Designing electrolytes with polymerlike glass-formingproperties and fast ion transport at low temperaturesQing Zhaoa

, Xiaotun Liua, Jingxu Zhengb, Yue Dengb, Alexander Warrena, Qiyuan Zhanga, and Lynden Archera,b,1

aRobert Frederick Smith School of Chemical and Biomolecular Engineering, Cornell University, Ithaca, NY 14853; and bDepartment of Materials Science andEngineering, Cornell University, Ithaca, NY 14853

Edited by Alexis T. Bell, University of California, Berkeley, CA, and approved August 27, 2020 (received for review March 10, 2020)

In the presence of Lewis acid salts, the cyclic ether, dioxolane(DOL), is known to undergo ring-opening polymerization insideelectrochemical cells to form solid-state polymer batteries withgood interfacial charge-transport properties. Here we report thatLiNO3, which is unable to ring-open DOL, possesses a previouslyunknown ability to coordinate with and strain DOL molecules inbulk liquids, completely arresting their crystallization. The strainedDOL electrolytes exhibit physical properties analogous to amor-phous polymers, including a prominent glass transition, elevatedmoduli, and low activation entropy for ion transport, but manifestunusually high, liquidlike ionic conductivities (e.g., 1 mS/cm) attemperatures as low as −50 °C. Systematic electrochemical studiesreveal that the electrolytes also promote reversible cycling of Limetal anodes with high Coulombic efficiency (CE) on both conven-tional planar substrates (1 mAh/cm2 over 1,000 cycles with 99.1%CE; 3 mAh/cm2 over 300 cycles with 99.2% CE) and unconven-tional, nonplanar/three-dimensional (3D) substrates (10 mAh/cm2

over 100 cycles with 99.3% CE). Our finding that LiNO3 promotesreversibility of Li metal electrodes in liquid DOL electrolytes by aphysical mechanism provides a possible solution to a long-standing puzzle in the field about the versatility of LiNO3 salt ad-ditives for enhancing reversibility of Li metal electrodes in essen-tially any aprotic liquid electrolyte solvent. As a first step towardunderstanding practical benefits of these findings, we create func-tional Lijjlithium iron phosphate (LFP) batteries in which LFP cath-odes with high capacity (5 to 10 mAh/cm2) are paired with thin (50μm) lithium metal anodes, and investigate their galvanostaticelectrochemical cycling behaviors.

electrolytes | ion transport | lithium batteries | thermal transition |coupled dynamics

Isolating and understanding the various roles played by theelectrolyte solvent in regulating ion transport in bulk liquid

phases and at emergent solid-state phases (solid electrolyte in-terphases, SEI) formed by electrochemical reactions and poly-merization of the electrolyte solvent are important in allrechargeable batteries that utilize aprotic liquid electrolytes(1–4). They are considered crucial steps toward practically rel-evant batteries that use alkali metals as anode to achieve high-energy and/or lower-cost portable storage of electrical energy(5–12). It is known for instance that with minimal assistance fromother electrolyte components (e.g., salts, organic molecular ad-ditives, particulate fillers, etc.), some cyclic ether solvents exhibitsuch high stability in extended contact with working Li and Naanodes that these solvents are often used as benchmark systemsfor evaluating other battery components (e.g., separator, artifi-cial SEI, and current collector architectures) that improve re-versibility of the Li anode (13–15).Electrolytes based on the cyclic ether 1,3-dioxalane (DOL) are

promising for Li metal anodes because DOL is known to facilitatereversible lithium deposition. It is believed that the electrochemicalreduction products (oligomeric/low molecular weight Li+ con-ducting polymer molecules) formed by DOL passivate the Li an-ode, preventing continuous degradation of the electrolyte (16–20).DOL also has a low room-temperature viscosity (0.59 mPa·s) (21),

wide liquidus range (from −95 to 78 °C) at atmospheric pressure,and higher lowest unoccupied molecular orbital energy (1.5 eV)than carbonate solvent (e.g., ethylene carbonate: −0.28 eV), all ofwhich suggest that it is of potential interest in batteries able tooperate stably under extreme weather conditions (22, 23). DOL is,however, conventionally not used as a stand-alone electrolyte sol-vent in Li-ion batteries due to its relatively poor oxidative stability(18) and instability with the most commonly used salt (LiPF6) (24).In Li-S batteries, DOL is nonetheless widely used in combinationwith linear ether and small quantities of LiNO3 salt, because thedimethoxyethane (DME) offers high solubility of electronically in-sulating Li-polysulfides (Li-PS) formed at the cathode, while theDOL, LiNO3, and dissolved Li-PS are thought to form a stable,protective film on the lithium metal, enabling extended cycling of aLi-S batteries with high reversibility (14, 19, 25, 26).Recently, we reported that addition of even small concentra-

tions of Lewis acidic salts (e.g., aluminum triflate [Al(CF3SO3)3])(18), or halides [e.g., AlI3 or AlF3] (27) to a liquid DOL elec-trolyte improves the thermal, mechanical, and electrochemicalstability of DOL by initiating ring-opening polymerization ofDOL inside a battery cell. Because the formed solid-state poly-mer electrolytes originate from liquids able to wet all compo-nents of the cell, the high interfacial resistances typical of solid-state electrolytes are avoided, and the high-oxidation stability isimproved. Motivated by this discovery we herein evaluate DOLas a stand-alone electrolyte solvent for Li metal batteries.

