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Besides NiFe-based catalysts, the performance of transition metal sulfides can also be enhanced by Fe incorporation, similar to the effect of Fe-ion doping. For instance, Liu's group reported a novel HER candidate, namely, a Fe-doped NiS2 nanosheet, with high activity and long-time durability. The theoretical and experimental results show that Fe3+ doping into the surface lattice of the NiS2 (002) crystal plane can reduce the activation energy of H2 formation. Further, Fe-doped NiS2 nanosheets can be a bifunctional electrocatalyst with high activities toward both OER and HER in the same media. Sun's group reported the exploitation of iron-doped nickel disulfide nanoarrays on Ti substrate via subsequent sulfidation treatment and used it as a bifunctional electrocatalyst with overall water-splitting capacity. In addition, Cao's group prepared a vertically oriented Fe-doped Ni3S2-nanosheet-based bifunctional electrocatalyst toward both HER and OER with outstanding activity and stability in alkaline electrolytes. Based on detailed experiments and theoretical simulations, it was proved that the electrochemically active surface area, water adsorption ability, and hydrogen adsorption energy of Ni3S2 can be regulated via Fe doping (Fig. 5). For transition metal selenides, the effect of Fe incorporation is also evident. Yu and co-workers described ultrathin Fe-doped NiSe2 nanowires (diameter: <1.7 nm) via a soft-template-mediated colloidal synthesis strategy, further revealing the amorphous hydroxide layers formed in situ that contributed toward the enhanced OER activity. Moreover, Zhao et al. incorporated both Fe dopants and Co vacancies into atomically thin CoSe2 nanobelts for oxygen evolution catalysis, and the resultant CoSe2–DFe–VCo exhibits much higher catalytic activity than other defect-activated CoSe2 and previously reported FeCo compounds. Deeper characterizations and theoretical calculations identify the most active centers of Co2 sites that are adjacent to the VCo-nearest Fe site. Fe doping and Co vacancy can synergistically tune the electronic states of Co2 to a near-optimal value, resulting in considerably decreased binding energy of OH* and consequently lowering the catalytic overpotential.

What's the electrolyte?

alkaline

Typical NASICON compounds (such as a NVP cathode and a NTP anode) suffering from severe performance degradation in aqueous electrolytes were employed. The crystal structures, X-ray diffraction (XRD) patterns and SEM images of the prepared NVP/C and NTP/C are shown in Fig. S7.† In Fig. S8,† all of the materials have two featured Raman shifts, in which the G-band (ordered graphitic structure) peak is at 1590 cm−1 and the D-band (disordered portion) is at 1352 cm−1. The thin carbon layer provides good electro-conductivity to NVP/C and NTP/C. In addition, the carbon content of the NVP/C and NTP/C composites measured by elemental analysis (Table S2†) and TG analysis (Fig. S9†) is approximately 3.5 wt%.

What's the cathode?

Typical NASICON compounds (such as a NVP cathode and a NTP anode) suffering from severe performance degradation in aqueous electrolytes were employed. The crystal structures, X-ray diffraction (XRD) patterns and SEM images of the prepared NVP/C and NTP/C are shown in Fig. S7.† In Fig. S8,† all of the materials have two featured Raman shifts, in which the G-band (ordered graphitic structure) peak is at 1590 cm−1 and the D-band (disordered portion) is at 1352 cm−1. The thin carbon layer provides good electro-conductivity to NVP/C and NTP/C. In addition, the carbon content of the NVP/C and NTP/C composites measured by elemental analysis (Table S2†) and TG analysis (Fig. S9†) is approximately 3.5 wt%.

What's the anode?

NTP

Typical NASICON compounds (such as a NVP cathode and a NTP anode) suffering from severe performance degradation in aqueous electrolytes were employed. The crystal structures, X-ray diffraction (XRD) patterns and SEM images of the prepared NVP/C and NTP/C are shown in Fig. S7.† In Fig. S8,† all of the materials have two featured Raman shifts, in which the G-band (ordered graphitic structure) peak is at 1590 cm−1 and the D-band (disordered portion) is at 1352 cm−1. The thin carbon layer provides good electro-conductivity to NVP/C and NTP/C. In addition, the carbon content of the NVP/C and NTP/C composites measured by elemental analysis (Table S2†) and TG analysis (Fig. S9†) is approximately 3.5 wt%.

What's the electrolyte?

The nanomesh after galvanostatic EMD deposition is shown in Fig. 2c–e. The electrodeposition resulted in conformal coating of the nanowires with a 8.5 ± 1 nm thick EMD layer. Despite the limited pore size of the nanomesh, the coating was uniformly distributed throughout the depth of the nanowire network, without closing the pores within and at the top of the nanomesh (Fig. 2d). Such open porosity of the cathode is crucial to ensure its high contact area with the battery electrolyte and facilitate transport of Li+ into and out of the electrode. The excellent conformality of EMD on the nanomesh network can be ascribed to both the relatively slow diffusion of reactants through the growing oxide layer, and the pre-deposited MnO2 seed layer, which increases the homogeneity and coulombic efficiency of EMD electrodeposition. X-ray Photoemission Spectroscopy (XPS) confirmed that the as-deposited EMD consists of hydrated MnO2 (Fig. S2†), in agreement with previous observations on the hydrous nature of pristine EMD deposited from aqueous electrolytes. Note that the total thickness of the EMD-coated 3D electrode is much higher than the maximum attainable thickness of electrodeposited MnO2 on planar electrodes (500 nm on Ni and 1.5 μm on carbon-coated TiN), where the thicker EMD layers easily delaminate.

What's the cathode?

The nanomesh after galvanostatic EMD deposition is shown in Fig. 2c–e. The electrodeposition resulted in conformal coating of the nanowires with a 8.5 ± 1 nm thick EMD layer. Despite the limited pore size of the nanomesh, the coating was uniformly distributed throughout the depth of the nanowire network, without closing the pores within and at the top of the nanomesh (Fig. 2d). Such open porosity of the cathode is crucial to ensure its high contact area with the battery electrolyte and facilitate transport of Li+ into and out of the electrode. The excellent conformality of EMD on the nanomesh network can be ascribed to both the relatively slow diffusion of reactants through the growing oxide layer, and the pre-deposited MnO2 seed layer, which increases the homogeneity and coulombic efficiency of EMD electrodeposition. X-ray Photoemission Spectroscopy (XPS) confirmed that the as-deposited EMD consists of hydrated MnO2 (Fig. S2†), in agreement with previous observations on the hydrous nature of pristine EMD deposited from aqueous electrolytes. Note that the total thickness of the EMD-coated 3D electrode is much higher than the maximum attainable thickness of electrodeposited MnO2 on planar electrodes (500 nm on Ni and 1.5 μm on carbon-coated TiN), where the thicker EMD layers easily delaminate.

What's the electrolyte?

On the basis of the change in the direction of the exfoliating voltage, electrochemical exfoliation of 2D antimony, bismuth and their compounds can be divided into DC voltage exfoliation and square-wave voltage exfoliation. DC voltage exfoliation uses a DC power supply to provide the voltage, and the voltage direction does not change during the exfoliation process. The process of preparing 2D antimony, bismuth and their compounds by the DC voltage exfoliation method is similar to that of cathode exfoliation graphene. The exfoliation process was conducted by using a large block of antimony or bismuth or its compound as the cathode and platinum wire or foil as the anode and in an organic solution. This process applies a voltage to the bulk of antimony, bismuth and their compounds at the cathode, drives the insertion of cations in the electrolyte, and uses jet force that produces hydrogen to promote the peeling of antimony, bismuth and their compounds. In 2017, Li et al. illustrated the possible mechanism of cathodic exfoliation to prepare 2D Sb nanosheets in 0.5 M Na2SO4 solution. Driven by the voltage, the Na+ intercalation layer that accumulates at the negative electrode enters the interlayer of the bulk Sb crystal. Because a large amount of the Na+ intercalation layer enters the interlayer of the Sb crystal, the exfoliation of Sb is eventually promoted by the expansion force (Fig. 4a). Furthermore, the positive and negative voltage and the size of the electrolyte cations affect the efficiency of exfoliation.

What's the cathode?

On the basis of the change in the direction of the exfoliating voltage, electrochemical exfoliation of 2D antimony, bismuth and their compounds can be divided into DC voltage exfoliation and square-wave voltage exfoliation. DC voltage exfoliation uses a DC power supply to provide the voltage, and the voltage direction does not change during the exfoliation process. The process of preparing 2D antimony, bismuth and their compounds by the DC voltage exfoliation method is similar to that of cathode exfoliation graphene. The exfoliation process was conducted by using a large block of antimony or bismuth or its compound as the cathode and platinum wire or foil as the anode and in an organic solution. This process applies a voltage to the bulk of antimony, bismuth and their compounds at the cathode, drives the insertion of cations in the electrolyte, and uses jet force that produces hydrogen to promote the peeling of antimony, bismuth and their compounds. In 2017, Li et al. illustrated the possible mechanism of cathodic exfoliation to prepare 2D Sb nanosheets in 0.5 M Na2SO4 solution. Driven by the voltage, the Na+ intercalation layer that accumulates at the negative electrode enters the interlayer of the bulk Sb crystal. Because a large amount of the Na+ intercalation layer enters the interlayer of the Sb crystal, the exfoliation of Sb is eventually promoted by the expansion force (Fig. 4a). Furthermore, the positive and negative voltage and the size of the electrolyte cations affect the efficiency of exfoliation.

What's the anode?

On the basis of the change in the direction of the exfoliating voltage, electrochemical exfoliation of 2D antimony, bismuth and their compounds can be divided into DC voltage exfoliation and square-wave voltage exfoliation. DC voltage exfoliation uses a DC power supply to provide the voltage, and the voltage direction does not change during the exfoliation process. The process of preparing 2D antimony, bismuth and their compounds by the DC voltage exfoliation method is similar to that of cathode exfoliation graphene. The exfoliation process was conducted by using a large block of antimony or bismuth or its compound as the cathode and platinum wire or foil as the anode and in an organic solution. This process applies a voltage to the bulk of antimony, bismuth and their compounds at the cathode, drives the insertion of cations in the electrolyte, and uses jet force that produces hydrogen to promote the peeling of antimony, bismuth and their compounds. In 2017, Li et al. illustrated the possible mechanism of cathodic exfoliation to prepare 2D Sb nanosheets in 0.5 M Na2SO4 solution. Driven by the voltage, the Na+ intercalation layer that accumulates at the negative electrode enters the interlayer of the bulk Sb crystal. Because a large amount of the Na+ intercalation layer enters the interlayer of the Sb crystal, the exfoliation of Sb is eventually promoted by the expansion force (Fig. 4a). Furthermore, the positive and negative voltage and the size of the electrolyte cations affect the efficiency of exfoliation.

What's the electrolyte?

First, we demonstrate the effect of DBSA as a small molecule electrolyte additive, which enhances the electrochemical doping mechanism and device characteristics of OECTs fabricated from conjugated polymers. We start with P3HT and extend the application to PBTTT and DPPT-TT.

What's the electrolyte?

