crystal . In particular, among the various decomposition products that are studied, Li2NSO2CF3 . Thus, thermodynamically, the formation of Li clusters is at least as favorable as the decomposition of BMP. However, due to a low barrier for the reaction of Li with BMP‐TFSI, which was already reported for the decomposition of BMP‐TFSI monolayers at graphite,2NCH3C4H9 due to the reaction of 2 Li atoms with BMP is probable. Along with the formation of Li2NCH3C4H9 two ethylene molecules desorbing into the gas phase are generated and Li2NCH3C4H9 retains the positive charge of BMP. The corresponding neutral moiety LiNCH3C4H9, which may form by releasing Li+, resembles in its structure and in the environment of N a secondary amine and thus might also have a similar N 1s core level binding energy. In order to see whether the computationally determined stable reaction products are compatible with experimental XPS results, a first estimate of the qualitative order of the N 1s core level BEs of the most stable decomposition products might be given by the atomic charges. Therefore, here we compare the number of electrons within Rwigs=0.741 Å around N of BMP, TFSI and different decomposition products within the BMP‐TFSI crystal:where nTFSI peak, which might be due to the decomposition products Li2NCH3C4H9 and/or Li2NSO2CF3. Note that a rigorous correlation of the atomic charges to the core level BE can be difficult, as final state effects might be crucial. Yet, a more detailed analysis of further decomposition products of BMP‐TFSI due to the reaction with Li and the corresponding core level shifts is beyond the scope of this article and will be discussed elsewhere.The trend tentatively suggests that among the compounds studied, BMP shows the highest BE of the N 1s electron, while LiCN exhibits the lowest one. According to that, we are inclined to assign the new peak at a BE of 398.5 eV to LiCN that forms due to a decomposition reaction of TFSI with Li. Besides, we propose that there is at least one other new peak at a BE presumably only slightly higher than the BE of the N3O4(111) pre‐covered by a BMP‐TFSI adlayer (1.5–2 ML), we also monitored the Co 2p regime to test for changes of the chemical state of the thin film model electrode   and after deposition of 0.35, 0.7, 1.1 and 3.1 MLE of Li. It is clearly visible that already after exposure of 0.35 MLE of Li (violet solid line) the two Co 2p maxima shift from around 779.7 and 794.1 eV to 780.2 and 796.0 eV, respectively. Simultaneously the two SCo2+ satellites significantly increase in intensity. For higher Li doses no significant changes occur. In addition, we show the corresponding XP difference of the XP spectra during Li deposition minus the spectrum of pristine Co3O4  at the BEs of Co3+. After post‐deposition of 0.7, 1.0 and 3.1 MLE of Li, the Co3+ loss is only slightly enhanced, and only a marginal decrease of the two SCo2+ satellites becomes visible, which is maybe due to damping effects induced by adsorbed Li species. The BE shift from Co3+ to Co2+, that is, an increasing amount of Co2+ relative to Co3+ can be ascribed to charge transfer from Li0 to Co3+ due to insertion or/and adsorption. Indeed, Yao et al.3O4. Their proposed mechanism includes Li insertion into available host sites, leading to LinCo3O4 up to n=3 and Co3+ to Co2+ reduction. For n=3 they found a mosaic structure with local Li2O and CoO character, for larger Li content the local transformation to Li2O and CoO and finally to Co0.In the XPS experiment, during the stepwise postdeposition of 0.15–4.0 MLE of Li on Coe Figure . Figure 3O4(111) before  to ∼1.2, that is, the ratio used for deconvolution of both pristine CoO and Co3O4. After Li deposition, fitting reveals a significant change of the ratio of Co3+ : Co2+ to 1 : 1.4. The full lithiation of the cobalt oxide from Co2+Co3+2O2−4 to Li+Co2+2Co3+O2−4 would result in an inversion of the relative amount of Co2+ to Co3+ species, hence from the experimental results we assume at least a partial lithiation of the cobalt oxide. However, based on the Co 2p spectra we cannot resolve wether this modification is the result of an insertion reaction of Li into Co3O4 to form LinCo3O4, of Li adsorption, a partial transformation of Co3O4 to CoO or a combination thereof.Peak deconvolution of pristine Coe Figure c and afte Figure d reveals3O4(111) allow us to gain further insights into the processes at the interface  in the analysis, the stable adsorption sites can be anticipated by looking at the electrostatic potential of the bare surface: the electrostatic potential (of the electron) mapped onto an isosurface of the charge density is shown in Figure 2+ atoms . The adsorption energy per Li atom is given relative to the energy of the bare surface and bulk Li metal: it amounts −3.47 eV  at a coverage of 0.9 Li atoms/nm2 and decreases to −3.01 eV  if all adsorption sites A are occupied (i. e. at a coverage of 3.5 Li atoms/nm2). This clearly reveals the anticipated strongly ionic character of the bond of Li to the substrate. Finally, the ionic character is shown by the calculation of the change of the atomic charges  due to Li adsorption at a coverage of 3.5 Li atoms/nm2. It shows that charge is transferred from Li atoms (Δq/atom=+0.9 e) to the topmost Co2+ atoms (Δq/atom=−0.11 e) and O atoms (Δq/atom=−0.16 e), but also to directly underneath located octahedrally coordinated Co atoms (Co3+) (Δq/atom=−0.04 e). The atomic charge of the latter atoms reduces from 1.33 e to 1.29 e, which corresponds to the value of Co2+ atoms in the bulk (see section 3.1).DFT‐D3 calculations of the adsorption (or insertion) of Li atoms at (or in) Coe Figure . Assumin2. The diffusion path passes a shallow local minimum or saddle point . It is associated with a barrier of 1.0 eV which is higher than the value that was reported for Li diffusion along octahedral sites in the Co3O4 bulk by PBE calculations (0.75 eV).3O4(111).3O4(111) surface slab have also been studied. The most stable surface structure in which Li is underneath the topmost O atoms is about 2.9 eV less stable than the adsorption of Li  in Figure Considering the fact that the diffusion barrier at the surface is rather high and diffusion into the bulk through the defect free (111) surface is difficult, we conclude that at least a small part of the Co0 (∼778 eV). Here, we tentatively assume that surface reaction products (SEI) and maybe inserted (adsorbed) Li in Co3O4 in the surface region (on the surface) and in particular at the step edges inhibit further Li diffusion into the bulk, or the diffusion is at least very slow.Different from the Li deposition on CoO(111) (pristine and IL pre‐covered),4Ti5O12(111)4+ partially transforms into Ti3+ species for TiO(110) and Li4Ti5O12(111), while Co2+ transforms into Co0. Also in the latter studies, the IL adayer/decomposition products are at least in parts permeable for Li, allowing for the reaction with TiO2(110), Li4Ti5O12(111) and CoO(111) surfaces. Differently from CoO(111), where Co2+ is transformed to Co0, the reduction of Ti4+ proceeds only to the Ti3+ stage. In the latter cases, reaction of Li with the substrate was found to take place both in the absence and the presence of an adsorbed IL adlayer,[We finally compare with previous results on oxidic model electrodes measured in our laboratory, i. e., the interaction of a pre‐adsorbed BMP‐TFSI monolayers with post‐deposited Li on TiO(110), adlayer,26, 323O4(111) thin film on Ir(100) covered with a pre‐adsorbed BMP‐TFSI adlayer (1.5–2 ML) at r.t. results in the gradual decomposition of BMP‐TFSI, which can be considered as the initial stage of the chemical SEI formation at the EEI in the absence of an electric field, i. e., at the open circuit potential (OCP). Both the anions and cations gradually decompose upon stepwise increasing the Li exposure. The efficient SEI formation at r.t. reflects a rather low barrier for IL decomposition upon interaction with Li. Simultaneously with IL decomposition (SEI formation) the chemical state of Co3O4(111) is modified in the near surface region, which we either assign to the insertion of Li into Co3O4 leading to LinCo3O4 structures , which partially show localized CoO−Li2O character,0. Thus, the conversion reaction is most likely inhibited (or diffusion is very slow) by IL/decomposition products (SEI) and/or inserted/adsorbed Li particularly at the steps. These experiments thus reproduce not only the initial stages of SEI formation (surface chemistry) but also the first lithiation step of Co3O4 under idealized conditions, much better