Significance

Liquid electrolytes with thermophysical properties analo-gous to solid polymers, but with exceptional liquidlike ionicconductivities, are formed spontaneously when moderateamounts (≤1 M) of inorganic salts coordinate strongly withsmall molecules in a conventional aprotic solvent. Specifi-cally, we report that electrolytes composed of the cyclic liq-uid ether, dioxolane (DOL), and containing the simple saltLiNO3 are able to completely bypass the liquid → crystallinesolid thermal transition, and to exhibit abnormally high bulkand interfacial ionic conductivities down to temperatures aslow as −50 °C. Through physical, spectroscopic, and ion-transport measurements it is shown that strong interac-tions between LiNO3 and DOL distort bonds in DOL, couplemotions of individual solvent molecules, and lower thethermodynamic activity of the electrolyte.

Author contributions: Q. Zhao and L.A. designed research; Q. Zhao, X.L., Y.D., andQ. Zhang performed research; Q. Zhao, J.Z., Y.D., A.W., and Q. Zhang contributed newreagents/analytic tools; Q. Zhao, X.L., Q. Zhang, and L.A. analyzed data; and Q. Zhao andL.A. wrote the paper.

The authors declare no competing interest.

This article is a PNAS Direct Submission.

Published under the PNAS license.1To whom correspondence may be addressed. Email: laa25@cornell.edu.

This article contains supporting information online at https://www.pnas.org/lookup/suppl/doi:10.1073/pnas.2004576117/-/DCSupplemental.

First published October 5, 2020.

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Comparing the physical, thermal, and electrochemical charac-teristics of DOL-based electrolytes with those of other chemis-tries of contemporary interest, we find that DOL enables fasterelectrode kinetics, including high exchange current density andlow charge-transfer resistance. Additionally, we find that the saltadditive, LiNO3, which cannot initiate ring-opening polymeri-zation of DOL, is able to drive the liquid electrolyte into a glassythermal state in which liquidlike bulk and interfacial ion-transport properties coexist with glassy thermal and rheologicalproperties. Optimization of the LiNO3 content in the DOLelectrolytes yields systems that enable stable lithium stripping/plating over thousands of cycles with CEs exceeding 99%, evenat high lithium-plating capacities (10 mAh/cm2)—among thehighest reported in the literature. We evaluate the DOL elec-trolytes in Li||LiFePO4 full cells in which a thin Li foil (thickness:50 μm) is paired with lithium iron phosphate (LFP) cathodeswith high areal capacity (5 to 10 mAh/cm2) to yield cells withnegative:positive (N:P) electrode capacity ratios of 2:1 and 1:1,respectively.

Results and DiscussionRegulating Kinetics of Electrolytes by the Solvent. Fig. 1A comparesthe ionic conductivities of carbonate- and ether-based electro-lytes composed of four solvents [ethylene carbonate (EC), di-methyl carbonate (DMC), DOL, DME] of contemporaryinterest. The same salt LiFSI is used in all cases because it hasbeen reported to offer high solubility, high ionic conductivity,and forms favorable interphases on lithium metal (28–30). Theconductivities of carbonate-based electrolytes, including typicalcommercial electrolytes (1 M LiPF6 in EC/DMC), are generallylower than 0.1 mS cm−1 at temperatures below −20 °C. Standardconductivities of aqueous KCl solutions are also listed for com-parison (31). The conductivity of the DOL electrolyte is ob-served to exceed 1 mS cm−1 at −50 °C. To our knowledge thelow-temperature ionic conductivity values for the DOL electro-lyte are among the highest reported for lithium batteries (32, 33).In addition, the increasing trends of conductivity at temperatureswell above the melting point are clearly distinguishing. We fitted theexperimental results using the Arrhenius equation σ = σ0e−Ea=KBT

