Chapter 3. Design of organic single-ion conductor based on electrophoretic
3.2. Interpenetrated polycations for anode-free lithium storage
3.2.3. Results and discussion
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Figure 3-10. Structural characterizations of interpenetrated polycations. (a) FTIR spectra of QSIC- B(Ph)4 and its components. (b) XPS N1s spectrum of the QSIC-B(Ph)4 showing pyrrolidinium cations.
(c) SEM and EDS elemental mapping images (C, O, N, B, and P atoms) of the QSIC-B(Ph)4. (d) Cross- sectional SEM image of QSIC-B(Ph)4. (e) High-magnification SEM images of QSIC-B(Ph)4 after solvent extraction, showing its semi-IPN structure and open pores. (f) Pore size distribution determined by mercury porosimeter for QSIC-B(Ph)4 after solvent extraction.
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The ion conduction behavior of the QSIC was examined as a function of its composition. Both ion conductivity and Li+ transference number (tLi+) tended to increase with the cationic PU content (Figure 3-11a). Further increase of cationic PU content more than 50% failed to fabricate a self-standing film. At a composition ratio of CPU-B(Ph)4/ETPTA = 50/50 (w/w), the QSIC showed the highest room temperature ion conductivity (4.4 × 10-3 S cm-1) and tLi+ (0.84), indicating that cationic PU makes selective transport of Li+. The single Li+ transport behaviors of QSIC were further investigated. As a model study, QSICs were prepared with various cationic PUs as a function of ionic radius of monovalent counter anion in the same composition ratio of cationic PU/ETPTA = 50/50 (w/w). It is well known that a large size of counter ion induces high solubility of that salt in a polar solvent and the same charge conditions.16,17 The tLi+ of resultant QSICs tended to increase until 0.84 with the counter anion size (Figure 3-11b), implying a significant role of the cationic PU and its counter anion to manipulate the Li+ transport. Figure 3-11c showed the respective ion conductivities of QSIC-B(Ph)4 as a function of LiPF6 concentrations. The responses of ion transport exhibited a conductivity plateau for concentrations below 10-2 M, which is explained by the surface charge-governed nanofluidic transport.18 At the surface- governed transport region (10-4 M), about 100 times higher conductivity was obtained for the QSIC (1.5
× 10-4 S cm-1) compared with that of bulk solution (1.6 × 10-6 S cm-1). These results demonstrate that charge screening nanofluidic in the electric double layer occurs at the vicinity of the cationic PU and its effectiveness could be controlled precisely.
To highlight the fast Li+ transport behavior through the QSIC-B(Ph)4, its Arrhenius plot of Li+ conductivity (estimated by ion conductivity × tLi+) was compared with those of liquid electrolyte with and without conventional polyethylene separator (Figure 3-11d and Figure 3-12). Due to the highly active state of Li cations, the QSIC-B(Ph)4 showed highest Li+ conductivity over a wide range of temperatures with the lowest activation energy compared with those of liquid electrolytes, even under a separator-free condition (black dots in Figure 3-11d). Note that higher ion conductance of QSIC- B(Ph)4 (8.8 × 10-6 S vs. 3.3 × 10-7 S of liquid electrolyte w/ separator) makes it possible to improve many kinds of polarizations and performance of batteries.
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Figure 3-11. Accelerated Li+ transport behaviors. (a) Ion conductivity and Li+ transference number of QSCI-B(Ph)4 as a function of its composition ratio. (b) Li+ transference number as a function of effective charge density of cationic PU. (c) Ion conductivity of QSIC-B(Ph)4 and control electrolyte as a function of the salt concentration. (d) Arrhenius plot of the QSIC-B(Ph)4 (vs. liquid electrolyte with and without membrane).
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Figure 3-12. Total ion and Li+ conductivities of (a) QSIC-B(Ph)4, liquid electrolyte (b) without and (c) with a conventional polyethylene membrane. Li+ conductivity was estimated by total ion conductivity
× Li+ transference number.