defined than the complex situation in a battery cell.Overall, the measurements in this work reveal that stepwise postdeposition of Li on a Co3+ shuttle) on a Co3O4(111) thin film model anode and also the changes of the chemical state of Co3O4(111). Our main results and conclusions are:In order to get deeper insights into the processes at the model electrode electrolyte interface in LIBs, we have investigated the interaction of ultrathin films of the ionic liquid BMP‐TFSI (model solvent/electrolyte) with Li (effect of the Li3O4.RPBE‐D3+U is a suitable method to describe both the interactions within BMP‐TFSI and Co3O4(111) thin films with extended flat island structures.XPS and STM experiments prove the formation of Co3O4(111) at r.t., except of a small amount of decomposition products which we tentatively assign to the interaction with the steps.Based on the XP spectra BMP‐TFSI adsorbs as molecular species on Co3O4(111) surface pre‐covered with a BMP‐TFSI adlayer instantaneously results in TFSI and BMP decomposition for the initial Li doses, reflecting the rather low reaction barrier for IL decomposition upon interaction with metallic Li. This process is considered as the initial stage of the chemical SEI formation at the EEI. XPS measurements reveal possible IL decomposition products of the TFSI anion such as LiCN, Li2NSO2CF3, LiF, Li2S, Li2O2, Li2O, etc. The decrease of the total peak area of adsorbate‐related peaks during Li deposition indicates that beside decomposition also desorption takes place.Stepwise postdeposition of small amounts of Li on a Co2NSO2CF3 and in particular LiCN are two of the most stable initial decomposition products. Decomposition products of BMP are less stable; thermodynamically, the formation of Li clusters is at least as favorable. Note, however, that due to a low barrier for the reaction of Li with BMP‐TFSI kinetic products such as Li2NCH3C4H9+ or LiNCH3C4H9 are possible, which can form by ethylene molecules desorbing into the gas phase in agreement with the observation of desorption in the XPS measurements.The DFT‐D3+U calculations reveal that the decomposition of TFSI is more favorable than the reaction of BMP with Li. Li2+ states relative to Co3+, which is attributed to charge transfer from Li0 to Co3+. This could be due to a partial lithiation of the cobalt oxide to LinCo3O4 and/or Li adsorption. Furthermore, we cannot exclude that also a partial transformation of Co3O4 to CoO takes place. Larger Li deposits do not induce further significant changes in the Co 2p range, most likely because surface reaction products and maybe inserted/adsorbed Li in Co3O4 in the surface region, in particular at the step edges may either inhibit further Li diffusion into the bulk or the diffusion is at least very slow. The formation of Co0 due to a conversion reaction of CoO to Co0 and Li2O could not be observed in the present experiments.After initial post‐deposition of Li to a pre‐adsorbed BMP‐TFSI adlayer, the XP Co 2p core level spectra show the formation of an increasing amount of Co3O4(111) surface is rather high and diffusion into the bulk through the defect free (111) surface is due to a prohibitively large barrier very unlikely. Hence, we assume that at least a small part of the Co3+ reduction seen in XPS experiments must be due to Li adsorption. Further contributions most likely stem from insertion reactions of Li into Co3O4 via step edges or defects.DFT‐D3+U calculations reveal that the diffusion barrier at the Co3O4(111) surface covered with an ultrathin IL layer. It provides valuable information about the basic processes occurring at the EEI which might be crucial for the further development of future batteries.In summary, the present surface chemistry model study reveals molecular details of the interactions of Li at the CoThe authors declare no conflict of interest.As a service to our authors and readers, this journal provides supporting information supplied by the authors. Such materials are peer reviewed and may be re‐organized for online delivery, but are not copy‐edited or typeset. Technical support issues arising from supporting information (other than missing files) should be addressed to the authors.SupplementaryClick here for additional data file.


XTCNQ  are investigated experimentally and theoretically. By tuning the number of F atoms in the acceptor molecules, the resulting complexes display a continuous down shift of the valence band maximum, conducting band minimum, and optical bandgap. The majority carriers convert from hole (DPTTA‐F1TCNQ), balanced hole, and electron (DPTTA‐F2TCNQ) to electron (DPTTA‐F4TCNQ). This result shows that band engineering can be realized easily in the donor–acceptor complex systems by tuning the electron affinity of the acceptor. The bandgaps of these three complexes vary from 0.31 to 0.41 eV; this narrow bandgap feature is crucial for achieving high thermoelectric performance and the unintentional doping in DPTTA‐F4TCNQ leads to the effective suppression of the bipolar cancelling effect on the Seebeck coefficient and the highest power factor.Three isostructural donor–acceptor complexes DPTTA‐F XTCNQ complexes leads to continuous tuning of the energy gaps and the species of majority carriers. Efficient band structural engineering effects induced by tuning the electron affinity of the acceptor component are thus shown.Variation of the number of F atoms in the acceptor units of a series of isostructural‐crystal DPTTA‐F KbT (Kb refers to Boltzmann constant and T is the absolute temperature). Such character is critically important for inorganic semiconductor to achieve superior thermoelectric (TE) performance, but still unexplored for organic counterpart.σS2, σ and S are electrical conductivity and Seebeck coefficient, respectively), which is due to the formation of gapless polaron band under the scenario of highly doped state.Efficient band engineering  of organic semiconductors is still a great challenge in organic electronics, although it has already been a general strategy to vary the electronic and optical performance of inorganic semiconductors to meeting the requirement on specialized device applications.XTCNQ , balanced hole and electron (DPTTA‐F2TCNQ) to electron (DPTTA‐F4TCNQ). Band structures, as well as the intermolecular electronic coupling were calculated. The observed tuning of the majority carrier species and electronic structures show that the band engineering of the DA complexes can be realized by tuning the electron affinity of the acceptor component.Recently, we reported two mixed stack DA complexes DPTTA‐Fnnulene2,,2,1, FXTXTCNQ suitable for single crystal structure analysis grew through a diffusion method according to the procedure reported previously.P1¯ space group. It should be noted that the F atoms in DPTTA‐F1TCNQ disordered over 2, 5‐positions and in DPTTA‐F2TCNQ over 2, 3, 5, and 6‐positions of the quinoid ring, respectively. So, F1TCNQ presents as a pseudo‐F2TCNQ molecule and F2TCNQ as a pseudo‐F4TCNQ molecule in the complex crystals, which leads to the almost identical structure of these crystals not only in the cell parameters but also the relative position between the donor and the acceptor molecules. FigureXTCNQs are sensitive to the charge transfer (CT) process (see the note of Table S2 in the Supporting Information for the bond labeling). The bond length ratio r value (r = c/(b + d)) can estimate the degree of charge transferr values are in between that of the neutral and fully ionized acceptor molecules, and the CT degrees derived to be 0.111, 0.213, 0.923 for DPTTA‐F1TCNQ, DPTTA‐F2TCNQ, and DPTTA‐F4TCNQ, respectively. The CT ratio increases with the number of F atoms in acceptor molecules, which can be further confirmed by the IR spectra . It should be noted that this result is different from what we previewed in DPTTA‐TCNQ complex, where the molecules maintain in the neutral state.1TCNQ, DPTTA‐F2TCNQ, and DPTTA‐F4TCNQ respectively, indicating the existence of intermolecular π–π interaction in the DA pairs.Evaluation of the single crystal structure is crucial for understanding the structure–property relationships of organic semiconductors. DPTTA, a sulfur‐bridged annulene, is a typical electron donor and was demonstrated to be one of the most suitable donor components for construction of DA complexes.1TCNQ together with the F2TCNQ and F4TCNQ based crystals were characterized with the fabrication of SCFETs for deriving the carrier mobilities. SCFETs with bottom‐gate top‐contact (BGTC) configuration were fabricated based on corresponding