(Fig. 1B). The results reported in Fig. 1C show that ion transport inthe DOL-based electrolytes involves a much lower activation energybarrier (Ea = 0.06 eV) than for any of the other electrolytes studied,including aqueous KCl. Taken together with the high ionic con-ductivity, the results mean that DOL is not only effective in pro-ducing large numbers of dissociated ion pairs, but that the mobility ofthe dissociated ions remains high over a wide temperature range.The Arrhenius preexponential factor (σ0) for the electrolytes are alsoreported in Fig. 1C. The value for the DOL electrolyte is substan-tially smaller than any of the other organic electrolytes, σ0 is relatedto the activation entropy (ΔS) in Eyring theory for liquids; the muchsmaller value for the DOL indicates that ionic motions are lessrandom (i.e., more correlated) than in the other electrolytes.It is conventionally assumed that dissociated ions in a simple

molecular fluid move in tandem with the molecules of the fluid.Comparison of the activation energy deduced from solvent flu-idity (1/viscosity) data provides a simple means of testing thishypothesis. We plot the shear viscosity of all electrolyte solventsused in the study in SI Appendix, Fig. S1A and report the acti-vation energies obtained by fitting the data using the Arrheniusexpression in SI Appendix, Fig. S1B. The results show that whilethe barrier for relative motion of DOL molecules is approxi-mately half of that for H2O, the activation energy for pure DOLis unremarkable relative to the other ether and carbonate liquids.This conclusion is evidently quite different from the one reachedby comparing the temperature-dependent viscosity for DOL andDME containing the LiFSI salt (Fig. 1D). At the same LiFSIconcentration, the viscosity of the DOL electrolyte is not only

substantially lower, but the temperature dependence is markedlyweaker. This is in accordance with the trends seen in theconductivity.The movements of molecules in liquids, especially at low

temperature, can be inferred from the dipole relaxation time (τ),determined from dielectric spectroscopy. We measured the di-electric spectra of the electrolytes used in the study and comparethe dielectric loss moduli (M″) in Fig. 1 E and F and SI Appendix,Fig. S2. At low temperature, clear and distinct loss maximaemerge in many of the electrolytes; the maxima shift to lowerfrequencies as temperature decreases. This behavior is well-known in soft materials and is associated with the slowdown ofmolecular relaxation dynamics. Comparison of M″ spectra forthe commercial carbonate (Fig. 1E) and DOL electrolyte(Fig. 1F), as well as with the full set of electrolytes (SI Appendix,Fig. S2) reveals that at temperatures above −60 °C, the DOLelectrolyte shows no obvious dielectric loss maxima, indicatingthat the selected frequencies (1 to 107 Hz) are too low to capturerelaxations in the DOL electrolyte. The Havriliak−Negami(H-N) equation can be used to fit the dielectric loss spectra andto extract characteristic relaxation times, τ, (Fig. 1G) for thevarious electrolytes (34). It is conventionally thought that when τis over 0.1 s in an electrolyte, motions are too slow to supportbattery operations. It is remarkable that τ for the DOL electro-lyte only crosses this threshold at −80 °C, and extraordinarily lowtemperature! In order to ascertain that these observations areelectrochemically meaningful, we studied lithium stripping andplating in the DOL cells down to −50 °C. The results reported inSI Appendix, Fig. S3 clearly show that whereas electrochemicalcells containing the DOL electrolytes can undergo fast lithiumstripping and plating at −50 °C, the Li electrodes are completelyinactivated in the commercial carbonate electrolyte at −50 °C.Electrochemical impedance spectroscopy (EIS) measurements(SI Appendix, Fig. S4) explain the observed behaviors—whereasthe charge-transfer resistance at −50 °C increased to levels onthe order of tens of thousands for the commercial electrolyte, thevalue is less than 100 Ω for the DOL electrolyte.The large differences in ion transport and thermal properties

observed in the DOL electrolytes motivated us to analyze otherkinetic properties of the materials. Exchange current densities(I0) of the electrolytes at a lithium metal electrode were obtainedfrom the respective Tafel plots (Fig. 2A and SI Appendix, Fig.S5). The results reported in Fig. 2B show that the I0 values forthe DOL electrolytes are much larger than any of the otherelectrolytes studied. These results are consistent with thoseobtained from EIS analysis of symmetric lithium cells (Fig. 2C).All Nyquist plots show a single semicircle shape, which can beeasily fitted a bulk resistance in series with a charge-transferresistance that is in parallel with a double-layer capacitance. Itis shown in Fig. 2D that the DOL electrolyte exhibits the smallestcharge-transfer resistance (∼10 Ω), indicating the DOL electro-lyte facilitates fast interfacial Li+ transport. Measurements of theoverpotential required to galvanostatically strip and plate Li(Fig. 2E and SI Appendix, Fig. S6) in the various electrolytesprovide a more straightforward method to assess differences intheir ion-transport characteristics aggregated over the electrolytebulk and interphases formed on the Li electrode. It is againapparent that cells that employ the DOL-based electrolytes shownoticeably lower polarizations (<30 mV) than those based on theother electrolytes. The Li+ transport number for the electrolyteswas quantified using direct current (DC) polymerization measure-ments. The results reported in Fig. 2F show that all electrolytesbased on the LiFSI salt have nearly the same transference number(∼0.6), which is higher than the corresponding values for thecommercial LiPF6-based electrolytes (35).We investigated the reversibility of Li stripping and plating

processes using Li||Cu electrochemical cells (SI Appendix, Fig.S7). In these experiments, a certain capacity of lithium (typically