In summary, a novel NVPF@5% rGO with in situ coated 3D carbon network has been successfully prepared via a two-step solid state CTR method and investigated for use as a cathode material for SIBs. The rGO carbon network architecture in NVPF@5% rGO could effectively construct ionic/electronic pathways and provide sufficient electrode–electrolyte contact area for rapid Na+/e− transport, which significantly improves the electrochemical performance of NaVPO4F. In addition, the stability of NVPF@5% rGO after long-term cycle testing was revealed by the results of the ex situ XRD and SEM analyses, which demonstrated the stable crystal structure of NaVPO4F and the sturdy construction of a robust carbon network. Furthermore, the results of the EIS, CV at various scan rates and GITT tests were implemented to study the electrode kinetic characteristics of NVPF@5% rGO. More importantly, the HC//NVPF@5% rGO full-cell system was fabricated with an HC anode and NVPF@5% rGO cathode, and exhibited superior high-rate capabilities (e.g., 81.8 mA h g−1 at 10C and 61.9 mA h g−1 at 20C) and ultralong cycling performance (82.71% capacity retention after 1500 cycles at 5C). Considering the easy synthesis route and impressive results, it is believed that this strategy can inspire the further development of high-rate and long-term cycle life electrode materials for SIBs.

What's the cathode?

NVPF@5% rGO

In summary, a novel NVPF@5% rGO with in situ coated 3D carbon network has been successfully prepared via a two-step solid state CTR method and investigated for use as a cathode material for SIBs. The rGO carbon network architecture in NVPF@5% rGO could effectively construct ionic/electronic pathways and provide sufficient electrode–electrolyte contact area for rapid Na+/e− transport, which significantly improves the electrochemical performance of NaVPO4F. In addition, the stability of NVPF@5% rGO after long-term cycle testing was revealed by the results of the ex situ XRD and SEM analyses, which demonstrated the stable crystal structure of NaVPO4F and the sturdy construction of a robust carbon network. Furthermore, the results of the EIS, CV at various scan rates and GITT tests were implemented to study the electrode kinetic characteristics of NVPF@5% rGO. More importantly, the HC//NVPF@5% rGO full-cell system was fabricated with an HC anode and NVPF@5% rGO cathode, and exhibited superior high-rate capabilities (e.g., 81.8 mA h g−1 at 10C and 61.9 mA h g−1 at 20C) and ultralong cycling performance (82.71% capacity retention after 1500 cycles at 5C). Considering the easy synthesis route and impressive results, it is believed that this strategy can inspire the further development of high-rate and long-term cycle life electrode materials for SIBs.

What's the anode?

HC

In summary, a novel NVPF@5% rGO with in situ coated 3D carbon network has been successfully prepared via a two-step solid state CTR method and investigated for use as a cathode material for SIBs. The rGO carbon network architecture in NVPF@5% rGO could effectively construct ionic/electronic pathways and provide sufficient electrode–electrolyte contact area for rapid Na+/e− transport, which significantly improves the electrochemical performance of NaVPO4F. In addition, the stability of NVPF@5% rGO after long-term cycle testing was revealed by the results of the ex situ XRD and SEM analyses, which demonstrated the stable crystal structure of NaVPO4F and the sturdy construction of a robust carbon network. Furthermore, the results of the EIS, CV at various scan rates and GITT tests were implemented to study the electrode kinetic characteristics of NVPF@5% rGO. More importantly, the HC//NVPF@5% rGO full-cell system was fabricated with an HC anode and NVPF@5% rGO cathode, and exhibited superior high-rate capabilities (e.g., 81.8 mA h g−1 at 10C and 61.9 mA h g−1 at 20C) and ultralong cycling performance (82.71% capacity retention after 1500 cycles at 5C). Considering the easy synthesis route and impressive results, it is believed that this strategy can inspire the further development of high-rate and long-term cycle life electrode materials for SIBs.

What's the electrolyte?

In summary, a novel NVPF@5% rGO with in situ coated 3D carbon network has been successfully prepared via a two-step solid state CTR method and investigated for use as a cathode material for SIBs. The rGO carbon network architecture in NVPF@5% rGO could effectively construct ionic/electronic pathways and provide sufficient electrode–electrolyte contact area for rapid Na+/e− transport, which significantly improves the electrochemical performance of NaVPO4F. In addition, the stability of NVPF@5% rGO after long-term cycle testing was revealed by the results of the ex situ XRD and SEM analyses, which demonstrated the stable crystal structure of NaVPO4F and the sturdy construction of a robust carbon network. Furthermore, the results of the EIS, CV at various scan rates and GITT tests were implemented to study the electrode kinetic characteristics of NVPF@5% rGO. More importantly, the HC//NVPF@5% rGO full-cell system was fabricated with an HC anode and NVPF@5% rGO cathode, and exhibited superior high-rate capabilities (e.g., 81.8 mA h g−1 at 10C and 61.9 mA h g−1 at 20C) and ultralong cycling performance (82.71% capacity retention after 1500 cycles at 5C). Considering the easy synthesis route and impressive results, it is believed that this strategy can inspire the further development of high-rate and long-term cycle life electrode materials for SIBs.

What's the cathode?

NVPF@5% rGO

Electrochemical measurements. All electrochemical measurements were made using an Autolab potentiostat (PGSTAT) and the custom-made three electrode electrochemical cell. The cells components are as follows: thin film deposited on FTO coated glass (working electrode), platinum dispersed on FTO coated glass (counter electrode), silver (Ag) wire (pseudo reference electrode), and 0.1 M tBuNH PF6 in acetonitrile (inert electrolyte). Cyclic voltammetry was measured by applying an oxidising potential to the working electrode and scanning the potential from −0.5 V to 1.2 V vs. Ag at a scan rate of 20 mV s−1, followed by reversing the direction. The current response waveform was measured simultaneously. Chronoamperometry was used to oxidise the polymer film during in situ spectroscopy measurements. It was performed by applying an excitation square-wave potential to the working electrode in the following configuration: 0 V for 20 seconds (neutral state), an oxidising potential (Vox) for 40 seconds (oxidised state), and 0 V for 20 seconds (back to neutral state). In situ Raman and UV-vis measurements were taken during the application of an oxidising potential when the current response reached steady state (usually after 5–10 seconds).

What's the electrolyte?

0.1 M tBuNH PF6 in acetonitrile

Various approaches have been employed to address these issues through the engineering of ex situ deposited surface coatings, solid electrolyte interphase (SEI) transplantation, electrolyte additives, and 3D host materials. These studies have been largely empirical with limited effort directed towards a fundamental understanding of the interplay between electrolyte composition, SEI formation, and lithium metal morphology. For example, it is important to quantify the amount of lithium lost into the SEI, soluble reduction products, “dead” lithium, and corrosion, and to identify how these loss mechanisms are mitigated by using additives or surface coatings. This knowledge will enable the development of commercially viable solutions. Empirical approaches have had some success in improving coulombic efficiency and cycle life, but in most cases it is unclear which loss mechanism is mitigated and how. We consider understanding these processes mechanistically to be the “Holy Grail” of lithium anode research, knowledge that may be extensible to other battery chemistries.

What's the anode?

Various approaches have been employed to address these issues through the engineering of ex situ deposited surface coatings, solid electrolyte interphase (SEI) transplantation, electrolyte additives, and 3D host materials. These studies have been largely empirical with limited effort directed towards a fundamental understanding of the interplay between electrolyte composition, SEI formation, and lithium metal morphology. For example, it is important to quantify the amount of lithium lost into the SEI, soluble reduction products, “dead” lithium, and corrosion, and to identify how these loss mechanisms are mitigated by using additives or surface coatings. This knowledge will enable the development of commercially viable solutions. Empirical approaches have had some success in improving coulombic efficiency and cycle life, but in most cases it is unclear which loss mechanism is mitigated and how. We consider understanding these processes mechanistically to be the “Holy Grail” of lithium anode research, knowledge that may be extensible to other battery chemistries.

What's the electrolyte?

A few investigations are recently reported based on dual redox-additive electrolytes, responsible for improving the capacitive performance via anodic and cathodic redox-active electrolytes at negative and positive electrode interfaces, respectively. For example, Zhong et al. constructed an AC-based supercapacitor with two separate redox-additive gel polymer electrolytes: PVA/H2SO4/HQ (in the anodic region) and PVA/H2SO4/MB (in the cathodic region). This approach enhanced the specific capacitance from ∼139 F g−1 to 563.7 F g−1 and specific energy from 4.67 W h kg−1 to 18.7 W h kg−1. Frackowiak et al. reported a significantly improved supercapacitor performance based on dual redox additives KI and VOSO4 in an aqueous H2SO4 electrolyte, respectively, in anodic and cathodic regions. Similarly, Chun et al. showed the effect of dual redox additives methyl viologen chloride and KBr added into a supporting electrolyte on the supercapacitor performance. In another report, Fan et al. examined the effect of redox additives KI and VOSO4, added together in a PVA/H2SO4 polymer electrolyte. The capacitor based on this electrolyte and AC electrodes offered a significantly higher capacitance of 1232.4 F g−1 as compared to that of the PVA/H2SO4 based device (156.4 F g−1). The selection of dual redox additives for redox-active electrolytes is based on various factors. Apart from the factors, namely high solubility, reversible electron-transfer kinetics and chemical/electrochemical stability, their different redox potentials are important criteria to choose the dual redox additives. While one additive with a higher potential is used for the positive electrode, the other one functions at the negative electrode. Simultaneous redox processes at two interfaces, during charge and discharge, are responsible for a high value of the overall pseudocapacitance. Based on the same criteria discussed above, the two redox additives DPA and KI are selected in the present study, which have different redox potentials ∼0.76 V and 0.53 V versus the SHE corresponding to their possible redox reaction(s), respectively.

What's the electrolyte?

PVA/H2SO4/HQ (in the anodic region) and PVA/H2SO4/MB (in the cathodic region)

A few investigations are recently reported based on dual redox-additive electrolytes, responsible for improving the capacitive performance via anodic and cathodic redox-active electrolytes at negative and positive electrode interfaces, respectively. For example, Zhong et al. constructed an AC-based supercapacitor with two separate redox-additive gel polymer electrolytes: PVA/H2SO4/HQ (in the anodic region) and PVA/H2SO4/MB (in the cathodic region). This approach enhanced the specific capacitance from ∼139 F g−1 to 563.7 F g−1 and specific energy from 4.67 W h kg−1 to 18.7 W h kg−1. Frackowiak et al. reported a significantly improved supercapacitor performance based on dual redox additives KI and VOSO4 in an aqueous H2SO4 electrolyte, respectively, in anodic and cathodic regions. Similarly, Chun et al. showed the effect of dual redox additives methyl viologen chloride and KBr added into a supporting electrolyte on the supercapacitor performance. In another report, Fan et al. examined the effect of redox additives KI and VOSO4, added together in a PVA/H2SO4 polymer electrolyte. The capacitor based on this electrolyte and AC electrodes offered a significantly higher capacitance of 1232.4 F g−1 as compared to that of the PVA/H2SO4 based device (156.4 F g−1). The selection of dual redox additives for redox-active electrolytes is based on various factors. Apart from the factors, namely high solubility, reversible electron-transfer kinetics and chemical/electrochemical stability, their different redox potentials are important criteria to choose the dual redox additives. While one additive with a higher potential is used for the positive electrode, the other one functions at the negative electrode. Simultaneous redox processes at two interfaces, during charge and discharge, are responsible for a high value of the overall pseudocapacitance. Based on the same criteria discussed above, the two redox additives DPA and KI are selected in the present study, which have different redox potentials ∼0.76 V and 0.53 V versus the SHE corresponding to their possible redox reaction(s), respectively.