microcrystals growing on SiO2/Si substrates. The detailed description of the fabrication process is available in the reference.Figure2 V−1 s−1 (μh), 0.24 cm2 V−1 s−1 (μe) for DPTTA‐F1TCNQ; 1.01 cm2 V−1 s−1 (μh), 0.27 cm2 V−1 s−1 (μe) for DPTTA‐F2TCNQ; and 0.11 cm2 V−1 s−1 (μh), 0.46 cm2 V−1 s−1 (μe) for DPTTA‐F4TCNQ. Compared to the results reported previously, the value of μh and μe of DPTTA‐F2TCNQ and DPTTA‐F4TCNQ is slightly lower, but still in a reasonable range considering the quality variation of the solution processed single crystals. The P channel and N channel threshold voltage (Vth) is −1.2 V, 18.4 V for DPTTA‐F1TCNQ; 43.7 V, −117.4 V for DPTTA‐F2TCNQ; 44.1 V, −43.1 V for DPTTA‐F4TCNQ respectively. The corresponding output curves of transistors based on DPTTA‐F1TCNQ, DPTTA‐F2TCNQ, and DPTTA‐F4TCNQ single crystals are presented in Figure S2 in the Supporting Information. Figure S3 in the Supporting Information shows the on/off current ratios (Ion/Ioff). The mobility distribution for these DPTTA based devices can be seen in Figure S4a,c,e in the Supporting Information. The structures of the microcrystals used for preparing devices were examined using X‐ray diffraction (XRD)  planes. For DPTTA‐F2TCNQ, the strong peaks at 8.24°, 16.41°, and 24.68° are indexed as (001), (002), and (003) respectively. In the case of DPTTA‐F4TCNQ, the strong peaks at 8.11°, 16.15°, and 16.39° correspond to (001), (002), and (003). The data show that the crystals all grow along a preferring direction parallel to the DA alternative stacking direction. The corresponding selected area diffraction patterns (SAED) of the microcrystals are shown in Figure S4b,d,f in the Supporting Information.Charge carrier mobility, electrical conductivity, and Seebeck coefficient are the most important parameters reflecting the semiconducting behaviors of the CT complexes. Here, the newly prepared DA complex DPTTA‐Fn is the charge carrier concentration, and μ is the carrier mobility.−1 for DPTTA‐F1TCNQ, DPTTA‐F2TCNQ, and DPTTA‐F4TCNQ respectively. Before deriving the charge carrier concentration via applying Equation 1TCNQ complex displays the lowest off‐current in the range of 2 × 10−10–4 × 10−10 A (under a VD of 10 V), corresponding to the largest energy gap of this material as the optical spectra revealed, which will be discussed below. Meanwhile the symmetric transfer curve of DPTTA‐F1TCNQ along VG = 0, shows that ambipolar transport behavior is around zero gate voltage and the material possess relatively balanced hole and electron concentration.2TCNQ complex devices show similar situation, but a much higher off‐current in the range of 2 × 10−7–4 × 10−7 A with VD = 10 V. The symmetric transfer curve and ambipolar transport character around VG = 0 V show a balance bipolar state of this material. What observed with the transfer curves and output curves of F4TCNQ devices are quite different from that of previous two. The symmetric ambipolar transfer curve is only observed around the VG of −30.6 V, showing that extra negative gate voltage is needed for inducing holes to compensate the dominating electrons. The electron domination behavior can be verified with the output characters in the p‐type regime, which shows depleting feature,VG applied from 0 to −60 V and get increasing again when VG exceeding −80 V . This could be interpreted with the fact that when complexing with F4TCNQ, the strongest electron acceptor, the resulting material possess the deepest conducting band minimum (CBM) according to the band structure calculation as discussed below, which leads to the unintentional n‐doping by moisture from the ambient circumstance. So, when we apply Equation ne/nh and μe/μh refer to the electron/hole concentration and mobilities, respectively.1TCNQ and DPTTA‐F2TCNQ the balanced hole and electron density is calculated to be 1.31 × 1016 and 3.4 × 1017 cm−3, while a dominating electron density of 8.15 × 1017 cm−3 is estimated for DPTTA‐F4TCNQ (TableThe conductivity can be defined asS directly reflects the polarity of the majority charge carrier (hole or electron) of unipolar materials. For bipolar materials, where electron and hole contribute to the electrical conductivity simultaneously, the overall Seebeck coefficient can be described by the conductivity weighted average according to the Equation Sh/Se and σh/σe refer to the Seebeck and electrical conductivity of the two different carrier species.ne = nh and Sh = −Se.S is determined by the mobility ratio between electron and hole. It seems that balanced bipolar character supposed for DPTTA‐F1TCNQ is conflicting with the positive S measured (830 µV K−1), as the mobility of electron is slightly larger than that of hole according to the SCFET devices characterizations. This indicates that carrier concentration of hole and electron is deviating from equilibrium, with the hole concentration exceeding electron concentration in this crystal. As the mobilities of the two carrier species are close to each other, this nonequilibrium could not be observed through the FETs characterization. The measured seebeck coefficient is 185 µV K−1 for DPTTA‐F2TCNQ. The absolute value of Se and Sh can be calculated to be 320 µV K−1 according to Equation S value of hole and electron in this ambipolar semiconductor, showing the detrimental bipolar effect on the TE performance. The negative Seebeck coefficient observed with DPTTA‐F4TCNQ is −443 µV K−1. The absolute value is obviously larger than that of DPTTA‐F2TCNQ showing the bipolar cancelling effect in this material has been minimized due to the dominant electron concentration induced by unintentional doping.Seebeck coefficient measurement is a powerful tool for characterizing the semiconducting materials.Figure4TCNQ. PF of DPTTA‐F4TCNQ is 0.33 µW m−1 K−2 at 300 K, which gradually increases to 0.62 µW m−1 K−2 at 400 K. The unintentional doping results in the electron dominant transport in F4TCNQ complex and suppress the bipolar effect on the Seebeck coefficient, which leads to highest TE performance in these DA complexes.To characterize the thermoelectric properties, the temperature dependence of the electrical conductivity and the Seebeck coefficient of those complexes were measured Figurea,b. The T can be observed and fitted with the Arrhenius equation in the temperature range of 300–400 K  and DPTTA‐F4TCNQ (0.098 eV)  is a critical parameter determining the TE performance. Here the bandgap was determined using infrared spectroscopy (IR) by applying the Tauc method: (αhν)n ∝ (hν − Eg). In this equation, α is the absorption coefficient, hν is the photon energy, and n = 2 and 1/2 for direct and indirect gaps. Eg is derived from a linear extrapolation to zero absorption in the plot.Figure1TCNQ, DPTTA‐F2TCNQ, and DPTTA‐F4TCNQ are 0.41, 0.33, 0.31 eV respectively  go deeper according the calculation  to F4TCNQ (35 meV) complexes were also reproduced in the quantum simulation  display narrow optical bandgap in the range of 0.31 to 0.41 eV, and continuous down shift of the VBM and CBM. This leads to a thermal activated bipolar character and majority carrier switching from hole dominated semiconductor in F1TCNQ complex to electron dominated semiconductor in F4TCNQ complex. This result shows that tuning of the molecular frontier energy level and the intermolecular interactions including the charge transfer degree, intermolecular electronic coupling can be easily realized in the DA complex systems, which leads to the efficient tuning of the band structure of organic semiconductors. Unintentional doping leads to the break of balanced bipolar transport which is important for achieving a high TE performance in these narrow bandgap organic semiconductors, similar to that of the inorganic counterpart.Isostructural DA complexes formed by DPTTA and FMaterial and Crystal: DPTTA was purchased from Solarmer Beijing. F1TCNQ, F2TCNQ, and F4TCNQ were purchased from TCI. They were all used without further purification.2/Si substrates.Single crystals based on DPTTA suitable for X‐ray single crystal structure characterization and measurement of thermoelectric properties were obtained through diffusion method. A solution of DPTTA was mixed in chlorobenzene and the corresponding acceptor (molar ratio 1:1) in the same solvent. After 24 h, the black shiny crystals were collected which can be used for IR, XPS, and UPS measurement. Redissolving the obtained crystals in chlorobenzene to get their saturated solutions. Then microcrystals suitable for FET device fabrication could be obtained by drop casting of the saturated solution