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1 mAh cm−2) was deposited on Cu first and subsequently strip-ped from the Cu electrode. The CE is calculated as the per-centage of the original Li electrodeposit plated on Cu that isremoved in the subsequent Li stripping cycle. Electrolytes basedon DOL exhibit both low polarization (∼40 mV) and very highCE values (∼99%) over 300 cycles (SI Appendix, Fig. S7E). Thevalues compare favorably with those obtained for the EC elec-trolyte (∼99% for 100 cycles), DME electrolyte, and DMCelectrolyte (CEs are reduced to <50% before 100 cycles). Re-markably, even at high current density (10 mA cm−2) and high Liplating capacity (3 mAh cm−2) (SI Appendix, Fig. S8), the DOL-based electrolytes still exhibit high CEs over 95%.

Regulating Stabilities of Electrolytes by Salt Addition. LiNO3 hasbeen reported in numerous studies to be an effective salt additivein improving the reversibility of Li metal electrodes in liquidelectrolytes (14, 19). It is commonly assumed that the beneficialeffects of LiNO3 come mostly from its role in regulating chem-istry and transport properties of interphases formed on the Lielectrode. Various literature has also explained the beneficialeffects of LiNO3 in terms of its influence on Li surface chemistryand solvation of intermediates formed at Li electrodes in liquidelectrolytes (14, 29, 36–38). Because Li already exhibits excep-tional reversibility in the DOL electrolyte, we hypothesized thatthese electrolytes would provide a good testbed for under-standing how LiNO3 works. Our results reported in Fig. 3 suggest

that in addition to the widely held view that LiNO3 functions asan interphase agent, LiNO3 also markedly alters the solvationenvironment and bulk electrolyte properties, particularly athigher concentration. The bulk, DC ionic conductivity of theLiFSI/DOL electrolytes containing LiNO3 decreases progres-sively as the LiNO3 content rises. However, even in electrolytescontaining 1 M LiNO3, the conductivity remains very high (∼1mS cm−1) at −50 °C (SI Appendix, Fig. S9). The activation energy(Ea) increases with the addition of LiNO3 and the pre-exponential factors (σ0) are also larger. To understand thesource of these behaviors we determined the bulk mechanicalmodulus of the DOL electrolytes with/without LiNO3 by mea-suring the speed of sound in the respective liquids. The resultsshown in Fig. 3B clearly show that LiNO3 increases the bulkmodulus of electrolytes, meaning that the compressibility of theliquid DOL is lowered by LiNO3, but the modulus does not showany obvious change in its temperature dependence. The resultssuggest that LiNO3 reinforces the electrolyte bulk perhapsthrough its ability to strongly interact with and electrostaticallycross-link DOL molecules. This sort of cross-linking would notonly slow down bulk ionic motions, as is already evidenced in theDC conductivity, but should also impact thermal properties ofthe electrolytes. We performed differential scanning calorimetry(DSC) to investigate the latter effect and the results are reportedin Fig. 3C. The LiFSI/DOL electrolyte crystallizes at −87 °C,which is close to the melting point of DOL (−95 °C). Addition of

Fig. 1. Ion-transport characteristics in electrolytes comprising different solvents. (A) DC conductivity versus temperature. The commercial electrolyte (Com.Ele.) is composed of 1 M LiPF6 dissolved in a symmetric EC/DMC (1:1 by volume) solvent blend. The electrolytes designated DME, DOL, DMC, and EC are eachcomposed of a 2 M LiFSI salt dissolved in the respective solvents. (B) Arrhenius fitting of the measured ionic conductivity versus temperature at relatively hightemperatures. Results for aqueous KCl electrolytes are also listed to facilitate comparisons. (C) Arrhenius preexponential factor (σ0) and Ea diffusion calculatedfrom the fits in B for all electrolytes used in the study. (D) Arrhenius fits of the shear viscosity versus temperature data for DOL and DME-based electrolytes.Dielectric loss modulus spectra as a function of temperature for electrolytes composed of (E) 1 M LiPF6 in EC/DMC and (F) 2 M LiFSI in DOL. (G) Variation incharacteristic relaxation times (τ) as a function of temperature. τ was obtained by fitting the frequency-dependent dielectric modul measured at eachtemperature using the H-N function.