What's the electrolyte?

H2SO4

In summary, we have demonstrated a flexible organic–inorganic composite solid electrolyte consisting of polymer PEO, Li salt (LiTFSI), ionic liquid (BMP-TFSI), and the ceramic ion conducting LATP particles. The addition of BMP-TFSI can not only decrease the interface impedance between the polymer matrix and LATP particles, but also enhance the stability of the composite electrolyte and Li metal anode. Benefitting from the synergistic effect of the organic–inorganic complex, the PBL-CSE membrane shows improved electrochemical properties. The ionic conductivity is above 10−4 S cm−1 at 30 °C. The electrochemical stability window is up to 5.0V (vs. Li+/Li) and the Li+ transference number is 0.48. The PBL-CSE exhibits superb interface stability and impressive capacity against a Li electrode, as well as effective suppression of Li dendrite growth. The assembled solid-state LiFePO4/Li batteries demonstrated excellent cycling stability with a high specific capacity and very large capacity retention at 0.5C at 45 °C, as well as 0.3C at 30 °C. Furthermore, LiFePO4/PBL-CSE/Li all-solid-state pouch cells also demonstrate high flexibility and safety. Thus, this work represents a promising composite electrolyte for high performance, safe and high energy density all-solid-state lithium metal batteries.

What's the anode?

Li metal

In summary, we have demonstrated a flexible organic–inorganic composite solid electrolyte consisting of polymer PEO, Li salt (LiTFSI), ionic liquid (BMP-TFSI), and the ceramic ion conducting LATP particles. The addition of BMP-TFSI can not only decrease the interface impedance between the polymer matrix and LATP particles, but also enhance the stability of the composite electrolyte and Li metal anode. Benefitting from the synergistic effect of the organic–inorganic complex, the PBL-CSE membrane shows improved electrochemical properties. The ionic conductivity is above 10−4 S cm−1 at 30 °C. The electrochemical stability window is up to 5.0V (vs. Li+/Li) and the Li+ transference number is 0.48. The PBL-CSE exhibits superb interface stability and impressive capacity against a Li electrode, as well as effective suppression of Li dendrite growth. The assembled solid-state LiFePO4/Li batteries demonstrated excellent cycling stability with a high specific capacity and very large capacity retention at 0.5C at 45 °C, as well as 0.3C at 30 °C. Furthermore, LiFePO4/PBL-CSE/Li all-solid-state pouch cells also demonstrate high flexibility and safety. Thus, this work represents a promising composite electrolyte for high performance, safe and high energy density all-solid-state lithium metal batteries.

What's the electrolyte?

The charging-discharging and cycling curves were obtained from the hybrid electrolyte cells, as displayed in Fig. 6c and d. The cells were charged/discharged with a constant current density of 0.3 mA cm−2 for the first cycle, 0.5 mA cm−2 for the second to sixth cycles, 1 mA cm−2 for the seventh to eleventh cycles and 1.4 mA cm−2 for longer cycles. The charging/discharging voltage profiles in Fig. 6c and d show that the cell with the laser-treated LLZTO sample has a relatively small overpotential even at a high current density of 1.4 mA cm−2. This result is consistent with the impedance result in Fig. 6a and b, as evident from the reduction of Nyquist curves. For the polished (pristine) LLZTO cell, the cell voltage was found to be unstable in the charging curve of the 8th cycle, implying the formation of an internal short-circuit caused by the propagation of metallic Li. In contrast, notably, the cell with the laser-treated LLZTO sample showed a stable cycling performance without any short-circuit signals (voltage noise and/or sudden drop) at a constant current density of 1.4 mA cm−2. The improved cycling stability of the laser-treated cell was confirmed from the cycling curves shown in Fig. 6e for long-term operation under 1.4 mA cm−2. The cell showed an excellent coulombic efficiency of 88.9% for 1st cycle, 100% for 3rd cycle, 100% for 7th cycle, 99.96% for 11th cycle, 99.92% for 50th cycle, 99.92% for 100th cycle, and 99.92% for 160th cycle compared with the conventional LIBs and a remarkable capacity retention of 96.7% over the 160th cycle. The mechanism to enhance the electrochemical performance in the laser-treated cell might be complicated. The surface change of LLZTO after the laser treatment includes the formation of Li2O2 and amorphous garnet, the possible reduction of LiOH and Li2CO3, and the morphological deformation. Based on our electronic structure analysis of the laser-treated LLZTO sample and its reduced electronic conductivity (Fig. S5c†), it is highly possible that the main origin of the remarkable improvement in the cycling stability of the laser-treated cell is the formation of the stable amorphous layer with a wide bandgap, which can significantly suppress the Li dendrite formation.

What's the cathode?

The charging-discharging and cycling curves were obtained from the hybrid electrolyte cells, as displayed in Fig. 6c and d. The cells were charged/discharged with a constant current density of 0.3 mA cm−2 for the first cycle, 0.5 mA cm−2 for the second to sixth cycles, 1 mA cm−2 for the seventh to eleventh cycles and 1.4 mA cm−2 for longer cycles. The charging/discharging voltage profiles in Fig. 6c and d show that the cell with the laser-treated LLZTO sample has a relatively small overpotential even at a high current density of 1.4 mA cm−2. This result is consistent with the impedance result in Fig. 6a and b, as evident from the reduction of Nyquist curves. For the polished (pristine) LLZTO cell, the cell voltage was found to be unstable in the charging curve of the 8th cycle, implying the formation of an internal short-circuit caused by the propagation of metallic Li. In contrast, notably, the cell with the laser-treated LLZTO sample showed a stable cycling performance without any short-circuit signals (voltage noise and/or sudden drop) at a constant current density of 1.4 mA cm−2. The improved cycling stability of the laser-treated cell was confirmed from the cycling curves shown in Fig. 6e for long-term operation under 1.4 mA cm−2. The cell showed an excellent coulombic efficiency of 88.9% for 1st cycle, 100% for 3rd cycle, 100% for 7th cycle, 99.96% for 11th cycle, 99.92% for 50th cycle, 99.92% for 100th cycle, and 99.92% for 160th cycle compared with the conventional LIBs and a remarkable capacity retention of 96.7% over the 160th cycle. The mechanism to enhance the electrochemical performance in the laser-treated cell might be complicated. The surface change of LLZTO after the laser treatment includes the formation of Li2O2 and amorphous garnet, the possible reduction of LiOH and Li2CO3, and the morphological deformation. Based on our electronic structure analysis of the laser-treated LLZTO sample and its reduced electronic conductivity (Fig. S5c†), it is highly possible that the main origin of the remarkable improvement in the cycling stability of the laser-treated cell is the formation of the stable amorphous layer with a wide bandgap, which can significantly suppress the Li dendrite formation.

What's the anode?

The charging-discharging and cycling curves were obtained from the hybrid electrolyte cells, as displayed in Fig. 6c and d. The cells were charged/discharged with a constant current density of 0.3 mA cm−2 for the first cycle, 0.5 mA cm−2 for the second to sixth cycles, 1 mA cm−2 for the seventh to eleventh cycles and 1.4 mA cm−2 for longer cycles. The charging/discharging voltage profiles in Fig. 6c and d show that the cell with the laser-treated LLZTO sample has a relatively small overpotential even at a high current density of 1.4 mA cm−2. This result is consistent with the impedance result in Fig. 6a and b, as evident from the reduction of Nyquist curves. For the polished (pristine) LLZTO cell, the cell voltage was found to be unstable in the charging curve of the 8th cycle, implying the formation of an internal short-circuit caused by the propagation of metallic Li. In contrast, notably, the cell with the laser-treated LLZTO sample showed a stable cycling performance without any short-circuit signals (voltage noise and/or sudden drop) at a constant current density of 1.4 mA cm−2. The improved cycling stability of the laser-treated cell was confirmed from the cycling curves shown in Fig. 6e for long-term operation under 1.4 mA cm−2. The cell showed an excellent coulombic efficiency of 88.9% for 1st cycle, 100% for 3rd cycle, 100% for 7th cycle, 99.96% for 11th cycle, 99.92% for 50th cycle, 99.92% for 100th cycle, and 99.92% for 160th cycle compared with the conventional LIBs and a remarkable capacity retention of 96.7% over the 160th cycle. The mechanism to enhance the electrochemical performance in the laser-treated cell might be complicated. The surface change of LLZTO after the laser treatment includes the formation of Li2O2 and amorphous garnet, the possible reduction of LiOH and Li2CO3, and the morphological deformation. Based on our electronic structure analysis of the laser-treated LLZTO sample and its reduced electronic conductivity (Fig. S5c†), it is highly possible that the main origin of the remarkable improvement in the cycling stability of the laser-treated cell is the formation of the stable amorphous layer with a wide bandgap, which can significantly suppress the Li dendrite formation.

What's the electrolyte?

In general, the solid electrolytes used to develop Na-ion solid state battery systems are based on inorganic solid materials; in particular, oxide-based materials that have high ionic conductivity and are electrochemically and thermally stable are being studied. Among oxide-based materials, Na-β-alumina materials have been mainly studied. Because of their high ionic conductivity, they have been used in all-solid-state batteries until recently. However, they have the disadvantage of being vulnerable to moisture. Another stable oxide-based material is Na3Zr2Si2PO12 (Na super-ionic conductor, NASICON). NASICON solid electrolytes are promising oxide-based Na-ion conducting materials, with a high ionic conductivity of over 10−4 S cm−1 at room temperature and stability to air and moisture. NASICON ceramics are electrochemically stable up to 7 V, making them suitable for use in high voltage batteries. Owing to its stability advantages, efforts have been made to use NASICON in solid-state NIB systems. However, there remains a critical problem in the resistance between solid particle interfaces caused by the fragile and rigid nature of the oxide material itself, regardless of how high the pressure of the solid electrolyte powder is, making it difficult to utilize it in a solid-state battery system.

What's the electrolyte?

NASICON

3.2.1 Cyclic voltammetry. CV studies have been performed first to optimize the operating potential difference range of the supercapacitor cells. For this purpose, the CV profiles of a typical cell (Cell#4) have been recorded for varying potential difference ranges as shown in Fig. 4a. A close inspection indicates that CV responses show almost rectangular-box like patterns without much deviation, indicating stable capacitive behavior, for the range from 0 to 2.5 V. Thus, all the cells have been electrochemically characterized up to 2.5 V in two electrode configuration, in the present study. This has been confirmed from GCD studies also, as discussed later. It may be noted that the stability window of the dual redox-active electrolyte film is high (∼6.2 V), as discussed above, and the charging voltage of the device is restricted to ∼2.5 V only. The most possible reason for the limited charging voltage of the device is related to the possible reaction(s) of the electrolyte ions with the surface functional groups (carbonyl, carboxyl, hydroxyl, etc.) attached to the activated carbon electrode surface. The electrochemical performance of all the supercapacitor cells (Cell#1 to Cell#4) has been compared, comparing their CV patterns, recorded at a scan rate of 10 mV s−1, as shown in Fig. 4b. Cell#1 (with the GPE without redox additives) shows almost a rectangular pattern, similar to the characteristics of capacitors with carbon electrodes, indicating the dominance of double-layer-type capacitive nature. The CV profiles of Cell#2 and Cell#3 (containing GPEs with single redox additives DPA and KI, respectively) illustrate distinct reversible redox peaks (Fig. 4b). The reversible redox peaks in Cell#2 are associated with the reversible conversion between diphenyl benzidine, DPB (colorless) and DPB (violet), which is initiated after irreversible redox transformation of DPA into DPB (colorless), as shown in the following Scheme (1).