onto the SiOStructure Characterization and Measurements: IR spectra were investigated using a TENSOR‐27 spectrometer (Bruker). XPS and UPS were performed by using an AXIS Ultra‐DLD ultrahigh vacuum photoemission spectroscopy system (Kratos Co.). TEM and SAED measurements were carried out on a JEM 2011 (JEOL). XRD was measured on Empyrean  with CuKα source (λ = 1.541 Å).X‐ray crystallographic data were collected on ST Saturn724+ (Rigaku), using CuKα radiation (λ = 1.541 Å).I–V characteristics of the SCFETs were recorded with a Keithley 4200 SCS and a Micromanipulator 6150 probe station at room temperature in air. The mobilities were calculated using Equation ID is the drain current, Ci is the capacitance per unit area of insulating layer which was 11 nF cm−1 in the experiment, L is the channel length, W is the channel width and Vth is the threshold voltage.2TCNQ and DPTTA‐F4TCNQ. The measurement of DPTTA‐F1TCNQ was unreliable in this system due to voltage overload, so it was measured with SB‐C‐01 Seebeck Measurement System . Sample crystals (about 0.3 × 3 mm) were fixed on the measuring boards with conductive carbon paint (SPI).The measurement of thermoelectric parameters  of the crystals was all carried out in vacuum. Electrical conductivity was measured by a standard four‐probe method using a Physical Property Measurement Systems instrument . Striped patterns  of gold electrodes (with thicknesses of 60 nm) were thermally evaporated onto the upper surface of the selected crystals (about 0.3 × 3 mm) then they were mounted on measuring boards via the connection of gold wires with a diameter of 25 µm and fixed with gold conductive paint (EMS). Seebeck coefficient was measured using the SB‐100 Seebeck Measurement System (MMR Tech.) for DPTTA‐FQuantum Simulation: Electronic structure calculations on the isolated DA complexes, ADA and DAD triads have been performed through B3LYP functional and 6–31G(d) basis sets implemented in Gaussian09. Band structure calculation was carried out by using the plane‐wave technique as implemented in CASTEP code. An 800 eV cutoff energy was used for the plane‐wave basis set, and norm conserving pseudopotentials were employed for all ions. The exchange correlation energy is described by the generalized gradient approximation (GGA) in the form of Perdew, Burke, and Ernzerhof (PBE).The authors declare no conflict of interest.Supporting InformationClick here for additional data file.


In this study, experimental data and theoretical calculations demonstrate that the recharge overpotential (ηRC) of Li–O2 battery is fundamentally dependent on the Li2O2 crystallization pathway which is intrinsically related to the microscopic structural properties of the growing crystals during discharge. The Li2O2 grown by concurrent surface reduction and chemical disproportionation seems to form two discrete phases that have been deconvoluted and the amount of Li2O2 deposited by these two routes is quantitatively estimated. Systematic analyses have demonstrated that, regardless of the bulk morphology, solution‐grown Li2O2 shows higher ηRC (>1 V) which can be attributed to higher structural order in the crystal compared to the surface‐grown Li2O2. Presumably due to a cohesive interaction between the electrode surface and growing crystals, the surface‐grown Li2O2 seems to possess microscopic structural disorder that facilitates a delithiation induced partial solution‐phase oxidation at lower ηRC (<0.5 V). This difference in ηRC for differently grown Li2O2 provides crucial insights into necessary control over Li2O2 crystallization pathways to improve the energy efficiency of a Li–O2 battery.Identification of a low‐potential decomposition pathway for lithium peroxide (Li Herein, a fundamental correlation between the lithium peroxide crystallization pathway and subsequent recharge mechanism is established by combination of theory and experiment. The lithium peroxide grown by concurrent disproportionation and surface reduction is found to form two coexisting phases. Quantitative deconvolution shows that the lithium peroxide grown on electrode surface decomposes at low overpotential by a partial solution phase oxidation. The overall reversible electrochemical reaction in a nonaqueous Li–O2 battery involves a two‐electron (2e−) redox of oxygen (O2) in the presence of lithium ions (Li+) (2Li+ + O2 + 2e− ↔ Li2O2(s); E0 = 2.96 V vs Li/Li+) producing solid lithium peroxide (Li2O2) as the discharge ) product, which subsequently decomposes in the reverse reaction during recharge ). Despite the promise of Li–O2 battery as a next generation storage system, low round‐trip energy efficiency (ratio of energy densities during DC and RC respectively) due to high RC overpotential  and measured (E) potential (η = |E − E0|); usually ηRC > 1 V) and as a whole poorly understood RC mechanism pose a significant challenge before its realization. Although several electrocatalysts and soluble redox‐mediators were used to reduce the ηRC, they could not truly improve the reversibility of oxygen‐electrochemistry, rather the solid catalysts exacerbated parasitic reactions and shuttling of redox‐mediators on the unprotected negative electrode led to passivation of Li metal. Moreover, use of an electrocatalyst seems redundant due to low kinetic overpotential for OER in Li–O2 battery as reported by density functional theory (DFT) calculations. Consistent with these theoretical studies, several experimental reports also found that the onset potential for Li2O2 oxidation is indeed very close to the equilibrium potential. However, despite the low kinetic overpotential for OER, the ηRC gradually increases during the course of RC and a two‐stage RC profile consisting of a low‐potential (<4 V vs Li/Li+) sloping region and a high‐potential (>4 V vs Li/Li+) flat plateau is usually observed for catalyst/redox‐mediator free Li–O2 batteries. The rise in ηRC has been qualitatively attributed to several interrelated factors such as physicochemical properties of the discharged Li2O2, stability of the electrode/electrolyte, presence of additives/impurities in the electrolyte etc. Unfortunately, the effects from these factors have not been decoupled yet and combining results from independent experiments is often difficult due to completely different experimental conditions. As a result, contradictory reports explaining decomposition pathway of Li2O2, number of steps involved, reaction intermediates, kinetics of OER and exact reaction interface remain specific to a particular system and provide inconsistent understanding about the overall RC mechanism of Li–O2 battery. Therefore, in order for realizing practical applications of Li–O2 battery by suppressing the increase of ηRC, it is essential to identify a general low energy decomposition pathway for Li2O2.Nonaqueous lithium–oxygen  compared to that (<0.5 V) of surface‐grown Li2O2. The difference in ηRC can be attributed to intrinsic structural differences in Li2O2 crystals grown by two different routes. The higher ηRC of solution grown Li2O2 can be attributed to relatively higher structural order in Li2O2 crystals grown by chemical disproportionation. In contrast, due to a possible cohesive interaction between the electrode surface and growing crystals, the surface‐grown Li2O2 seems to possess microscopic structural disorder and improved charge transport properties that facilitate a delithiation induced partial solution‐phase oxidation at relatively lower ηRC (<0.5 V). Our observations suggest that ensuring surface‐growth of Li2O2 by strategic modification of electrode and proper choice of electrolyte can simultaneously provide a large DC capacity with high RC efficiency to mitigate the capacity‐rechargeability trade‐off in Li–O2 battery.Herein, we combine several in situ and ex situ experimental measurements with theoretical modelling to systematically decouple the crucial factors affecting the RC of Li–Oies LiO2.9 Howe22.1N‐methyl‐2‐pyrrolidone (NMP) solvent with 20 wt% LITHion (Ion Power) as the binder) pasted on Toray carbon paper , 1 m lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in tetraethylene glycol dimethyl ether  and lithium foil  as the negative electrode. TEGDME has been chosen as the representative solvent due to its acceptable stability in Li–O2 battery, good compatibility with Li metal and intermediate value of the Gutmann donor number  that is known to be a major solvent property influencing the ORR and OER mechanisms. The configuration of a typical Li–O2 coin cell is schematically shown in Figure Coin type CR2032 