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0.25 M LiNO3 causes the peak to shift to −91 °C and to disap-pear entirely in electrolytes LiNO3 to 0.5 M; in these systems, thecrystallization transition is replaced by a glass transition regionat −121 °C (Tg). However, in the subsequent heating process, amassive, sharp reentrant freezing peak located at −90 °C is ob-served, which is followed by a normal melting peaking locatedat −58 °C. The latter sequence of observations is well-known inglassy materials that undergo a cold crystallization process whenquenched too quickly into the glassy zone (39, 40). We note fi-nally that when the LiNO3 content increases to 1 M, only a re-versible glass transition with Tg of −119 °C is observed. Theobservation of cold crystallization and vitrification of DOLelectrolytes in the presence of LiNO3 is well-known in studies ofsupercooled semicrystalline polymers, but is not expected for asimple molecular fluid such as DOL. The stability of electrolytesagainst polymerization was assessed using a combination of vi-sual interrogation of their flowability (SI Appendix, Fig. S10) andNMR analysis (SI Appendix, Fig. S11). The results show that theelectrolytes do not polymerize to any meaningful extent on thetimescales of our experiments. Therefore, it is thought to provideadditional evidence that the LiNO3 is strongly coordinated withDOL molecules, constraining their motions and producing solid,polymerlike thermal characteristics in a liquid electrolyte.The interactions of different ingredients in the electrolytes are

considered to generate because of electron-lacking N-atom inLiNO3, which will attack both DOL and LiFSI in the electrolytes(Fig. 3D). This hypothesis is further proved by vibrational spec-troscopy. Attenuated total reflection–Fourier-transform infraredspectroscopy (ATR-FTIR) measurements (Fig. 3E) show thataddition of LiFSI to DOL produces multiple changes. First, thevibration associated with the H-C-C deformation splits into twopeaks (41, 42), indicating that cyclic DOL molecules are highlystrained, and undergo a shape change as a result of their inter-actions with LiFSI. This interaction increases with addition ofLiNO3. Second, the IR peak located at about 920 cm−1, whichcan be assigned to the C-C + C-O ring stretching vibration, canstill be observed when LiFSI is present in the electrolyte, butgradually disappears and shifts to lower wavenumber as the con-centrations of LiNO3 is increased. These changes are amplified inthe corresponding Raman spectra (Fig. 3F). The C-O-C bendingpeaking associated with DOL and the S-N-S (775 cm−1) vibration

of LiFSI combine into one broad peak after adding LiFSI into DOl(43). Upon addition of LiNO3, this peak progressively shifts tohigher energy, which indicates the combined interactions of threecomponents. The phenomenon is also observed in highly concen-trated LiFSI/DMC electrolytes, when DMC solvent is coordinatedwith salt (44). Finally, the Raman band at 1,070 cm−1 associatedwith NO3

− in anhydrous LiNO3 shifts to 1,053 cm−1 in the presenceof DOL, particularly when the concentration of NO3

− reaches0.5 M. This change has been previously observed in phases ofLiNO3·xH2O (45), indicating that DOL molecules are also highlycoordinated with LiNO3. The cumulative results from our vibra-tional spectroscopy therefore appear to conclusively demonstratethat both the structure and dynamics of the DOL molecule arealtered by LiNO3.

Electrochemical Performances of Polymerlike Electrolytes. Li||Cuelectrochemical cells were assembled to study the efficiency oflithium plating/stripping in LiFSI/DOL electrolytes containingLiNO3. Fig. 4B reports the CE values in electrolytes with dif-ferent amounts of LiNO3. The CEs increase from 98.9 to 99.1%in electrolytes containing 0.5 and 1 M LiNO3, respectively. Thecorresponding galvanostatic lithium plating/stripping profiles forthe electrolytes with 0.5 M LiNO3 are provided in Fig. 4A. Theresults show that small levels of polarization are maintained forover 900 cycles. Increasing the capacity of lithium platting to 3mAh cm−2, the LiFSI/DOL electrolytes containing 0.5 and 1 MLiNO3 exhibit high CEs (>99%) for over 350 cycles (Fig. 4 C andD). In order to determine how the reversibility of Li changes athigh lithium deposition capacity, we deposited 10 mAh cm−2 ofLi at a current density of 5 mA cm−2 on a commercial three-dimensional carbon cloth current collector and studied its cyclingbehavior (Fig. 4 E and F). A high CE of 99.3% is retained for atleast 100 cycles. The uniform deposition of lithium contributes tothe high CEs (SI Appendix, Fig. S12). We progressively increasedthe Li plating rate from 1 to 9 mA cm−2, with the capacitymaintained at 3 mAh cm−2, and measured the CE values. Again,high CE values (>99%; SI Appendix, Fig. S13) are observed.When the concentration of LiFSI is reduced to 0.5 or 1 M, theLi||Cu electrochemical cells still show high CE of ∼99% after200 cycles when LiNO3 is added (0.5 to 1 M) (SI Appendix, Fig.S14). We note further that similar benefits of LiNO3 are