What's the electrolyte?

Electrochromic devices (ECDs) have been widely investigated for application in next-generation displays and smart windows thanks to their highly efficient optical transmittance modulation properties. However, several challenges such as chemical and environmental instabilities and leakage of electrolytes limit their practical applications. In this paper, we report a simple and efficient approach for synthesising ultraviolet (UV)-cured poly(methyl methacrylate) (PMMA) gels that can be used as safe electrolytes. The ECDs fabricated with the 10 min UV-cured PMMA gel electrolyte deliver remarkable device performances with a wide optical transmittance transition (ΔT) of 51.3% at a wavelength of 550 nm under −1.2–0 V bias range and fast switching times (Δt) of 1.5 s and 2.0 s for bleaching and colouration, respectively. In addition, excellent operational stability of 98.9% after 11500 cycles and environmental stability at a wide temperature range of −20 to +70 °C are exhibited. Moreover, a smart electrochromic window system, including an ECD connected with an Arduino circuit, is developed. These smart windows can change colour by simultaneously monitoring the illumination and UV intensities of sunlight.

What's the electrolyte?

ultraviolet (UV)-cured poly(methyl methacrylate) (PMMA)

Electrochromic devices (ECDs) have been widely investigated for application in next-generation displays and smart windows thanks to their highly efficient optical transmittance modulation properties. However, several challenges such as chemical and environmental instabilities and leakage of electrolytes limit their practical applications. In this paper, we report a simple and efficient approach for synthesising ultraviolet (UV)-cured poly(methyl methacrylate) (PMMA) gels that can be used as safe electrolytes. The ECDs fabricated with the 10 min UV-cured PMMA gel electrolyte deliver remarkable device performances with a wide optical transmittance transition (ΔT) of 51.3% at a wavelength of 550 nm under −1.2–0 V bias range and fast switching times (Δt) of 1.5 s and 2.0 s for bleaching and colouration, respectively. In addition, excellent operational stability of 98.9% after 11500 cycles and environmental stability at a wide temperature range of −20 to +70 °C are exhibited. Moreover, a smart electrochromic window system, including an ECD connected with an Arduino circuit, is developed. These smart windows can change colour by simultaneously monitoring the illumination and UV intensities of sunlight.

What's the electrolyte?

UV-cured PMMA gel

Li metal is regarded as the best candidate for anode materials because of its high theoretical capacity and negative electrode potential. However, due to the continuous parasitic side reactions and messy growing Li dendrites, the practical use of the Li metal as an anode is seriously limited. An in situ formed artificial solid electrolyte interface (SEI) can commendably solve the above-mentioned problems; however, inevitable cracks and fractures are found during their long-term service due to the existence of inorganic compounds (or alloys) in the artificial-SEI. Herein, a self-repairing alloy for protecting the Li metal anode was prepared via a facile in situ reaction. A long-term cycling life of more than 1800 h and 1400 h was obtained for the self-repairing alloy protected Li metal anode at a practical current density of 2 mA cm−2 and 5 mA cm−2. Even after increasing the deposition capacity to 15 mA h cm−2, no dendrite formation was detected in the self-repairing alloy protective Li anode. This is expected to be a promising strategy in achieving a stable Li metal anode/electrolyte interface and providing a new way of thinking for the development of Li metal batteries.

What's the anode?

Li metal

Li metal is regarded as the best candidate for anode materials because of its high theoretical capacity and negative electrode potential. However, due to the continuous parasitic side reactions and messy growing Li dendrites, the practical use of the Li metal as an anode is seriously limited. An in situ formed artificial solid electrolyte interface (SEI) can commendably solve the above-mentioned problems; however, inevitable cracks and fractures are found during their long-term service due to the existence of inorganic compounds (or alloys) in the artificial-SEI. Herein, a self-repairing alloy for protecting the Li metal anode was prepared via a facile in situ reaction. A long-term cycling life of more than 1800 h and 1400 h was obtained for the self-repairing alloy protected Li metal anode at a practical current density of 2 mA cm−2 and 5 mA cm−2. Even after increasing the deposition capacity to 15 mA h cm−2, no dendrite formation was detected in the self-repairing alloy protective Li anode. This is expected to be a promising strategy in achieving a stable Li metal anode/electrolyte interface and providing a new way of thinking for the development of Li metal batteries.

What's the electrolyte?

Li metal is regarded as the best candidate for anode materials because of its high theoretical capacity and negative electrode potential. However, due to the continuous parasitic side reactions and messy growing Li dendrites, the practical use of the Li metal as an anode is seriously limited. An in situ formed artificial solid electrolyte interface (SEI) can commendably solve the above-mentioned problems; however, inevitable cracks and fractures are found during their long-term service due to the existence of inorganic compounds (or alloys) in the artificial-SEI. Herein, a self-repairing alloy for protecting the Li metal anode was prepared via a facile in situ reaction. A long-term cycling life of more than 1800 h and 1400 h was obtained for the self-repairing alloy protected Li metal anode at a practical current density of 2 mA cm−2 and 5 mA cm−2. Even after increasing the deposition capacity to 15 mA h cm−2, no dendrite formation was detected in the self-repairing alloy protective Li anode. This is expected to be a promising strategy in achieving a stable Li metal anode/electrolyte interface and providing a new way of thinking for the development of Li metal batteries.

What's the anode?

Li metal

As indicated by the above results, the Li/Al-ion electrolyte exhibits better cycle stability than the Al-ion electrolyte. To further investigate the electrolytes, the voltammograms of these electrolytes measured at different scan rates were analysed using the Randles–Sevick equation (Fig. S7†). The calculated Li-ion diffusion coefficients are 2.84 × 10−12 cm2 s−1 and 7.96 × 10−13 cm2 s−1 in the WO3 and the Ti-V2O5 films, and the Al-ion diffusion coefficients are 2.72 × 10−14 cm2 s−1 and 7.99 × 10−15 cm2 s−1 in the WO3 and the Ti-V2O5 films, respectively. The obtained diffusion coefficients are all two magnitudes lower for Al3+ than Li+, owing to the strong electrostatic interaction between trivalent Al3+ and the host lattice, which slows the ion diffusion in the films. For the mixed Li/Al-ion electrolyte, the obtained diffusion curves are between those of the pure Li-ion electrolyte and the pure Al-ion electrolyte (Fig. S7d, h†). Such a performance is much better compared with the coefficient of pure Al-ions in the films, indicating better ion diffusion in the lattices and promoting the cycle stabilities. Previous reports also indicated that the mixed Li/Al-ion electrolyte can promote the cycling performance in Al-anode batteries. The utilization of mixed ion electrolytes can be a promising direction and deserves further investigation.

What's the anode?

Al

As indicated by the above results, the Li/Al-ion electrolyte exhibits better cycle stability than the Al-ion electrolyte. To further investigate the electrolytes, the voltammograms of these electrolytes measured at different scan rates were analysed using the Randles–Sevick equation (Fig. S7†). The calculated Li-ion diffusion coefficients are 2.84 × 10−12 cm2 s−1 and 7.96 × 10−13 cm2 s−1 in the WO3 and the Ti-V2O5 films, and the Al-ion diffusion coefficients are 2.72 × 10−14 cm2 s−1 and 7.99 × 10−15 cm2 s−1 in the WO3 and the Ti-V2O5 films, respectively. The obtained diffusion coefficients are all two magnitudes lower for Al3+ than Li+, owing to the strong electrostatic interaction between trivalent Al3+ and the host lattice, which slows the ion diffusion in the films. For the mixed Li/Al-ion electrolyte, the obtained diffusion curves are between those of the pure Li-ion electrolyte and the pure Al-ion electrolyte (Fig. S7d, h†). Such a performance is much better compared with the coefficient of pure Al-ions in the films, indicating better ion diffusion in the lattices and promoting the cycle stabilities. Previous reports also indicated that the mixed Li/Al-ion electrolyte can promote the cycling performance in Al-anode batteries. The utilization of mixed ion electrolytes can be a promising direction and deserves further investigation.

What's the electrolyte?

Li/Al-ion

As indicated by the above results, the Li/Al-ion electrolyte exhibits better cycle stability than the Al-ion electrolyte. To further investigate the electrolytes, the voltammograms of these electrolytes measured at different scan rates were analysed using the Randles–Sevick equation (Fig. S7†). The calculated Li-ion diffusion coefficients are 2.84 × 10−12 cm2 s−1 and 7.96 × 10−13 cm2 s−1 in the WO3 and the Ti-V2O5 films, and the Al-ion diffusion coefficients are 2.72 × 10−14 cm2 s−1 and 7.99 × 10−15 cm2 s−1 in the WO3 and the Ti-V2O5 films, respectively. The obtained diffusion coefficients are all two magnitudes lower for Al3+ than Li+, owing to the strong electrostatic interaction between trivalent Al3+ and the host lattice, which slows the ion diffusion in the films. For the mixed Li/Al-ion electrolyte, the obtained diffusion curves are between those of the pure Li-ion electrolyte and the pure Al-ion electrolyte (Fig. S7d, h†). Such a performance is much better compared with the coefficient of pure Al-ions in the films, indicating better ion diffusion in the lattices and promoting the cycle stabilities. Previous reports also indicated that the mixed Li/Al-ion electrolyte can promote the cycling performance in Al-anode batteries. The utilization of mixed ion electrolytes can be a promising direction and deserves further investigation.

What's the electrolyte?

Al-ion electrolyte.

In addition, the long-term cycling test at 5C was performed to investigate the cycling performance of all the NVPF-based materials shown in Fig. 3d. The NVPF@5% rGO gave the highest initial discharge capacity of 83.2 mA h g−1, compared with the NVPF (74.6 mA h g−1), NVPF@2.5% rGO (76.4 mA h g−1), and NVPF@7.5% rGO (79.9 mA h g−1). It also exhibited an excellent capacity retention of 94.06% after 1000 cycles, achieving a high average coulombic efficiency (CE) of 99.6% throughout. Therefore, apart from its outstanding rate capability, the NVPF@5% rGO also exhibited an extremely durable cycling performance. Furthermore, due to the high conductivity and uniformity of the rGO carbon network, the ΔV values of NVPF@5% rGO were smaller than that of others in each cycle (inset Fig. 3d). The NVPF@7.5% rGO had a slight increase in polarization compared to the NVPF@5% rGO at 1000 cycles, which could be caused by the uneven carbon network. But the polarization of NVPF increased almost linearly during cycling, which leads to a low CE and poor cycling performance. Moreover, although NVPF@2.5% rGO with an incomplete carbon network can inhibit the increase of polarization before 400 cycles, it still exhibits high polarization after more cycles. The schematic diagram of the Na+ and e− transport pathways in NVPF@5% rGO is shown in Fig. 3e. As shown, because of the robust 3D continuous conductive network structure, the NVPF@5% rGO could not only achieve a highly efficient, ultrafast and continuous transmission of Na+ and e−, but could also build an effective buffer to accommodate volume changes during charge (desodiation)/discharge (sodiation) processes. In addition, the large specific surface area of the rGO carbon network in NVPF@5% rGO could serve as electrolyte storage micro-pools and transport channels, which could provide sufficient and timely electrolyte replenishment for ultrafast charging/discharging.