cells were assembled using Ketjenblack (KB) as the positive electrode prepared by KB‐slurry (in an chanisms.10 The2 cells with KB electrodes where DC of the cells is carried out at four different current densities (jDC) 25, 50, 100, and 200 mA g−1 up to a fixed capacity (Q) of 1 mAh. On the contrary, all the cells are recharged at 100 mA g−1 with cut‐off conditions of 100% RC or 4.5 V versus Li/Li+. DC/RC currents for all the cells are normalized to the mass of KB (≈0.5±0.05 mg cm−2) in the electrode. During discharge, all the cells show a nearly constant‐potential plateau in the DC curve with gradual increase in average ηDC from 223 mV at 25 mA g−1 to 316 mV at 200 mA g−1. A magnified view of the DC curves in the range of 0.4–0.6 mAh is shown in Figure S1 in the Supporting Information. This increase in average ηDC with jDC is characteristic of iR drop due to buildup of insulating solid Li2O2 (having ionic and electronic conductivities in the range of 10−19–10−20 S cm−1 at 25 °C) during DC of the cells. In contrast, despite recharged at the same current density (jRC), all the cells show more pronounced difference in RC behavior. Qualitatively, all the RC curves consist of two stages; an initial sloping region with gradual increase in ηRC usually up to 20–40% state of charge (SOC) followed by a flat high‐voltage plateau above 4 V versus Li/Li+. From Figure 2O2 decomposed at low ηRC (<1 V) depends on the DC current. Nonetheless, the potential in the plateau region remains quite similar. In addition to galvanostatic RC, we have also recharged the cells by anodic linear sweep voltammetry  from 2.95 to 4.7 V versus Li/Li+ after galvanostatic DC at different jDC up to a QDC of 1 mAh. The LSV curves over the full range of potential are provided in Figure S2 in the Supporting Information, whereas a magnified section from 2.95 to 4 V versus Li/Li+ is shown in Figure + and the other at >4.2 V versus Li/Li+, are observed for all the cells. Consistent with the trend in sloping region of galvanostatic curves, the intensity of the LSV peak at ≈3.35 V versus Li/Li+ also increases with increase in jDC. These observations clearly show that larger fraction of Li2O2 deposited at relatively higher jDC can be decomposed at lower ηRC.Figure nsities jDC 25, 50nsities jDC 25, 502.22O2 crystals in Li–O2 battery predicted a transition from large toroids (at lower discharge rates) to thin‐film/nanoparticles (at higher discharge rates) at a critical jDC implying that large toroids and thin‐film/nanoparticles cannot coexist in the DC product. Consistent with this hypothesis, the DC current dependent RC profile of Li–O2 battery was previously attributed to morphological differences observed in Li2O2 deposited at different jDC. These studies demonstrated that thin‐film/nanoparticle‐like Li2O2 formed at higher jDC can be decomposed at lower ηRC sloping region, whereas large toroidal Li2O2 formed at lower jDC gets decomposed at higher ηRC (>1 V) flat plateau. However, it is not understood why a two‐stage RC profile, which according to those reports should be indicative of a mixed morphology in deposited Li2O2, is observed over a wide range of jDC. Moreover, a later study has evinced that a sufficiently dry glyme‐ether based electrolyte is not expected to show any jDC dependent morphological differences in the discharged Li2O2, rather a film‐like Li2O2 is expected at all jDC. As a result, sharp correlation between jDC‐dependent Li2O2 morphology and subsequent RC profile seems ambiguous. These contradictions in reported data lead to inconsistency in understanding the DC/RC process and ask for careful reconsideration of all the qualitative and quantitative parameters that affect the operation of Li–O2 battery. Besides, combining results from different independent experiments to provide a complete idea about DC/RC processes is often difficult due to high sensitivity of Li–O2 electrochemistry towards variation in experimental conditions. Therefore, here we have carried out several in situ and ex situ analyses under the same experimental set up inside a super‐dry room (maximum dew point −60 °C) in order to maintain similar effects by external factors in all of our experiments. The scanning electron microscopy (SEM) images in Figure jDC. Possibly due to low water content (<20 ppm H2O by Karl Fischer titration) in our tetraglyme based electrolyte, no large particle/toroidal Li2O2 is observed at any current density from 25 to 200 mA g−1. In fact, all the electrodes look indistinguishable in the SEM images after DC up to a fixed QDC of 1 mAh. A closer look at the DC products by transmission electron microscopy (TEM), shown in Figure jDC the deposited Li2O2 is composed of small nano‐sized conglomerated particles (30–40 nm diameter) that appear to be a film at low magnification. Another important revelation from TEM images is that the deposition is not conformal, rather particles of Li2O2 are nonuniformly deposited as islands leaving uncovered carbon surfaces around the deposits  studies during RC observed a non‐linear decay of peak intensity for Li2O2 where in the beginning, the peak intensity decays very slowly followed by a rapid decay at high RC potential. The slow decay in XRD peak intensity gives an indirect evidence for decomposition of Li2O2 with smaller domain size (or nano‐sized/amorphous Li2O2) at low RC potential, whereas the sharp decay directly indicates disappearance of larger domains of Li2O2  at flat high‐voltage plateau. This explanation of the two‐stage RC profile seems quite reasonable since higher electronic as well as ionic conductivities of amorphous Li2O2 make it feasible to decompose at low ηRC, whereas crystalline Li2O2 that behaves like an insulator decomposes at high voltage plateau. However, exact mechanism of simultaneous deposition of amorphous and crystalline Li2O2 at a particular DC condition is not understood. Moreover, jDC dependent variation in Li2O2 domain size is also questionable in a sufficiently dry ether‐based electrolyte. Figure −1. Superposition of peaks for both (100) and (101) reflections of Li2O2 discharged at all current densities is quite evident from the XRD patterns. The average domain sizes of Li2O2 discharged at four different jDC vary in the ranges 13.7–14.2 and 9.7–11.1 nm for (100) and (101) reflections respectively as displayed in Figure jDC and domain size can be established from these results due to experimental uncertainty over the small differences in full width half maximum (FWHM) values for corresponding peaks of the four samples. However, an apparent difference in domain size (≈3 nm) between (100) and (101) reflections for all the samples indicates anisotropy in the crystal structure of deposited Li2O2. An important corollary of the XRD data shown in Figure 2O2 deposited at different jDC. Having said that, we do not disregard the effects of Li2O2 crystallinity on the RC profile of Li–O2 battery, we rather emphasize the limitation of ex situ XRD analysis in identifying subtle differences in jDC dependent structural features of electrochemically formed Li2O2. Instability of carbon electrode and nonaqueous liquid electrolyte towards reduced oxygen species is a serious problem in Li–O2 battery leading to buildup of insulating solid parasitic products that increase the impedance of the cell. As the cell‐impedance rises, decomposition of Li2O2 should require high ηRC. In order to investigate any effect of parasitic side reactions on the jDC dependent RC profile of Li–O2 battery, we have quantified the amount of Li2O2 after DC at different jDC up to a fixed QDC of 1 mAh by titrating the electrodes with titanium (IV) oxysulfate (TiOSO4) solution. The yields of Li2O2 after DC have been shown in Figure jDC 25, 50, 100, and 200 mA g−1 respectively. Lower than theoretical yield of Li2O2 is ascribed to parasitic reactions as mentioned above and the Fourier transform infrared spectrum (FTIR) of the KB electrode discharged at 100 mA g−1 in Figure S5 in the Supporting Information shows the existence of lithium carbonate (Li2CO3) and lithium carboxylates (RCO2Li) as the parasitic products. It has to be mentioned here that the parasitic products may have detrimental effect in increasing the ηRC of the cells but quite similar yield of Li2O2 (hence similar extent of parasitic reactions) after DC at different jDC should exhibit comparable effects on ηRC.A nanoscale continuum model for the growth of LijDC dependent difference in average thickness of the deposited Li2O2. A possible correlation between ηRC and thickness of Li2O2 deposit was