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Such unusual ion-regulating phenomena of the QSIC-B(Ph)4 were examined by FTIR analysis.
P-F vibrations of PF6 anion are well-known to be highly sensitive to the ion association status. FTIR spectra of liquid electrolyte and QSIC show characteristic peaks at 840 and 838 cm-1, respectively, assigned to the P-F vibration (Figure 3-13a), exhibiting an appreciable shift of the P-F peak in QSIC.
This result indicates that the PF6- anions could be electrostatically trapped by the positively charged pyrrolidinium groups of the PU.19,20 The local chemical environments of Li+ in QSIC were further investigated using nuclear magnetic resonance (Figure 3-13b). A singlet 7Li peak at -0.42 ppm was observed in the spectrum of control electrolyte. In contrast, a narrow downshifted 7Li signal at -0.36 ppm was observed in QSIC-B(Ph)4, which is attributed to the presence of more mobile 7Li species resulting from the effective dissociation from the contact ion pairs and cation-anion aggregates.21,22 In addition, shoulder peaks are present in the 7Li spectrum of QSIC, indicating that Li+ may exist in different local chemical environments compared to that of control electrolyte. This point is strongly supported by the FTIR spectra as shown in Figure 3-13c. The characteristic IR peak of carbonyl vibrations of carbonates is highly sensitive to coordination structures.23,24 The control electrolyte showed 2 sharp peaks at 1770 and 1791 cm-1, which are assigned to the Li+-coordinated and free carbonates, respectively. However, when Li salt was introduced into the QSIC, the intensity of coordinated carbonyl vibration increases with a clear downshift to 1764 cm-1. These results verify that cationic PU electrostatically restricts mobility of PF6- anions and generates highly dissociated Li salts, resulting in large density of the ideally solvated Li+ compared to the Li salt in conventional liquid electrolyte. The low dissociation degree of conventional liquid electrolytes16 causes an aggregated solvation structures and sluggish Li+ mobility, showing relatively higher mobility of anions with unwanted interfacial side reactions25 and polarizations of the battery26,27 (Figure 3-12). Meanwhile, we prepared a LiBr-containing QSCI-B(Ph)4 using the same composition to explore the versatility (Figure 3-14 and Table 3-1). The LiBr-based QSIC-B(Ph)4 showed a deshielded narrow 7Li NMR peak and improved tLi+ and Li+ conductivity, demonstrating that commonly regarded nonideal salts can be ion conductor candidates when coupled with QSIC design.
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Figure 3-13. FTIR and 7Li NMR spectra of QSIC showing (a) a down shifted IR peak of P-F vibration, (b) a deshielded 7Li NMR peak, and (c) more solvated Li+ structure, comparing with liquid electrolyte.
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Figure 3-14. 7Li NMR spectra of QSIC-B(Ph)4 containing 0.1 M LiBr and liquid electrolyte (0.1 M LiBr in EC/PC).
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Table 3-1. Li+ transference number and Li+ conductivity of QSIC-B(Ph)4 containing 0.1 M LiBr (vs.
liquid electrolyte, 0.1 M LiBr in EC/PC).
Electrolyte Li+ transference number Li+ conductivity (S cm-1)
0.1 M LiBr in EC/PC 0.18 1.2 x 10-4
QSIC-B(Ph)4 w/ LiBr 0.72 3.0 x 10-4
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Based on the facile and selective Li+ transport of QSIC described above, we explored its feasibility as a single-ion conductor for potential use in Li batteries. As expected in the remarkable John newman’s works,26,27 application of a single-ion conductor to the battery could improve the electrode kinetics of both cathode and anode as well as exclude the side reactions of anions, which almost previous studies focused on. To investigate cathode kinetics, QSIC-B(Ph)4 is paired with LiNi0.8Co0.1Mn0.1O2 (NCM811) half-cell. As a model study, NCM811 cathode was prepared with thickness of 10 μm, same with that of single NCM811 particle, to facilitate electrochemical activation of NCM811 particles. The QSIC-B(Ph)4 paired half-cell showed a very close open circuit voltage with that of liquid electrolyte-containing half-cell (QSIC-B(Ph)4 = 3.53 V, liquid electrolyte = 3.54 V vs.