What's the electrolyte?

NVPF@5% rGO

Dr Hongfei Li obtained his B.S. degree and M.S. degree in materials science and engineering from Central South University and Tsinghua University, respectively. After that, he received his PhD degree from City University of Hong Kong. Now, he is an associate professor at Songshan Lake Materials Laboratory. His research focuses on aqueous batteries, flexible and wearable energy storage devices, and polymer electrolytes. He has published more than 68 scientific papers with over 3800 total citations and an h-index of 36.

What's the electrolyte?

polymer

We are able to clarify that the electrochemical doping of the polymer is still predominantly caused by the formation of ion pairs between charged polymer and PF6− anions, as DBSA itself causes irreversible doping of the polymer (Fig. 4). Fig. 4a shows the behaviour of DBSA in an electrochemical cell with no other salts present by exposing a P3HT thin film to 0.1 M DBSA in ACN electrolyte and measuring in situ ERRS. We observe the typical hole polaron signature in P3HT even at 0 V, indicating the ability of DBSA to chemically dope the polymer film when no other salts are present. We monitor the CC peak position as a function of oxidising potential and find the doping level has a minor dependence on applied bias indicated by the CC peak downshifting from 1425 cm−1 to 1417 cm−1 (Fig. 4b); the electric field causes more DBSA anions (dissociated from protons) to drift and penetrate into the P3HT film causing further doping. After discharge (0 V) the film is irreversibly doped and the CC peak position remains at 1417 cm−1 (green triangle Fig. 4a). Further evidence for irreversible p-type doping caused by DBSA is provided by changes in the UV-Vis absorbance spectra, HOMO level, and work function (see Fig. S2, ESI†). Thus, we elucidate that the mixed ionic electrolyte (with DBSA and TBA:PF6) causes reversible electrochemical doping of the polymer caused by the formation of ion pairs between charged polymer and PF6− anions. This suggests that the effect of DBSA in the mixed electrolyte is to modify the interaction between the polymer and the electrolyte allowing electrochemical doping at a lower voltage. DBSA has previously been used as a surfactant added to PEDOT:PSS formulations to improve wettability. We believe that DBSA could be having a similar effect on P3HT, interacting with the polymer and acting as a surfactant to enhance the strength of the interactions between the polymer and the electrolyte. In other words, DBSA may modulate the RED between the polymer and the solvent. We demonstrated the effect of RED between the solvent and the polymer on oxidation onset during cyclic voltammetry measurements: there is a lower RED between DCM and P3HT compared to ACN and P3HT, which results in a reduction in the oxidation onset of 0.28 V (Fig. 2). Thus, the RED between the polymer and the electrolyte solvent appears to influence the injection barrier into the film. By adding DBSA to the electrolyte, it may reduce the RED, thus causing a higher mass in the polymer film (solvent and ions), ultimately decreasing the injection barrier into the film.

What's the electrolyte?

0.1 M DBSA in ACN

We are able to clarify that the electrochemical doping of the polymer is still predominantly caused by the formation of ion pairs between charged polymer and PF6− anions, as DBSA itself causes irreversible doping of the polymer (Fig. 4). Fig. 4a shows the behaviour of DBSA in an electrochemical cell with no other salts present by exposing a P3HT thin film to 0.1 M DBSA in ACN electrolyte and measuring in situ ERRS. We observe the typical hole polaron signature in P3HT even at 0 V, indicating the ability of DBSA to chemically dope the polymer film when no other salts are present. We monitor the CC peak position as a function of oxidising potential and find the doping level has a minor dependence on applied bias indicated by the CC peak downshifting from 1425 cm−1 to 1417 cm−1 (Fig. 4b); the electric field causes more DBSA anions (dissociated from protons) to drift and penetrate into the P3HT film causing further doping. After discharge (0 V) the film is irreversibly doped and the CC peak position remains at 1417 cm−1 (green triangle Fig. 4a). Further evidence for irreversible p-type doping caused by DBSA is provided by changes in the UV-Vis absorbance spectra, HOMO level, and work function (see Fig. S2, ESI†). Thus, we elucidate that the mixed ionic electrolyte (with DBSA and TBA:PF6) causes reversible electrochemical doping of the polymer caused by the formation of ion pairs between charged polymer and PF6− anions. This suggests that the effect of DBSA in the mixed electrolyte is to modify the interaction between the polymer and the electrolyte allowing electrochemical doping at a lower voltage. DBSA has previously been used as a surfactant added to PEDOT:PSS formulations to improve wettability. We believe that DBSA could be having a similar effect on P3HT, interacting with the polymer and acting as a surfactant to enhance the strength of the interactions between the polymer and the electrolyte. In other words, DBSA may modulate the RED between the polymer and the solvent. We demonstrated the effect of RED between the solvent and the polymer on oxidation onset during cyclic voltammetry measurements: there is a lower RED between DCM and P3HT compared to ACN and P3HT, which results in a reduction in the oxidation onset of 0.28 V (Fig. 2). Thus, the RED between the polymer and the electrolyte solvent appears to influence the injection barrier into the film. By adding DBSA to the electrolyte, it may reduce the RED, thus causing a higher mass in the polymer film (solvent and ions), ultimately decreasing the injection barrier into the film.

What's the electrolyte?

(with DBSA and TBA:PF6)

The cycle stability in terms of the charge capacity for the two kinds of films in the Li/Al-ion electrolyte was also evaluated as shown in Fig. 3a and b (details for the in situ measurement of both the charge capacity and the transmittance are provided in the ESI†). After 1000 cycles, the capacity retention was about 78% for the Ti-V2O5 film and 72% for the WO3 film. The cycle stability in terms of the electrochromic performance is obtained by monitoring the optical contrast between the colored and the bleached state of the films. As shown in Fig. 3c and d, after 1000 cycles, the Ti-V2O5 film still delivered a transmittance contrast of up to 24.8% at 400 nm, and for the WO3 film, a contrast up to 47.4% was delivered at 780 nm. These values show that 69.5% of the optical modulation can be retained for Ti-V2O5 and 65.4% can be retained for the WO3 film after 1000 cycles. As a comparison, the Al-ion electrolyte, which is commonly used in Al-based EES devices, was applied to replace the Li/Al-ion electrolyte. As shown in Fig. S6,† for both the Ti-V2O5 and the WO3 films, the optical contrast in the transmittance spectra drops rapidly within 100 cycles, indicating that the mixed ion electrolyte effectively promotes the cycle stability of the electrochromic films.

What's the electrolyte?

Li/Al-ion

The cycle stability in terms of the charge capacity for the two kinds of films in the Li/Al-ion electrolyte was also evaluated as shown in Fig. 3a and b (details for the in situ measurement of both the charge capacity and the transmittance are provided in the ESI†). After 1000 cycles, the capacity retention was about 78% for the Ti-V2O5 film and 72% for the WO3 film. The cycle stability in terms of the electrochromic performance is obtained by monitoring the optical contrast between the colored and the bleached state of the films. As shown in Fig. 3c and d, after 1000 cycles, the Ti-V2O5 film still delivered a transmittance contrast of up to 24.8% at 400 nm, and for the WO3 film, a contrast up to 47.4% was delivered at 780 nm. These values show that 69.5% of the optical modulation can be retained for Ti-V2O5 and 65.4% can be retained for the WO3 film after 1000 cycles. As a comparison, the Al-ion electrolyte, which is commonly used in Al-based EES devices, was applied to replace the Li/Al-ion electrolyte. As shown in Fig. S6,† for both the Ti-V2O5 and the WO3 films, the optical contrast in the transmittance spectra drops rapidly within 100 cycles, indicating that the mixed ion electrolyte effectively promotes the cycle stability of the electrochromic films.

What's the electrolyte?

Al-ion

The wettability of electrodes is well known as an important characteristic governing the interaction between the electrolyte and electrode as well as the charge transfer at the interface. As shown in Fig. S6,† the water contact angle (WCA) of COG is measured to be 0°, and the total spreading and penetration time (TSPT) of a water drop is about 3.17 s, indicating the inherent hydrophilicity of the COG. This should be attributed to the presence of some heteroatoms in natural reed straw, leading to heteroatom doping during the carbonization process and hence improved hydrophilicity of carbon materials. In contrast, the TSPT of COG@Zn, porous COG@MnO2 and bulk COG@MnO2 is reduced to 0.92 s, 0.42 s and 0.54 s, respectively. The results indicate that the surface wettability of the electrodes has been enhanced from hydrophilic to super-hydrophilic with the incorporation of MnO2 and Zn nanosheets. As is well known, MnO2 and Zn are inherently hydrophilic. In this case, the nanosheet structure increases the roughness of the material's surface, thereby further improving the hydrophilicity of the material. This favors the accessibility of the electrolyte at the surface of the electrodes and the charge transfer between them, giving rise to enhanced electrochemical performance.

What's the electrolyte?

The Li-metal anode is considered an essential component for obtaining the expected Li–S battery performance. This is because of Li's extremely low negative electrochemical potential, very high theoretical specific capacity, and low density. However, the Li-metal anode has additional issues. Safety concerns arise because of the possibility of cell damage or manufacturing failures due to Li-metal's high reactivity, as well as due to the dendrite formation throughout cycling processes which can cause short circuits if dendrites reach the cathode. The electrolyte (salt and solvent) plays a vital role for rapid lithium transport and providing stability even near the highly reactive lithium surface. At the Li surface, electrolyte decomposition due to reduction reactions triggered by their low electrochemical stability leads to the growth of a multicomponent passivating film. This film, known as the solid electrolyte interphase (SEI) may have beneficial passivation properties, or, quite the opposite, may be a cause of irreversible capacity loss. The specific film behavior depends on its chemical composition, structure, and thickness. Moreover, it has been found the SEI properties depend largely on the solvent, salt, and additives used as well as on the electrode structure. Therefore, understanding the role of the electrolyte in SEI formation and its properties should be a useful strategy for controlling dendrite formation, reducing the effects of anode degradation and stabilizing the surface, thus, potentially improving the performance, cycling and safety of the batteries.

What's the anode?