reported earlier for a flat electrode and it was argued that relatively thinner parts of Li2O2 film could be decomposed at low ηRC. However, interestingly, the positions of anodic LSV peaks (in the ranges of <3.5 and >4 V vs Li/Li+) of a Li2O2 film of <3 nm thickness grown on a flat electrode resemble the positions of the LSV peaks for the oxidation of several tens of nanometer thick Li2O2 deposited on KB electrode as well as on other high surface area carbons reported previously. Therefore, exact correlation between RC profile and Li2O2 thickness seems inconclusive. Besides, the stochasticity of Li2O2 layer thickness due to nonuniform deposition also makes it difficult to establish a quantitative relationship between thickness and RC potential. Nonetheless, based on the observation of nonuniform deposition, we have modelled the nucleation‐and‐growth of Li2O2 in our system in order to provide a quantitative insight about the relation between jDC dependent Li2O2 deposition and subsequent RC profile.The structural, morphological and quantitative analyses demonstrated above could only point out a 2.32 battery is known to involve multiple steps of ORR in nonaqueous conditionsThe DC of Li–Oonditions8c(1)(Here the asterisk (*) represents the species adsorbed on the electrode surface)2O2 as observed in the TEM images. In our model, reaction (1), i.e., formation of surface bound LiO2* is regarded as the first step of ORR. This ORR step is followed by concurrent nucleation (reaction (3)) and growth (reaction (4)) of Li2O2 on the KB surface. Negative free energies for both surface reduction and solution phase disproportionation indicate that simultaneous occurrence of these two reactions is thermodynamically feasible. Here we specify that nucleation refers to surface‐growth that accounts for the coverage of Li2O2 layer on the electrode surface, whereas growth means precipitation of Li2O2 from solution phase on top of the nucleation sites. Li2O2 layer with thickness >10 nm (maximum electron tunneling distance through Li2O2 layer) observed in TEM images also indicates involvement of solution mediated vertical growth of Li2O2 in our cells. The solubility, if any, of intermediate LiO2 in sufficiently dry LiTFSI/TEGDME electrolyte (<20 ppm H2O) has been verified by cathodic LSV using rotating ring disk electrode  under 3‐electrode set‐up. From the LSV data shown in Figure S6 in the Supporting Information, a pronounced reduction current is observed at the KB disk below 2.76 V versus Li/Li+ and simultaneously an oxidizing current is observed at the GC ring held at 3.5 V versus Li/Li+. The reduction current corresponds to ORR at the KB disk while the ring current originates from the oxidation of solvated LiO2 (sol.) that diffuses from the disk to the ring due to the rotation of the electrode. The emergence of ring current confirms the solubility, at least partially, of LiO2 in dry LiTFSI/TEGDME electrolyte (<20 ppm H2O) and the results are consistent with previously reported studies. Higher binding energy of lithium as well as of oxygen leading to preferential nucleation at randomly distributed defect sites results in nonuniform nucleation of Li2O2 on carbon surface. Density functional theory (DFT) calculations showed that the adsorption energies at the defect sites can be higher by ≈1.5 and ≈4.4 eV for lithium and oxygen respectively compared to a perfect defectless site. KB indeed manifests considerable defects and consists of oxygen functional groups that can act as the adsorption sites for both oxygen and lithium. The Raman spectrum of KB in Figure S7 in the Supporting Information shows D‐band  to G‐band  intensity ratio (ID/IG) of 1.06 that certainly confirms presence of defects in the carbon structure. X‐ray photoelectron spectroscopic (XPS) analysis of the KB surface estimates oxygen to carbon ratio (O/C) of 0.027 where oxygen functionalities are identified to be C—O , C=O  and —COO−  groups that are shown by deconvoluted XPS data in Figure S8 in the Supporting Information and the physicochemical properties of KB are summarized in Table S1 in the Supporting Information. According to our proposed mechanism, nucleation of Li2O2 starts preferentially at the defect sites of KB electrode and concurrent nucleation‐and‐growth occurs progressively until complete coverage of Li2O2 layer on the electrode surface. A schematic representation of the proposed DC mechanism is shown in Figure Based on these reaction steps we propose a nucleation‐and‐growth mechanism that agrees with the nonuniform deposition of Lifeasible.6c Herr1, mol m−2 s−1) and (3) In order to model the nucleation‐and‐growth mechanism described above, we have considered the kinetic equations describing the rates of reactions (1) (k1 (m s−1) and k2 (s−1) are the rate constants for reduction of oxygen (O2) and surface adsorbed lithium superoxide (LiO2*) respectively, cO2 (mol m−3) and cLiO2 (mol m−2) are the concentrations of oxygen in the electrolyte and surface concentration of LiO2* on the electrode respectively, α1 and α2 are the charge transfer symmetry factors approximated to be 0.5, η1 and η2 are the overpotentials for the respective steps assumed to be η1 ≈ η2 = ηDC. F (96 485 C mol−1), R (8.314 J mol−1 K−1), and T (298 K) have their usual meaning. The solvation kinetics of LiO2* is represented byrs (mol m−2 s−1) is the solvation rate of LiO2* and ks (s−1) is the solvation rate constant. These equations show that unlike nucleation rate (r2), which depends on the DC overpotential (ηDC), the growth rate depends on the rate of LiO2* solvation (rs) and its supersaturation in the electrolyte. For simplicity, here we consider same rates of reactions at the superficial surface and also inside the bulk of the porous electrode. We further assume that since nucleation occurs by surface reduction, the nucleating layer on the electrode surface can have a maximum average thickness (h0) limited by electron tunneling length and based on this assumption the areal growth rate of Li2O2 can be described aswhere Aelectrode is the electrochemically active surface area  of the KB electrode  of a symmetric KB|KB cell and the calculation is shown in the supporting information), MLi2O2 and ρLi2O2 are the molar mass (45.88 g mol−1) and bulk density (2.31 × 106 g m−3) of Li2O2 respectively and as previously mentioned, h0 is the maximum average thickness of the nucleating layer. Full derivation of Equation (ts) when ALi2O2=Aelectrode. Figure S10 in the Supporting Information shows DC of the cells at four jDC from 25 to 200 mA g−1 up to the cut‐off potential of 2 V versus Li/Li+. A sharp decay in DC potential up to 2 V versus Li/Li+ (sudden‐death) is observed at the end of DC for all the cells in Figure S10 in the Supporting Information. This sudden‐death denotes the complete coverage of the electrode surface by insulating Li2O2 as the DC product. Had the sudden death been due to charge transport limitation through a conformal layer of Li2O2, as observed in flat electrodes, the QDC would be inversely proportional to jDC. Figure S11 in the Supporting Information shows a plot of QDC versus jDC for the KB electrodes where the inverse relationship, shown by the gray solid line, significantly deviates from the experimental data points.where 2.4ts versus ηDC for the cells discharged at jDC from 25 to 200 mA g−1 in Figure ηDC for different jDC are obtained by subtracting the equilibrium potential (E0) of 2.96 V versus Li/Li+ from the DC potentials (E) in Figure 2O2 layer is reported to be ≈10 nm, presence of defects and grain boundaries can increase the tunneling length. Therefore, to fit the experiemntal data in Figure h0 values in the range 10–20 nm. A representative fitting of the experimental data points considering h0 value of 15 nm in Equation is calculated by multiplying ECSA of KB with θLi2O2.We have estimated the amount of charge (QDC/SL) stored in the surface grown layer in terms of deposited Li2O2 from this ALi2O2′ and different h0 values at each jDC from 25 to 200 mA g−1 for a capacity of 1 mAh. Expectedly, QDC/SL increases with jDC as shown in Figure QRC < 4 V) at low RC potential (<4 V vs Li/Li+) during RC  of the cells discharged up to 1 mAh at respective jDC. It is interesting to see that QDC/SL for a layer of 15 nm and QRC < 4 V measured by both galvanostatic RC and anodic LSV are quantitatively comparable. That means the 15 nm thick Li2O2 layer grown by surface reduction decomposes at potential <4 V versus Li/Li+. Conversely, based on this argument, any Li2O2 deposited by chemical disproportionation decomposes at