Fig. 2. Electrochemical kinetics of electrolyte comprising different solvents. (A) Tafel plots derived from cyclic voltammetry measurements. (B) Exchange-current densities (I0) calculated from Tafel plots. The exchange-current density was calculated by fitting the linear region of Tafel plots (from 50 to 30 mV,and −50 to −30 mV). (C) EIS analysis. (D) Charge-transfer resistance (Rct) calculating by fitting the data in C with an equivalent circuit model (D, Inset). (E)Voltage response obtained from 10-min discharge and 10-min charge cycles at 1 mA cm−2. (F) Comparisons of Li+ transport number (τLi+) computed from DCpolarization measurements at 5 mV using the Bruce–Vincent method. All of the data are obtained from symmetrical lithium batteries, in which the glass fiber(grade A) is used as separator to guarantee wetting by all electrolyte solvents used in the study.

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observed in the EC, DMC, and DME electrolytes (SI Appendix,Fig. S15), but again the most sustained benefits are for the ether-based electrolyte. Meanwhile, when other salts [LiCF3SO3 andLiN(CF3SO2)2-LiTFSI] are dissolved in DOL to study lithiumdeposition, the CEs are also largely improved with the additionof LiNO3 (SI Appendix, Fig. S16).Both salts (LiFSI and LiNO3) that facilitate high Li revers-

ibility in DOL-based electrolytes have low oxidation stability,which would normally be thought of as disadvantageous becausethey would exacerbate the already poor stability of DOL incontact with a high-voltage cathode (18). Surprisingly, we ob-serve that the strong coordination of DOL with LiNO3 enhancesboth the reductive and oxidative stability of DOL in an elec-trochemical cell (SI Appendix, Fig. S17). The results reported inSI Appendix, Fig. S17 A and B show that the LiNO3 significantlylowers the initial discharge capacity in Li||Cu cells, which is re-lated to SEI formation/parasitic degradation of the electrolyteprior to the onset of Li metal plating. On the basis of our earlierfindings, we hypothesize that LiNO3 coordination lowers thereactivity of DOL at all potentials.We evaluated the electrochemical performance of lithium

batteries containing the 2 M LiFSI-0.5 M LiNO3-DOL electro-lyte. These assessments were deliberately performed under strictconditions (e.g., thin lithium with thickness of 50 μm paired withhigh-loading commercial cathodes). The CE of lithium plating/stripping measured in thin-Li||Cu electrochemical cells is similarto those observed in cells in which the typical large excess (thicklithium) are used (Fig. 5 A–C). We assembled Li||LFP batteriescontaining cathodes with different active material mass loadingsto initiate and understand failure modes of the DOL electrolytes

in practical electrochemical cells. In low mass-loading LFPcathodes, a capacity retention of ∼80% is observed after 500cycles (SI Appendix, Fig. S18). At a higher LFP loading of32 mg cm−2, the Li||LFP cells exhibit capacities of 4.6, 4.0, 3.3,2.8, and 2.3 mAh cm2 at current density 0.5, 1.0, 1.5, 2.0, and 2.5mA cm−2, respectively (Fig. 5D). Fig. 5E shows that when thecurrent density decreases to 0.5 mA cm−2, the discharge capacityincreases to 4.6 mAh cm2, indicating the high current toleranceof optimized electrolytes. This battery can also operate at a tem-perature of −30 °C (SI Appendix, Fig. S19). Fig. 5F shows prelim-inary cycling performance of Li//LFP battery at different anode-to-cathode capacity ratio (N:P = 1:1 or 2:1). The capacity retentionsare close to (N:P= 1:1) or higher than 80% (N:P = 2:1) after 50 cyclescompared to the third cycle. The performance can be improved byfurther optimizing the structures of cathodes.