Li-metal

The Li-metal anode is considered an essential component for obtaining the expected Li–S battery performance. This is because of Li's extremely low negative electrochemical potential, very high theoretical specific capacity, and low density. However, the Li-metal anode has additional issues. Safety concerns arise because of the possibility of cell damage or manufacturing failures due to Li-metal's high reactivity, as well as due to the dendrite formation throughout cycling processes which can cause short circuits if dendrites reach the cathode. The electrolyte (salt and solvent) plays a vital role for rapid lithium transport and providing stability even near the highly reactive lithium surface. At the Li surface, electrolyte decomposition due to reduction reactions triggered by their low electrochemical stability leads to the growth of a multicomponent passivating film. This film, known as the solid electrolyte interphase (SEI) may have beneficial passivation properties, or, quite the opposite, may be a cause of irreversible capacity loss. The specific film behavior depends on its chemical composition, structure, and thickness. Moreover, it has been found the SEI properties depend largely on the solvent, salt, and additives used as well as on the electrode structure. Therefore, understanding the role of the electrolyte in SEI formation and its properties should be a useful strategy for controlling dendrite formation, reducing the effects of anode degradation and stabilizing the surface, thus, potentially improving the performance, cycling and safety of the batteries.

What's the electrolyte?

The Li-metal anode is considered an essential component for obtaining the expected Li–S battery performance. This is because of Li's extremely low negative electrochemical potential, very high theoretical specific capacity, and low density. However, the Li-metal anode has additional issues. Safety concerns arise because of the possibility of cell damage or manufacturing failures due to Li-metal's high reactivity, as well as due to the dendrite formation throughout cycling processes which can cause short circuits if dendrites reach the cathode. The electrolyte (salt and solvent) plays a vital role for rapid lithium transport and providing stability even near the highly reactive lithium surface. At the Li surface, electrolyte decomposition due to reduction reactions triggered by their low electrochemical stability leads to the growth of a multicomponent passivating film. This film, known as the solid electrolyte interphase (SEI) may have beneficial passivation properties, or, quite the opposite, may be a cause of irreversible capacity loss. The specific film behavior depends on its chemical composition, structure, and thickness. Moreover, it has been found the SEI properties depend largely on the solvent, salt, and additives used as well as on the electrode structure. Therefore, understanding the role of the electrolyte in SEI formation and its properties should be a useful strategy for controlling dendrite formation, reducing the effects of anode degradation and stabilizing the surface, thus, potentially improving the performance, cycling and safety of the batteries.

What's the anode?

Li-metal

In order to achieve higher power/energy densities for electrochemical energy storage, dual-ion batteries (DIBs), relying on the migration of both cations and anions to store charge concurrently, have been developed. Compared with lithium ion batteries, dual-ion migration behavior can shorten the charge carrier distance. From a cation perspective, lithium ions, as a charge carrier, have contributed to energy storage for several decades. However, many issues have also been raised, such as safety, excessive resource consumption, and environmental pollution. In the process of searching for substitutes, sodium ions, zinc ions, aluminum ions, and other metal ions have been investigated as charge carriers. Non-lithium electrochemical energy storage systems have shown signs of progress. However, no matter how they were improved, metal dendrites have still remained present during the anodic process. To address this limitation, the non-metal ionic liquid EMIm+[PF6]− has been employed as an electrolyte to realize charge transport. Unfortunately, due to the larger radius of EMIm+ (vertical: 4.3 Å; horizontal: 7.6 Å), it is not possible for traditional graphite anodes to provide a high intercalation capacity. Selecting new electrode materials to achieve the intercalation/deintercalation of EMIm+ can be a route to overcome this bottleneck limitation.

What's the anode?

graphite

In order to achieve higher power/energy densities for electrochemical energy storage, dual-ion batteries (DIBs), relying on the migration of both cations and anions to store charge concurrently, have been developed. Compared with lithium ion batteries, dual-ion migration behavior can shorten the charge carrier distance. From a cation perspective, lithium ions, as a charge carrier, have contributed to energy storage for several decades. However, many issues have also been raised, such as safety, excessive resource consumption, and environmental pollution. In the process of searching for substitutes, sodium ions, zinc ions, aluminum ions, and other metal ions have been investigated as charge carriers. Non-lithium electrochemical energy storage systems have shown signs of progress. However, no matter how they were improved, metal dendrites have still remained present during the anodic process. To address this limitation, the non-metal ionic liquid EMIm+[PF6]− has been employed as an electrolyte to realize charge transport. Unfortunately, due to the larger radius of EMIm+ (vertical: 4.3 Å; horizontal: 7.6 Å), it is not possible for traditional graphite anodes to provide a high intercalation capacity. Selecting new electrode materials to achieve the intercalation/deintercalation of EMIm+ can be a route to overcome this bottleneck limitation.

What's the electrolyte?

EMIm+[PF6]−

To better visualize the local contact loss distribution around the cathode particles, a single NMC particle and its surrounding voids in the sample cycled 50 times were isolated from the larger volume and are shown in Fig. 4(f). The spherical NMC particle at the center is shown in grey, and the surrounding voids are shown in red. The voids are much more concentrated on the right side of the NMC particle, suggesting that the contact loss areas caused by the cathode particle volume shrinking will mostly be on one side of the particle, rather than uniformly distributed around the particle. This is expected because the presence of voids on one side of the particle is sufficient to release the strain caused by the particle volume change. This concentration of the decohered area in the cathode particle may aggravate its effect on capacity loss as now all the Li ions going in and out of the area under the contact loss have a much larger distance to travel before they reach the solid electrolyte. This is in contrast to if the contact loss area were homogeneously dispersed across the cathode particle. In the latter case, diffusion gradients parallel to the surface, resulting from blocked-out surface area for Li to enter or leave the particle, would rapidly fade out in the cathode particle.

What's the cathode?

To better visualize the local contact loss distribution around the cathode particles, a single NMC particle and its surrounding voids in the sample cycled 50 times were isolated from the larger volume and are shown in Fig. 4(f). The spherical NMC particle at the center is shown in grey, and the surrounding voids are shown in red. The voids are much more concentrated on the right side of the NMC particle, suggesting that the contact loss areas caused by the cathode particle volume shrinking will mostly be on one side of the particle, rather than uniformly distributed around the particle. This is expected because the presence of voids on one side of the particle is sufficient to release the strain caused by the particle volume change. This concentration of the decohered area in the cathode particle may aggravate its effect on capacity loss as now all the Li ions going in and out of the area under the contact loss have a much larger distance to travel before they reach the solid electrolyte. This is in contrast to if the contact loss area were homogeneously dispersed across the cathode particle. In the latter case, diffusion gradients parallel to the surface, resulting from blocked-out surface area for Li to enter or leave the particle, would rapidly fade out in the cathode particle.

What's the electrolyte?

Even though the SEI film is supposed to act as an insulator being a barrier for electron transfer, at the initial nucleation stages its structure may favor electron transport. This may occur because of the differences between the amorphous character of nucleating crystals and their bulk theoretical structures, as has been found in some bulk insulator materials where changes in the electronic conductivity were observed in ultra-thin films showing a semiconductor behavior. This capability of conducting electrons of imperfectly formed SEI nuclei and phases could contribute to the growth mechanism that keeps the SEI forming until it reaches hundreds of nanometers. Another important reason for the continuous SEI growth is the presence of radical charged species that result from the SEI reactions and are able to diffuse toward the vicinity of the electrolyte inducing further reactions. Several experimental and theoretical works have studied the various stages of the SEI formation and growth process while other authors have proposed the formation of artificial SEI layers by pretreating the anode before the battery is assembled. For the early stages of the SEI formation it is crucial to improve the understanding of electrolyte decomposition by characterizing the reaction mechanisms of electrolytes typically used in battery systems. It is known that carbonate solvents have low stability when implemented with Li-metal anodes, while ether-based solvents like dimethoxyethane (DME) and dioxalane (DOL) have shown better stability with respect to the Li-metal anode and are frequently used for these battery systems. However, solvent and salt molecules decomposition are still observed and in spite of the advances reached both from experimental and theoretical studies, the mechanisms of these processes are not yet well understood.

What's the anode?

Li-metal

Even though the SEI film is supposed to act as an insulator being a barrier for electron transfer, at the initial nucleation stages its structure may favor electron transport. This may occur because of the differences between the amorphous character of nucleating crystals and their bulk theoretical structures, as has been found in some bulk insulator materials where changes in the electronic conductivity were observed in ultra-thin films showing a semiconductor behavior. This capability of conducting electrons of imperfectly formed SEI nuclei and phases could contribute to the growth mechanism that keeps the SEI forming until it reaches hundreds of nanometers. Another important reason for the continuous SEI growth is the presence of radical charged species that result from the SEI reactions and are able to diffuse toward the vicinity of the electrolyte inducing further reactions. Several experimental and theoretical works have studied the various stages of the SEI formation and growth process while other authors have proposed the formation of artificial SEI layers by pretreating the anode before the battery is assembled. For the early stages of the SEI formation it is crucial to improve the understanding of electrolyte decomposition by characterizing the reaction mechanisms of electrolytes typically used in battery systems. It is known that carbonate solvents have low stability when implemented with Li-metal anodes, while ether-based solvents like dimethoxyethane (DME) and dioxalane (DOL) have shown better stability with respect to the Li-metal anode and are frequently used for these battery systems. However, solvent and salt molecules decomposition are still observed and in spite of the advances reached both from experimental and theoretical studies, the mechanisms of these processes are not yet well understood.

What's the electrolyte?

Even though the SEI film is supposed to act as an insulator being a barrier for electron transfer, at the initial nucleation stages its structure may favor electron transport. This may occur because of the differences between the amorphous character of nucleating crystals and their bulk theoretical structures, as has been found in some bulk insulator materials where changes in the electronic conductivity were observed in ultra-thin films showing a semiconductor behavior. This capability of conducting electrons of imperfectly formed SEI nuclei and phases could contribute to the growth mechanism that keeps the SEI forming until it reaches hundreds of nanometers. Another important reason for the continuous SEI growth is the presence of radical charged species that result from the SEI reactions and are able to diffuse toward the vicinity of the electrolyte inducing further reactions. Several experimental and theoretical works have studied the various stages of the SEI formation and growth process while other authors have proposed the formation of artificial SEI layers by pretreating the anode before the battery is assembled. For the early stages of the SEI formation it is crucial to improve the understanding of electrolyte decomposition by characterizing the reaction mechanisms of electrolytes typically used in battery systems. It is known that carbonate solvents have low stability when implemented with Li-metal anodes, while ether-based solvents like dimethoxyethane (DME) and dioxalane (DOL) have shown better stability with respect to the Li-metal anode and are frequently used for these battery systems. However, solvent and salt molecules decomposition are still observed and in spite of the advances reached both from experimental and theoretical studies, the mechanisms of these processes are not yet well understood.

What's the anode?