higher potential (>4 V vs Li/Li+) flat plateau of the RC curve. This observation shows an unprecedented deconvolution of RC potentials of Li–O2 battery based on mechanism of Li2O2 nucleation‐and‐growth and provides a quantitative evidence that electrochemically deposited Li2O2 on the surface of the electrode decomposes at low ηRC. It is important to note that this rationale does not specify the effect of any particular morphology of Li2O2 in explaining the RC curve of Li–O2 battery. Rather it just deconvolutes the effects of two concurrent Li2O2‐growth mechanisms on ηRC. However, this general explanation can also be extended to justify the apparent effects of Li2O2 morphology on RC profile of Li–O2 battery. The change in morphology of Li2O2 from large toroids to conformal thin‐film is actually the shift in reaction mechanism from predominant solution phase disproportionation towards electrochemical reduction on the electrode surface. As the morphology changes from toroids to thin‐film, the fraction of surface‐grown Li2O2 increases and consequently more Li2O2 is decomposed at relatively lower ηRC. Therefore, the apparent effect of Li2O2 morphology on RC potential is actually a corollary of fundamental growth mechanism described above. Although the visible morphology of Li2O2 was reported to change sharply at a particular DC condition, here we reiterate that depending on thermodynamic feasibility, the co‐occurrence of the two underlying growth processes is possible over a wide range of DC conditions and the equilibrium shifts quite steadily with small changes in external factors. As a result, DC condition dependent changes in the fractions of sloping region and the high voltage plateau of a two‐stage RC profile can be simply correlated to the change in equilibrium between the Li2O2 growth mechanisms regardless of the bulk morphology. According to the nonuniform nucleation‐and‐growth mechanism explained in Figure 2O2 in the lateral direction maintaining a fixed ratio of the two concurrent growth processes at a particular DC condition. Therefore, the ratio of the sloping region and flat plateau should not change with the DOD if the DC conditions remain unchanged. To verify this hypothesis, we have discharged and recharged three more cells up to QDC of 1.5, 2, and 3 mAh at jDC of 100 mA g−1. The galvanostatic DC/RC curves shown in Figure S13a in the Supporting Information demonstrate similar two‐stage RC profile for all the cells. Quite interestingly, all the RC curves have overlapped upon normalization of QDC in Figure S13b in the Supporting Information. XRD data of the discharged electrodes shown in Figure S14a in the Supporting Information demonstrates gradual increase in Li2O2 peak intensity that signifies the addition of more Li2O2 with increase in DOD. However, no significant growth of individual grains in Li2O2 crystals is observed as shown in Figure S14b in the Supporting Information. This means that the fundamental crystallization mechanism of Li2O2 does not change with DOD and maintains same recharge efficiency. Summarizing all these results it is understood that promoting reaction (3) (surface reduction of LiO2*) over reaction (4) (solution phase disproportionation of LiO2 (sol.)) during DC can contribute to low ηRC of the cell. The competition between these two reactions represented by ks/k2 ratio depends on ηDC (or jDC) and properties of the electrode as well as of the electrolyte. We have already explained the effect of ηDC (or jDC) dependent growth process of Li2O2 on ηRC. Next, in order to prove generality of our hypothesis, we have further modified KB electrode as well as the electrolyte to manipulate the ks/k2 ratio in our system. It is known that defects and surface functional groups on carbon electrode promote surface‐growth of Li2O2, whereas controlled amount of water (H2O) as additive in the electrolyte increases the solubility of LiO2 (sol.) and enhances the rate of disproportionation. We have oxidized KB, hereafter denoted as KBox, in H2SO4 + HNO3 acid mixture to introduce defects and surface oxygen functional groups that are expected to promote surface‐reduction of LiO2*. On the other hand, to decrease the propensity for surface reduction, KB is reduced, denoted as KBr, in Ar+H2 (4% H2) gas mixture at 1000 °C to remove oxygen functional groups. Raman spectra and XPS results of KBox and KBr are shown in Figures S7 and S8 in the Supporting Information respectively. The ID/IG and O/C ratios for both the samples are summarized in Table S1 in the Supporting Information. In short, these results clearly indicate significant enhancement of oxygen functionality (O/C = 0.098) and disorder (ID/IG = 1.54) in KBox and on the contrary, after reduction at high temperature the structure has become more ordered (ID/IG = 0.71) and oxygen content is much decreased (O/C = 0.003) in KBr. KB, KBox, and KBr are discharged and recharged at 100 mA g−1 up to 1 mAh and the DC/RC curves are shown in Figure r < KB < KBox, the sloping region becomes more extended with lower ηRC. The span of sloping region extends from ≈25% SOC in KBr to ≈33% SOC in KB to ≈60% SOC in KBox. The extension of sloping region can certainly be corroborated to enhanced surface reduction of LiO2* with increase in surface‐functionality of the electrode. Morphologies of the as prepared and discharged KBox and KBr electrodes are shown in Figure S15a–d in the Supporting Information. Figure S15b in the Supporting Information shows predominant film‐like deposition of Li2O2 on the superficial surface of KBox electrode. On the contrary, disk‐like Li2O2 with roughly 100 nm diameter and 20 nm thickness, as shown in Figure S15d in the Supporting Information, are formed on KBr electrode. The large sized disk‐like particles give more evidence in favor of a dominant solution phase growth of Li2O2 in KBr electrode. To further enhance the solution phase disproportionation, KB electrodes have been discharged and recharged in LiTFSI/TEGDME electrolyte with different concentrations of H2O added as solvating agent. Toroidal DC products in the H2O containing electrolytes, as shown in Figure S15e,f in the Supporting Information, certainly indicate enhanced solution mediated discharge upon addition of H2O. Figure 2O. In particular the length of the sloping region decreases from ≈33% SOC in dry electrolyte (<20 ppm H2O) to ≈27% SOC in 1000 ppm to ≈18% SOC in 5000 ppm H2O containing electrolytes. To provide a quantitative validation of our hypothesis that solution phase disproportionation decreases the surface coverage of Li2O2, we have calculated the dependence of θLi2O2 on ηDC (or jDC) and ks/k2 ratio for a fixed DC time. In our calculation we have considered commonly observed ηDC values of 0–0.5 V and ks/k2 values 0–105 at a DC time of 10 h. Here ks/k2 = 0 implies exclusive surface reduction of LiO2*, whereas ks/k2 = 105 means greater contribution from solution phase disproportionation of LiO2 (sol.). The contour plot in Figure ηDC, θLi2O2 increases with decrease in ks/k2 value. On the other hand, increase in ηDC that enhances the rate of LiO2* reduction (reaction (3)) also increases θLi2O2 at a fixed value of ks/k2. This clearly demonstrates that with the enhancement of LiO2* reduction, more fraction of electrode surface gets covered with Li2O2. Although increase in θLi2O2 decreases the ηRC as explained above, at the same time the electrode surface becomes passivated with insulating layer of Li2O2 that causes early death of DC resulting in capacity‐rechargeability trade‐off in Li–O2 battery. In Figure ts) at different ηDC and ks/k2 values. The results unequivocally demonstrate that dominant surface reduction of LiO2* leads to early death and hence manifests low QDC.2.52 battery, it is necessary to understand why and how the surface‐grown Li2O2 decomposes at relatively lower ηRC. Therefore, we now focus on understanding and establishing the RC mechanisms that correlate to the decompositions of surface and solution‐grown Li2O2 at different RC potentials. The apparent sharp transition from a sloping trend to a flat plateau in the RC curve  in anodic LSV  with QRC. A high rate (rg/O2) of oxygen evolution at the beginning of RC confirms the decomposition of Li2O2 below 3.5 V versus Li/Li+. However, the rg/O2 value is lower than that calculated for a 2e−/O2 process and more surprisingly, although the RC potential gradually increases, rg/O2 shows a decreasing trend up to ≈15% SOC that corresponds to ≈3.5 V versus Li/Li+. This means either the rate of Li2O2 decomposition (rLi2O2) is lower than expected