ConclusionIn summary, we report that electrolytes composed of the cyclicether, DOL, exhibit exceptional physical, thermal, and electro-chemical characteristics, including high bulk and interfacial ionicconductivities down to −50 °C and low Ea barriers for iontransport. Above a threshold concentration of ∼0.5 M, additionof LiNO3 to a DOL-based electrolyte causes the electrolyte totransition to a highly correlated but amorphous state in whichcrystallization is completely arrested, molecular relaxation slowsdown, but the high ionic conductivities are preserved. By meansof physical, spectroscopic, and ion-transport measurements wefind that strong interactions between LiNO3 and DOL distortbonds in DOL, couple motions of individual molecules, but donot produce ring opening. The resultant structured liquid

Fig. 3. Transport, thermal, and spectra properties of DOL-based electrolytes containing LiNO3 as additive. (A) Arrhenius plots of ionic conductivity versustemperature for electrolytes with different concentrations of LiNO3. (Inset) The values of σ0 and Ea provided were obtained by fitting the experimental data(symbols) using the Arrhenius ion-transport model (lines). (B) Effect of LiFSI and LiNO3 salt content on the temperature-dependent bulk modulus of DOL-based liquid electrolytes. (C) Heat flow as a function of temperature obtained from DSC analysis of the same electrolytes in B. (D) Proposed interactions of theelectrolyte components. (E) ATR-FTIR and (F) Raman spectra of designed electrolytes. The electrolyte designated as DOL Ele. is a 2 M solution of LiFSI in theelectrolyte solvent DOL Sol. We added different concentrations of LiNO3 (0.25, 0.5, 1 M) to the DOL Ele. to explore the effect of LiNO3 on ion transport,compressibility, and thermal properties of the electrolytes.

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electrolytes facilitate highly reversible lithium stripping/plating,over thousands of cycles with CE exceeding 99%, even at Lithroughputs as high as 10 mAh cm−2. An initial assessment of theDOL electrolytes in Li||LiFePO4 full cells with N:P ratios of 2:1and 1:1 demonstrates the potential of DOL electrolytes in Libatteries with wide temperature- and voltage-stability windows.

MethodsMaterials Preparation and Characterizations. All electrolytes are prepared inargon glove box (Inert Inc), in which the content of O2 and H2O content arelower than 0.5 ppm. The company and purity of raw materials: lithiumbis(fluorosulfonyl)amide [LiFSI, 99% purity, (Oakwood Products Inc. [OAK])],lithium nitrate (LiNO3, 99.99% metal basis, Chem-Impex Int’l. Inc.), 1,3-Dioxolane (DOL, 99.8%, Sigma-Aldrich), DME (99.5%, Sigma-Aldrich), DMC(≥ 99%, Sigma-Aldrich), EC (99%, Sigma-Aldrich), 1.0 M LiPF6 in EC/DMC =50/50 (vol/vol) (battery grade, Sigma-Aldrich). DC conductivity, dielectric lossmodulus, and EIS test are performed using a dielectric/impedance spec-trometer (Novocontrol Broad band). For conductivity and dielectric lossmodulus tests, the temperature increases from −90 to 70 °C and the test isoperated every 10 °C. The viscosity is tested using Vibro SV-10 Viscometer.DSC is tested using TA Instruments DSC Q2000. The sound speeds of elec-trolytes are measured by DMA density meter. Renishaw inVia confocalRaman microscope is used for Raman tests of electrolytes (excitation wave-length: 785 nm). ATR-FTIR spectra were conducted using a Thermo Scientificspectrometer. Spectra (1H NMR and 13C NMR) are performed by dissolvingelectrolytes in dimethyl sulfoxide-d6. All of the experiments are conductedless than 5 d after the preparation of electrolytes.

Fitting of DC Ionic Conductivity. Arrhenius equation σ = σ0e−Ea=KBT (Fig. 1B)was used to fit the conductivity of electrolytes, where σ0 is preexponential

factor, KBis Boltzmann constant, Ea is the activation energy, T is theabsolute temperature. Further studies on preexponential factor arebased on the Eyring model, in which the transport is dominated bymolecular contacts. The ionic conductivity follows the Nernst–Einsteinequation,

σ = 2e|z|FC0

KBT× D, [1]

in which the diffusivity follows the Stokes–Einstein equation,

D = KBT6πμa

, [2]

and μ can be calculated by the Eyring model,

μ = ra( )2NAh

Vexp

ΔG0

RT[ ]. [3]

In Eqs. 1–3, KB is Boltzmann’s constant, T is absolute temperature, F isFaraday constant, |z| is valence of ions, C0 is the concentration of ions, a isthe radius of spherical ions, μ is the viscosity of electrolyte, and r isthe distance between ions, V is the volume of the spherical ions, NA isAvogadro constant, h is Planck constant, R is the gas constant, and Gibbsfree energy ΔG = ΔH − T × ΔS. Thus, as most of the parameters are con-stant for various electrolytes, the ionic conductivity can be roughlysimplified as

σ = AeΔS=R × e(-Ea)=RT .

Therefore the preexponential factor is related with entropy; the low pre-exponential factor indicates the low entropy.