Li-metal

The electrochemical performance of the 400-KOH-Ti3C2 anode was measured in a three-electrode system with the 2 M H2SO4 solution serving as the electrolyte. To determine the potential range of the 400-KOH-Ti3C2 electrode in this acid electrolyte, the CV curves of the 400-KOH-Ti3C2 electrode were collected in various potential windows of −0.7 to 0.3 V, −0.6 to 0.3 V, −0.5 to 0.3 V and −0.4 to 0.3 V (vs. Ag/AgCl) at a scan rate of 5 mV s−1 (Fig. S6†). It is found that the potential window of the 400-KOH-Ti3C2 electrode can be stably extended to −0.6 V (vs. Ag/AgCl) without the decomposition of the electrolyte. Thus, the potential range of the 400-KOH-Ti3C2 electrode is fixed at −0.6–0.3 V (vs. Ag/AgCl) in 2 M H2SO4 solution. To further evaluate the stable potential window of the 400-KOH-Ti3C2 electrode, we collected CV curves at various scan rates in the potential window of −0.6–0.3 V (vs. Ag/AgCl) as shown in Fig. 2a. The shape of CV curves can be well retained even at a high scan rate of 200 mV s−1, indicating good rate capability of the 400-KOH-Ti3C2 electrode. Except for CV measurements, the galvanostatic discharge–charge profiles of the 400-KOH-Ti3C2 electrode at various current densities were recorded as displayed in Fig. 2b. The discharge–charge curves exhibit a triangular symmetric shape with an inappreciable deviation, suggesting high reversibility and prominent pseudocapacitive behavior. Based on the above discharge–charge curves, the specific capacitance of the 400-KOH-Ti3C2 electrode is calculated to be 334.5 F g−1 at 1 A g−1. Even at 100 A g−1, the 400-KOH-Ti3C2 electrode still retains a specific capacitance of 93.3 F g−1 (Fig. 2c). Such high specific capacitance of the 400-KOH-Ti3C2 electrode should be attributed to the large interlayer distance, which provides more accessible active sites for the electrolyte ionic diffusion and electrochemical reaction. To illustrate this view, EIS measurements were performed at a potential of −0.1 V (vs. Ag/AgCl) for the 400-KOH-Ti3C2 electrode. The corresponding Nyquist plot of the 400-KOH-Ti3C2 electrode (Fig. 2d) shows a negligible semicircle and a quite small charge transfer resistance (Rct) of 1.2 Ω, which is obviously lower than that of original Ti3C2Tx (1.8 Ω, Fig. S7†), revealing a fast ion transport and good electronic conductivity for the 400-KOH-Ti3C2 electrode. The Nyquist plots are fitted with the equivalent circuit model shown in Fig. S8.† In addition, the 400-KOH-Ti3C2 electrode shows a stable cycle performance at 4 A g−1, maintaining a capacitance retention of about 87% over 40000 cycles and almost 100% coulombic efficiency. Such excellent cycling stability is comparable to that of most recently reported electrode materials for supercapacitors (see Table S1†). Furthermore, the SEM images of the 400-KOH-Ti3C2 electrode after 5000 cycles show that the layered structure of 400-KOH-Ti3C2 is well retained after cycling (Fig. S9†). It is noteworthy that the capacitance of the 400-KOH-Ti3C2 electrode decreases significantly in the first few hundred cycles, which is likely related to the phase transition of the electrode material.

What's the anode?

400-KOH-Ti3C2

The electrochemical performance of the 400-KOH-Ti3C2 anode was measured in a three-electrode system with the 2 M H2SO4 solution serving as the electrolyte. To determine the potential range of the 400-KOH-Ti3C2 electrode in this acid electrolyte, the CV curves of the 400-KOH-Ti3C2 electrode were collected in various potential windows of −0.7 to 0.3 V, −0.6 to 0.3 V, −0.5 to 0.3 V and −0.4 to 0.3 V (vs. Ag/AgCl) at a scan rate of 5 mV s−1 (Fig. S6†). It is found that the potential window of the 400-KOH-Ti3C2 electrode can be stably extended to −0.6 V (vs. Ag/AgCl) without the decomposition of the electrolyte. Thus, the potential range of the 400-KOH-Ti3C2 electrode is fixed at −0.6–0.3 V (vs. Ag/AgCl) in 2 M H2SO4 solution. To further evaluate the stable potential window of the 400-KOH-Ti3C2 electrode, we collected CV curves at various scan rates in the potential window of −0.6–0.3 V (vs. Ag/AgCl) as shown in Fig. 2a. The shape of CV curves can be well retained even at a high scan rate of 200 mV s−1, indicating good rate capability of the 400-KOH-Ti3C2 electrode. Except for CV measurements, the galvanostatic discharge–charge profiles of the 400-KOH-Ti3C2 electrode at various current densities were recorded as displayed in Fig. 2b. The discharge–charge curves exhibit a triangular symmetric shape with an inappreciable deviation, suggesting high reversibility and prominent pseudocapacitive behavior. Based on the above discharge–charge curves, the specific capacitance of the 400-KOH-Ti3C2 electrode is calculated to be 334.5 F g−1 at 1 A g−1. Even at 100 A g−1, the 400-KOH-Ti3C2 electrode still retains a specific capacitance of 93.3 F g−1 (Fig. 2c). Such high specific capacitance of the 400-KOH-Ti3C2 electrode should be attributed to the large interlayer distance, which provides more accessible active sites for the electrolyte ionic diffusion and electrochemical reaction. To illustrate this view, EIS measurements were performed at a potential of −0.1 V (vs. Ag/AgCl) for the 400-KOH-Ti3C2 electrode. The corresponding Nyquist plot of the 400-KOH-Ti3C2 electrode (Fig. 2d) shows a negligible semicircle and a quite small charge transfer resistance (Rct) of 1.2 Ω, which is obviously lower than that of original Ti3C2Tx (1.8 Ω, Fig. S7†), revealing a fast ion transport and good electronic conductivity for the 400-KOH-Ti3C2 electrode. The Nyquist plots are fitted with the equivalent circuit model shown in Fig. S8.† In addition, the 400-KOH-Ti3C2 electrode shows a stable cycle performance at 4 A g−1, maintaining a capacitance retention of about 87% over 40000 cycles and almost 100% coulombic efficiency. Such excellent cycling stability is comparable to that of most recently reported electrode materials for supercapacitors (see Table S1†). Furthermore, the SEM images of the 400-KOH-Ti3C2 electrode after 5000 cycles show that the layered structure of 400-KOH-Ti3C2 is well retained after cycling (Fig. S9†). It is noteworthy that the capacitance of the 400-KOH-Ti3C2 electrode decreases significantly in the first few hundred cycles, which is likely related to the phase transition of the electrode material.

What's the electrolyte?

2 M H2SO4

Using density functional theory (DFT), we calculated the free energy of hydrogen (ΔGH) on various coating materials as well as their electronic conductivity, and used these two parameters to screen for the best candidate material. Materials with suitable hydrogen adsorption (neither too strong nor too weak) located at the top of volcano plots would have optimum activity. Hence, materials with either very strong (highly negative −ΔGH) or very weak bonding (highly positive +ΔGH) should be ideal candidates to suppress the HER. The high ΔGH of Al2O3 (Fig. 1b) suggested its suitability whereas TiO2, despite its high ΔGH, became highly catalytic towards the HER once lithiated into LiTi2O4 (as characterized by a low ΔGH). This increased catalytic activity according to the degree of lithiation might explain why LTO in non-aqueous batteries effectively catalyzed the decomposition of trace H2O present in electrodes and electrolyte. With a highly negative ΔGH, ZnO would also serve as a good candidate for HER suppression. However, its high electronic conductivity accelerated the charge transfer between the electrolyte and electrodes, which kinetically favors the electrochemical reduction of water. Conversely, Al2O3 had a large bandgap (5.33 eV) (Fig. 1c), which suggested very poor electrical conductivity that prevented charge transfer, in sharp comparison with the bandgaps of lithiated LiTi2O4 (0 eV) and ZnO (1.48 eV), respectively. In fact, the latter two were well known for their metallic and semiconducting behaviors (Fig. 1d and S1†). Therefore, the insulator Al2O3 would be the best candidate surface for HER suppression.

What's the electrolyte?

To facilitate the practical realization of sodium-ion batteries, the energy density, determined by the output operating voltage and/or capacity, needs to be improved to the level of commercial Li-ion batteries. Herein, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 is synthesized by incorporating Ca2+ into the NaO6 octahedron of Na[Ni0.5Mn0.5]O2 and its potential use as a cathode material for high energy density SIBs is demonstrated. The ionic radius of calcium (≈1.00 Å) is similar to that of sodium (≈1.02 Å); hence, it is energetically favorable for calcium to occupy sites in the sodium layers. Within a wide operating voltage range of 2.0–4.3 V, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 exhibits a reversible O3–P3–O3 phase transition with small volume changes compared to Ca-free Na[Ni0.5Mn0.5]O2 because of the strong interaction between Ca2+ and O2− and delivers a high reversible capacity of 209 mA h g−1 at 15 mA g−1 with improved cycling stability. Moreover, Ca substitution improves the practically useful aspects such as thermal and air stability. A prototype pouch full cell with a hard carbon anode shows an excellent capacity retention of 67% over 300 cycles. Thus, this study provides an efficient and simple method to boost the performance and applicability of layered oxide cathode materials for practical applications.

What's the cathode?

Na[Ni0.5Mn0.5]O2

To facilitate the practical realization of sodium-ion batteries, the energy density, determined by the output operating voltage and/or capacity, needs to be improved to the level of commercial Li-ion batteries. Herein, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 is synthesized by incorporating Ca2+ into the NaO6 octahedron of Na[Ni0.5Mn0.5]O2 and its potential use as a cathode material for high energy density SIBs is demonstrated. The ionic radius of calcium (≈1.00 Å) is similar to that of sodium (≈1.02 Å); hence, it is energetically favorable for calcium to occupy sites in the sodium layers. Within a wide operating voltage range of 2.0–4.3 V, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 exhibits a reversible O3–P3–O3 phase transition with small volume changes compared to Ca-free Na[Ni0.5Mn0.5]O2 because of the strong interaction between Ca2+ and O2− and delivers a high reversible capacity of 209 mA h g−1 at 15 mA g−1 with improved cycling stability. Moreover, Ca substitution improves the practically useful aspects such as thermal and air stability. A prototype pouch full cell with a hard carbon anode shows an excellent capacity retention of 67% over 300 cycles. Thus, this study provides an efficient and simple method to boost the performance and applicability of layered oxide cathode materials for practical applications.

What's the anode?

hard carbon

To facilitate the practical realization of sodium-ion batteries, the energy density, determined by the output operating voltage and/or capacity, needs to be improved to the level of commercial Li-ion batteries. Herein, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 is synthesized by incorporating Ca2+ into the NaO6 octahedron of Na[Ni0.5Mn0.5]O2 and its potential use as a cathode material for high energy density SIBs is demonstrated. The ionic radius of calcium (≈1.00 Å) is similar to that of sodium (≈1.02 Å); hence, it is energetically favorable for calcium to occupy sites in the sodium layers. Within a wide operating voltage range of 2.0–4.3 V, O3-type Na0.98Ca0.01[Ni0.5Mn0.5]O2 exhibits a reversible O3–P3–O3 phase transition with small volume changes compared to Ca-free Na[Ni0.5Mn0.5]O2 because of the strong interaction between Ca2+ and O2− and delivers a high reversible capacity of 209 mA h g−1 at 15 mA g−1 with improved cycling stability. Moreover, Ca substitution improves the practically useful aspects such as thermal and air stability. A prototype pouch full cell with a hard carbon anode shows an excellent capacity retention of 67% over 300 cycles. Thus, this study provides an efficient and simple method to boost the performance and applicability of layered oxide cathode materials for practical applications.