with a similar decreasing trend or the evolved oxygen somehow remains trapped inside the electrode during the course of RC. In order to quantify rLi2O2, we have titrated the KB electrode at different SOCs by TiOSO4 solution and the results are shown in Figure rLi2O2 values are quite high, roughly close to the theoretical value of 2.38 × 10−8 g s−1, at the beginning of RC and remain significantly higher than rg/O2 up to ≈30% SOC. These results evince that the decomposition of Li2O2 is considerably facile at low ηRC sloping region. The difference between rg/O2 and rLi2O2 can be attributed to generation of highly reactive singlet oxygen (1O2) that reacts with electrode as well as with electrolyte and does not evolve to be detected by OEMS despite continuous decomposition of Li2O2 during RC. These oxidations of electrode and electrolyte by 1O2 produce parasitic side products in the form of Li2CO3, RCO2Li, etc. We have also measured the change in mass of the whole cell in situ during DC/RC. Ideally, any change in the mass of a cell should account for the consumption and evolution of oxygen during DC and RC respectively. We have estimated a ≈1.97 e−/O2 ORR from the mass of consumed O2 during DC of the cell and the trend of O2 consumption is shown in Figure S17 in the Supporting Information. On the other side Figure rw) of the cell during RC. Generally, the value of rw should account for the decrease in mass due to oxygen evolution from decomposition of Li2O2. Interestingly, despite high value of rLi2O2 the graph in Figure rw up to ≈30% SOC and follows the trend of oxygen evolution. This observation clearly indicates entrapment of oxygen and accumulation of insulating side products that balance the decrease in mass due to Li2O2 decomposition. Nevertheless, the non‐catalytic decomposition of Li2O2 below 3.5 V versus Li/Li+ provides optimism about achieving Li–O2 batteries with high round‐trip energy efficiency. Considering the high energy barrier for a direct 2e− oxidation, stepwise 1e− process with low delithiation barrier seems to be a preferred oxidation pathway for Li2O2. Advantageously, defect induced structural disordering of surface grown Li2O2 lowers both the formation and migration energies of Li+ vacancy leading to facile delithiation and formation of non‐stoichiometric Li2−xO2 (0 < x < 1) phase. Despite the energetically favorable delithiation, according to first‐principles calculations, oxygen evolution from the solid Li2−xO2 species remains as the highest energy‐demanding step during Li2O2 decomposition. Alternatively, a concerted delithiation of Li2O2 clusters leading to disintegration and dissolution of the non‐stoichiometric Li2−xO2 at a critical value of x may instead provide a solution‐phase oxidation pathway on the bare electrode surface at relatively low ηRC in the range 0.3–0.4 V. An anodic LSV has been carried out on a pre‐discharged KB electrode using RRDE set up similar to that described for Figure S6 in the Supporting Information. The current responses of the KB disk swept from 2.9 to 4.5 V versus Li/Li+ and the GC ring held at 3.5 V versus Li/Li+ are shown in Figure 2O2 at ≈3.4 V versus Li/Li+ is associated with the occurrence of ring current at the GC ring electrode. This observed ring current directly provides evidence in favor of a partial solution mediated oxidation of the Li2O2 that is deposited by surface reduction. The oxidation in the solution phase is likely to be energetically favorable over the solid phase oxidation due to circumvention of the charge transport limitation of insulating solid Li2O2. Interestingly, the rationale for decomposition of surface‐grown Li2O2 at low ηRC can be extended to explain the apparent electro‐catalytic effect in Li–O2 battery. Any catalyst that shows strong binding towards Li+ and O2 is found to promote surface‐growth of Li2O2 and subsequently exhibits low ηRC. Therefore, regardless of the electrode type, low ηRC is fundamentally linked to the Li2O2 growth on the surface of the electrode during DC. Now the question is why the potential gradually increases and forms a slope before reaching a flat plateau. One possible reason is the structural nonuniformity in the surface‐grown Li2O2 layers that may show a range of RC potential. According to DFT calculations, the interaction between carbon and the Li2O2 layers changes as the distance from the carbon surface increases. As a result, the degrees of disorder in the surface‐grown Li2O2 is also expected to change and may exhibit gradual change in RC potential. Another obvious reason for increase in RC potential is the accumulation of parasitic side products that increase the impedance of the cell. After complete decomposition of surface grown Li2O2 when the RC potential reaches the plateau and ηRC becomes high (>1 V), the solution grown structurally more ordered and stoichiometric Li2O2 starts decomposing. Based on first‐principles calculations, the insulating bulk Li2O2 becomes conductive only when the ηRC is high enough to activate the delithiation process and the charge transport arises from the mixture of Li vacancies and hole polaron's contributions. Therefore, oxidation of Li2O2 in this case also presumably starts with delithiation and progresses via non‐stoichiometric phases. In fact, decomposition of Li2O2 by forming Li vacancy in the crystal was indeed observed by an in situ XRD study during RC of a Li–O2 battery. However, unlike decomposition of surface grown Li2O2, no solution mediated oxidation has been observed for the Li2O2 grown by disproportionation. High enough ηRC to permit successive delithiation and oxygen evolution from the Li2O2 crystal before it disintegrates and dissolves in the electrolyte may be the reason why no solution mediated oxidation is observed in this case. It is interesting that although the galvanostatic RC curve reaches a flat‐potential (≈4.2 V vs Li/Li+) plateau at ≈30% SOC and roughly maintains the potential up to ≈80% SOC, the values of rg/O2 and rLi2O2 are not consistent with this Coulometric data. Both rg/O2 and rLi2O2 are quite low at the beginning of the plateau and steadily increases up to 60–70% SOC where in fact rLi2O2 reaches the theoretical value. These low values of rg/O2 and rLi2O2 can be attributed to delithiation or Li vacancy creation process in stoichiometric and structurally more ordered solution grown Li2O2 at the initial part of the plateau. This is consistent with the oxygen evolution trend during RC of solution grown large toroidal Li2O2 that also shows very low rg/O2 at ≈5–20% SOC as demonstrated in Figure S18 in the Supporting Information. After this initial delithiation process, formation of non‐stoichiometric phases and gradual decrease in both particle size and thickness of Li2O2 deposits leading to improved charge transport can be one of the reasons behind the high OER rate at the later stage of RC. A possible mechanism for the decomposition of Li2O2 is schematically described in Figure 2O2, high amount of carbon dioxide (CO2) is evolved due to decomposition of parasitic products. From the ratio of consumed oxygen to evolved oxygen measured by in situ mass measurement and OEMS respectively, we have calculated an OER/ORR ratio of 0.47. Considering ≈70% yield of Li2O2 after DC, a correction of OER/ORR ratio gives a value of 0.67 which still indicates loss of oxygen during RC due to parasitic reactions. As a result, the overall mass change rate (rw) of the electrode does not follow the trend of rLi2O2 rather shows excellent consistency with the cumulative rate of gas evolution . Detailed mechanistic study identifies a delithiation induced partial solution phase decomposition of surface grown Li2O2 to be a non‐catalytic low energy oxidation pathway. On the contrary, regardless of bulk morphology, solution‐grown Li2O2 decomposes at recharge potential >4 V versus Li/Li+. Therefore, any discharge condition that promotes predominant surface growth of Li2O2 should be beneficial for efficient recharge of the cell. We have further shown that the equilibrium between these two growth processes and the subsequent recharge profile can be controlled by the properties of the electrode as well as of the electrolyte. In short, our results provide important insights into necessary control over fundamental crystallization pathways of Li2O2 by proper choice of electrode and electrolyte to improve the energy efficiency of Li–O2 battery.In summary, we have investigated the recharge mechanism of Li–OThe authors declare no conflict of interest.Supporting InformationClick here for additional data file.