Li/Li+), confirming that QSIC-B(Ph)4 could utilize every active particles. The cyclic voltammogram of the cathode half-cell exhibited reversible Li intercalation and deintercalation through the QSIC-B(Ph)4
with smaller potential difference between cathodic and anodic peak potentials than that of control cell (Figure 3-15a). Moreover, the QSIC-B(Ph)4 effectively alleviates the cell polarization upon the repeated current stimuli during charge and discharge reaction in the galvanostatic intermittent titration technique analysis (Figure 3-15b), indicating the improvement in cathode kinetics by its high Li+ conductivity.
This facile redox kinetics of QSIC was further verified by Cu|Li asymmetric cell. Figure 3-15c showed stable Li plating and stripping of QSIC-B(Ph)4 cell with high peak current density, whereas that of control cell exhibited a low peak current density with an intense anion decomposition at 1.9 V vs. Li/Li+ (inset of Figure 3-15c). Some decrease of oxidation onset potential is observed at QSIC-B(Ph)4 than that of liquid electrolyte, but it seems to a reasonable level of electrochemical stability window (>5 V vs. Li/Li+) for operating commercial high voltage cathodes (Figure 3-16). In addition, Tafel plot of control electrolyte cell presents a low exchange current density of 0.010 mA cm-2, indicating its rather sluggish Li extraction kinetics (Figure 3-15d). By comparison, the exchange current density rises to 0.019 mA cm-2 with the QSIC, demonstrating that the accelerated Li+ through QSIC provides adequate amount of Li+ for redox and thus facilitate both cathode and anode kinetics.
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Figure 3-15. Facile electrochemical kinetics of LiGQ. (a) CV and (b) GITT profiles of NCM811 half- cell containing QSIC (vs. liquid electrolyte). (c) CV profiles and (d) its Tafel plots of Cu|Li cell containing QSIC (vs. liquid electrolyte).
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Figure 3-16. LSV profiles of QSIC-B(Ph)4 and liquid electrolyte at a scan rate of 0.1 mV s-1 showing oxidation onset potentials about >5V vs. Li/Li+.
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Another noteworthy advantage of the QSIC is the substantially low anion activity, even in the same condition of using the carbonate-based solvents (vs. liquid electrolyte). Suppressed side reactions of anions and homogeneously fast flow of Li+ are needed to ensure the longer cycle life of Li metal anodes.7,8 Li|Li symmetric cells containing QSIC-B(Ph)4 and control liquid electrolyte were examined by galvanostatic Li plating and stripping cycle test at a current density and capacity of 1 mA cm-2 and 1 mAh cm-2, respectively (Figure 3-17a). The cell with control electrolyte showed irreversible voltage fluctuations with a large overpotential, revealing that the solid-electrolyte interphase (SEI) layers formed by the liquid electrolyte is not sufficiently stable and thus accelerates growth of Li dendrites.28 In sharp contrast, the cell with the QSIC exhibited stable cycle performance over 200 h without an increase of overpotential and modest mitigation in voltage peaking, indicating the reduced concentration polarizations.29 The clear difference on the interfacial resistance was further investigated by electrochemical impedance spectroscopy after the cycle test (Figure 3-17b). The QSIC has the lower interfacial resistance (145 Ω cm2) relative to the liquid electrolyte (362 Ω cm2), indicating that the QSIC enables the formation of a stable SEI layer on Li metal, even though carbonate solvents are utilized.
Moreover, the smooth and dense texture of Li metal after the cycle test verifies that the QSIC suppressed the dendritic Li growth compared with that of the control cell Figure 3-17c).