Fig. 4. Electrochemical stability of electrolytes containing LiNO3 as an additive. (A) Li stripping and plating profiles for Li//Cu electrochemical cells and (B) the cor-responding CEs as a function of cycle number at a fixed current density of 1 mA cm−2 and for a Li plating capacity of 1 mAh cm−2 per cycle. (C and D) Results fromanalogous experiments to A and B) except a higher Li plating capacity per cycles (3 mAh cm−2 at 1 mA cm−2) was employed in the measurements. (E and F) Gal-vanostatic cycling characteristics of high-capacity/high-rate Li half-cells composed of Li//nonplanar (carbon cloth) current collector. The Li plating capacity per cycle is 10mAh cm−2 and the current density is 5 mA cm−2. The profiles for A, C, and E are for the electrolyte previously designated DOL Ele. fortified with 0.5 M LiNO3.

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Calculation of Relaxation Times. The experimentally measuredM″ spectra werefitted with the generalized H-N model near the maximum region as follows:

M″ = M∞Ms[(M∞ −Ms)sin λϕ]Aλ

M2s A2λ + 2Aλ(M∞ −Ms)Ms cos λϕ + (M∞ −Ms)2

,

where

A = [1 + 2(ωτ)1−α sin πα

2+ (ωτ)2(1−α)]1=2,

ϕ = tan−1[ (ωτ)1−α cos απ21 + (ωτ)1−α sin απ

2

].Ms and M∞ are the electric storage modulus M′ at ω→0 and ω→∞, re-spectively; τ is the relaxation time, and α and λ are exponents that describethe symmetric and asymmetric distribution of the relaxation times. The fit-ting was performed with lsqcurvefit function using MATLAB.

Calculation of Bulk Modulus. The bulk modulus of electrolytes is calculated byNewton–Laplace equations.

c =Ks

ρ

√,

where c is the speed of sound and ρ is the density of electrolytes.

Electrochemical Cells Test. Coin 2032 cells model were assembled forelectrochemical tests. Lithium foil acts as anode and Al2O3 coated Celgardacts as separator. (In cells using LiFSI/EC electrolyte and symmetrical Li//Licells, glass fiber acts as separator to guarantee the full wetness.) Gal-vanostatic lithium stripping/plating tests (Li//Cu [or carbon cloth] elec-trochemical cells and Li//Li symmetrical electrochemical cells) andgalvanostatic discharge/charge tests (Li//LFP battery) were operated atroom temperature using Neware battery tester. The transference num-ber of Li+ is tested by a DC polarization according to the reports usingsolartron battery tester (35). Cyclic voltammetry profiles are obtainedfrom A CH 600E electrochemical workstation. Low-loading LFP cathodeswere prepared by mixing LFP powder, super P conductivities, and poly-vinylidene fluoride binder with weight ratio of 80:10:10. High-loadingLFP cathode on carbon cloth is prepared according to our previousreport (46).

Data Availability. All study data are included in the article and SI Appendix.

ACKNOWLEDGMENTS. This work was supported by the Department ofEnergy Basic Energy Sciences Program through Award DE-SC0016082 andthe NSF through Award DMR-1609125. The characterizations of electronimages are supported by the Cornell Center for Materials Research withfunding from the NSF Materials Research Science and Engineering Centers(MRSEC) program (Grant DMR-1719875).

Fig. 5. Electrochemical characterization of Li||Cu and Li||LFP cells employing thin (50-μm) lithium foil as the negative electrode and thick LFP as the cathode.(A and B) Galvanostatic Li stripping and plating profiles as a function of cycle index for Li||Cu cells at a current density of 1 mA cm−2 and Li capacity per cycle of5 mAh cm−2 (A) and 1 mAh cm−2 (B). (C) CEs deduced from the cycling profiles at 1 mA cm−2, as a function of cycle index. (D) Discharge/charge profiles forLi||LFP electrochemical cells at current densities ranging from 0.5 to 2.5 mA cm−2. (E) Cycling performance of Li||LFP cells at different current density. (F)Cycling performance and corresponding CEs of Li||LFP cells with different N:P ratio. The N:P ratio is varied by changing the mass loading of LFP in the cathode.The results denoted in red correspond to Li//LFP cells in which the 50-μm Li foil is paired with a commercial LFP cathode with mass loading ∼32 mg cm−2, for anN:P ratio of ∼2. Cells were subjected to a brief formation process in which they were discharged/charged at a relatively low current density of 0.5 mA cm−2 forthe first two cycles, whereafter they were galvanostatically cycled at 1.5 mA cm−2. The results denoted with blue symbols utilized LFP with double thestandard active material mass loading (i.e., an LFP mass loading of approximately 64 mg cm−2), corresponding to the theoretical N:P ratio of unity. For thesemeasurements, the formation process consisted of discharging and charging the cells at a current density of 1 mA cm−2 for the first two cycles, followed byextended cycling at a fixed rate of 2 mA cm−2.

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