To elucidate this interfacial behavior, the SEI layers on the Li metal were investigated by X- ray photoelectron microscopy (XPS) analysis. The C1s spectra showed that the SEI layer formed by both QSIC and liquid electrolyte have the typical carbonaceous species (C=C, C-O, C=O, and CO32-) (Figure 3-17d). These organic species are well known to originate from decomposition of carbonate- based solvents.30 However, major differences lie in the significant variation in the F and P contents in these two interphases. QSIC exhibited two smaller peaks tentatively assigned to Li-F (685.1 eV) and LixPFyOz (687.6 eV) in the F1s spectra (Figure 3-17e). It is believed that LiF and LixPFyOz arise from decomposition of LiPF6. Notably, the SEI formed by QSIC had smaller F content than that of control electrolyte in any depth (Figure 3-18), demonstrating the continuous decomposition of PF6 is sufficiently suppressed on the QSIC-Li interface during the repetitive Li plating and stripping. This unique SEI layer evidently enables the interfacial stabilization of Li metal, thereby suppressing the Li dendrite growth and the accompanied electrolyte consumption. This is in agreement with our ion regulation results that QSIC could trap and restrict motion of the anions.
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Figure 3-17. Stable solid-electrolyte interphase on the QSIC-Li metal anode interface. (a) Voltage profiles of Li|Li symmetric cell with QSIC, comparing that of liquid electrolyte. (b) EIS profiles of the Li|Li symmetric cells after the cycle test. (c) Morphology of Li anode surfaces after the cycle test. XPS (d) C1s and (e) F1s spectra of the SEI formed on the QSIC-Li metal anode interface (vs. liquid electrolyte).
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Figure 3-18. XPS depth profiles of the SEI formed on (a) liquid electrolyte-Li metal and (b) QSIC- B(Ph)4-Li metal interfaces.
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Application of the QSIC as a single-ion conducting electrolyte for lithium-ion battery and anode-free lithium storage was explored. As a proof-of-concept, 200 um Li metal anode (half-cell) and Cu foil (anode-free cell) were assembled with QSIC and NCM811 cathode. Figure 3-19a compares the discharge rate capability of the half-cell containing QSIC and control electrolyte, in which the discharge current density was varied from 0.2 to 2.0 C at a fixed charge current density of 0.2 C. The QSIC cell showed higher discharge capacities than control cell over a wide range of discharge current densities.
In high rate cycling test at a current density of 1.0 C, the QSIC cell showed stable cycling performance during 250 cycles with 88.6 % retention, compared to the control cell (59.8 %) (Figure 3-19b). These results demonstrated that the QSIC featuring the high Li+ conductivity facilitated the charge and discharge reaction during cycling, which is highly consistent with the previous model studies on kinetics.
The electrochemical compatibility of QSIC was further explored by anode-free lithium storage. Anode- free full-cell configuration is known to a promising alternative high energy density battery, however, has suffered from unsatisfied capacity retention owing to limited reversibility of anode.7,8 Figure 3-19c shows a substantial improvement in gravimetric capacity of the anode-free cell containing QSIC compared to the control electrolyte. Furthermore, the QSIC anode-free cell showed higher capacity retention with cycling than the control electrolyte cell (Figure 3-19d). These results confirm that the synergistic effect of facile Li+ conduction and interfacial stabilization of QSIC contribute the improvement in reversibility of anode-free lithium storage. Future works will be devoted to exploring the full-cell design with the QSIC to realize a practical level of high energy density battery, along with further efforts to improve Columbic efficiency by fine-tuning the electrophoretic effect.
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Figure 3-19. Application to half cell and anode-free lithium storage. (a) Rate performance and (b) high rate cycle test of the NCM811 half-cell with QSIC and liquid electrolyte. The cycle performance was conducted at a discharge current density of 1.0 C. (c) Voltage profile and (d) cycle performance of the anode-free full-cell configuration with QSIC and liquid electrolyte.
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