What's the electrolyte?

To overcome this limitation, analysis of Differential Capacity Plots (DCPs) has been performed to identify electrochemical charge/discharge reactions and track their evolution during cycling. Fig. 6 shows the DCPs for the different Si/Ni–Sn/Al/C composites at cycles 2, 80 and 200 corresponding to composite activation, maximum capacity and end-of-cycling states. For Si-B, at the 2nd lithiation (Fig. 6a), three reduction peaks are observed in the cathodic branch. The peak at 0.17 V is attributed to the reaction potential of amorphous Si with lithium to form LixSi alloys of approximate composition LiSi and Li7Si3. Silicon amorphisation is known to occur during the first lithiation of crystalline Si anodes. The two other reduction peaks at 0.31 and 0.62 V are attributed to the formation of LiySn alloys: Li7Sn2 and LiSn, respectively. On delithiation, multiple oxidation peaks are detected and are attributed to the decomposition of Li7Si3 at 0.51 V and different LiySn alloys: Li7Sn2 (or Li22Sn5), Li5Sn2, LiSn and Li2Sn5 at 0.47, 0.58, 0.74 and 0.81 V, respectively. The two phases Li7Sn3 and Li13Sn5 have compositions close to Li5Sn2 and could be involved in the delithiation process at 0.58 V. For both composites Si-AB and Si-A, poorly defined cathodic peaks are found at the 2nd lithiation, but DCPs exhibit significant features in delithiation. For Si-AB, two oxidation peaks are clearly identified at 0.47 and 0.58 V attributed to the decomposition of Li7Sn2 (or Li22Sn5) and Li5Sn2 alloys, respectively. For Si-A, a unique broad oxidation peak is observed at 0.51 V assigned to the decomposition of Li7Si3. Interestingly, no signal related to the decomposition of LiySn alloys is detected for the composite Si-A, suggesting that Ni3Sn2 does not react with lithium.

What's the anode?

Si

To overcome this limitation, analysis of Differential Capacity Plots (DCPs) has been performed to identify electrochemical charge/discharge reactions and track their evolution during cycling. Fig. 6 shows the DCPs for the different Si/Ni–Sn/Al/C composites at cycles 2, 80 and 200 corresponding to composite activation, maximum capacity and end-of-cycling states. For Si-B, at the 2nd lithiation (Fig. 6a), three reduction peaks are observed in the cathodic branch. The peak at 0.17 V is attributed to the reaction potential of amorphous Si with lithium to form LixSi alloys of approximate composition LiSi and Li7Si3. Silicon amorphisation is known to occur during the first lithiation of crystalline Si anodes. The two other reduction peaks at 0.31 and 0.62 V are attributed to the formation of LiySn alloys: Li7Sn2 and LiSn, respectively. On delithiation, multiple oxidation peaks are detected and are attributed to the decomposition of Li7Si3 at 0.51 V and different LiySn alloys: Li7Sn2 (or Li22Sn5), Li5Sn2, LiSn and Li2Sn5 at 0.47, 0.58, 0.74 and 0.81 V, respectively. The two phases Li7Sn3 and Li13Sn5 have compositions close to Li5Sn2 and could be involved in the delithiation process at 0.58 V. For both composites Si-AB and Si-A, poorly defined cathodic peaks are found at the 2nd lithiation, but DCPs exhibit significant features in delithiation. For Si-AB, two oxidation peaks are clearly identified at 0.47 and 0.58 V attributed to the decomposition of Li7Sn2 (or Li22Sn5) and Li5Sn2 alloys, respectively. For Si-A, a unique broad oxidation peak is observed at 0.51 V assigned to the decomposition of Li7Si3. Interestingly, no signal related to the decomposition of LiySn alloys is detected for the composite Si-A, suggesting that Ni3Sn2 does not react with lithium.

What's the electrolyte?

A sodium-based analog of 3a, i.e. sodium rhodizonate 3b, was tested by Goodenough et al. with an anode based on a liquid sodium–potassium alloy (NaK). The advantage of this anode is its dendrite-free nature, making it safer compared to pristine potassium. In a DME-based KFSI solution, Qm was ∼120 mA h g−1 at 125 mA g−1, with ∼50 mA h g−1 delivered at 1.25 A g−1. The capacity retention was ca. 75% after 100 cycles.

What's the anode?

A sodium-based analog of 3a, i.e. sodium rhodizonate 3b, was tested by Goodenough et al. with an anode based on a liquid sodium–potassium alloy (NaK). The advantage of this anode is its dendrite-free nature, making it safer compared to pristine potassium. In a DME-based KFSI solution, Qm was ∼120 mA h g−1 at 125 mA g−1, with ∼50 mA h g−1 delivered at 1.25 A g−1. The capacity retention was ca. 75% after 100 cycles.

What's the electrolyte?

The Zn symmetric and Zn–Ti asymmetric cells were assembled as CR2030 coin cells to measure the electrochemical performances. Glass fiber and a 3 M Zn(CF3SO3)2 aqueous solution were used as the separator and the electrolyte, respectively. Before cell fabrication, Ti foil, bare Zn foil and PAM/PVP-coated Zn were cut into disk-shaped electrodes. The symmetric cells consist of two bare Zn foils (or PAM/PVP-coated Zn foils) separated by a glass fiber separator. The cells underwent galvanostatic charge and discharge cycling at a current density of 0.2–10 mA cm−2 on a Neware battery testing instrument. These experiments were performed to evaluate the stripping/plating behavior and cycling stability of the bare Zn and PAM/PVP-coated Zn anodes. Furthermore, Zn–Ti asymmetric cells were fabricated to investigate the coulombic efficiency of Zn stripping/plating. The asymmetric cell was initially discharged for 10 min, and then to 2.0 V at various current densities from 0.2–20 mA cm−2. To obtain the corrosion potential (Ecorr) and corrosion current density (Icorr), Tafel plots of the Zn symmetric cells were recorded on an electrochemical workstation (CHI660E, China) at a scan rate of 10 mV s−1. Chronoamperograms (CAs) of the bare Zn and the PAM/PVP-coated Zn symmetric cells were recorded under a −200 mV overpotential. The Zn–AC hybrid supercapacitors were fabricated with AC as the cathode, bare Zn or PAM/PVP-coated Zn foil as the anode, and 3 M Zn(CF3SO3)2 aqueous solution as the electrolyte. The cathode slurry was fabricated by mixing the cathode active materials, Super P, and PTFE at a mass ratio of 80:10:10. The resulting slurry was rolled into a thin film with a thickness of about 120 μm. Finally, the film was cut into discs with a diameter of 10 mm, pressed into a Ti mesh (15 mm in diameter) and dried in a vacuum oven at 100 °C for 6 h. The electrochemical performance was evaluated by cyclic voltammetry on an electrochemical workstation. The galvanostatic charge–discharge processes of the Zn–AC hybrid supercapacitor were recorded on a Neware battery testing system.

What's the cathode?

AC

The Zn symmetric and Zn–Ti asymmetric cells were assembled as CR2030 coin cells to measure the electrochemical performances. Glass fiber and a 3 M Zn(CF3SO3)2 aqueous solution were used as the separator and the electrolyte, respectively. Before cell fabrication, Ti foil, bare Zn foil and PAM/PVP-coated Zn were cut into disk-shaped electrodes. The symmetric cells consist of two bare Zn foils (or PAM/PVP-coated Zn foils) separated by a glass fiber separator. The cells underwent galvanostatic charge and discharge cycling at a current density of 0.2–10 mA cm−2 on a Neware battery testing instrument. These experiments were performed to evaluate the stripping/plating behavior and cycling stability of the bare Zn and PAM/PVP-coated Zn anodes. Furthermore, Zn–Ti asymmetric cells were fabricated to investigate the coulombic efficiency of Zn stripping/plating. The asymmetric cell was initially discharged for 10 min, and then to 2.0 V at various current densities from 0.2–20 mA cm−2. To obtain the corrosion potential (Ecorr) and corrosion current density (Icorr), Tafel plots of the Zn symmetric cells were recorded on an electrochemical workstation (CHI660E, China) at a scan rate of 10 mV s−1. Chronoamperograms (CAs) of the bare Zn and the PAM/PVP-coated Zn symmetric cells were recorded under a −200 mV overpotential. The Zn–AC hybrid supercapacitors were fabricated with AC as the cathode, bare Zn or PAM/PVP-coated Zn foil as the anode, and 3 M Zn(CF3SO3)2 aqueous solution as the electrolyte. The cathode slurry was fabricated by mixing the cathode active materials, Super P, and PTFE at a mass ratio of 80:10:10. The resulting slurry was rolled into a thin film with a thickness of about 120 μm. Finally, the film was cut into discs with a diameter of 10 mm, pressed into a Ti mesh (15 mm in diameter) and dried in a vacuum oven at 100 °C for 6 h. The electrochemical performance was evaluated by cyclic voltammetry on an electrochemical workstation. The galvanostatic charge–discharge processes of the Zn–AC hybrid supercapacitor were recorded on a Neware battery testing system.

What's the anode?

the bare Zn and PAM/PVP-coated Zn

The Zn symmetric and Zn–Ti asymmetric cells were assembled as CR2030 coin cells to measure the electrochemical performances. Glass fiber and a 3 M Zn(CF3SO3)2 aqueous solution were used as the separator and the electrolyte, respectively. Before cell fabrication, Ti foil, bare Zn foil and PAM/PVP-coated Zn were cut into disk-shaped electrodes. The symmetric cells consist of two bare Zn foils (or PAM/PVP-coated Zn foils) separated by a glass fiber separator. The cells underwent galvanostatic charge and discharge cycling at a current density of 0.2–10 mA cm−2 on a Neware battery testing instrument. These experiments were performed to evaluate the stripping/plating behavior and cycling stability of the bare Zn and PAM/PVP-coated Zn anodes. Furthermore, Zn–Ti asymmetric cells were fabricated to investigate the coulombic efficiency of Zn stripping/plating. The asymmetric cell was initially discharged for 10 min, and then to 2.0 V at various current densities from 0.2–20 mA cm−2. To obtain the corrosion potential (Ecorr) and corrosion current density (Icorr), Tafel plots of the Zn symmetric cells were recorded on an electrochemical workstation (CHI660E, China) at a scan rate of 10 mV s−1. Chronoamperograms (CAs) of the bare Zn and the PAM/PVP-coated Zn symmetric cells were recorded under a −200 mV overpotential. The Zn–AC hybrid supercapacitors were fabricated with AC as the cathode, bare Zn or PAM/PVP-coated Zn foil as the anode, and 3 M Zn(CF3SO3)2 aqueous solution as the electrolyte. The cathode slurry was fabricated by mixing the cathode active materials, Super P, and PTFE at a mass ratio of 80:10:10. The resulting slurry was rolled into a thin film with a thickness of about 120 μm. Finally, the film was cut into discs with a diameter of 10 mm, pressed into a Ti mesh (15 mm in diameter) and dried in a vacuum oven at 100 °C for 6 h. The electrochemical performance was evaluated by cyclic voltammetry on an electrochemical workstation. The galvanostatic charge–discharge processes of the Zn–AC hybrid supercapacitor were recorded on a Neware battery testing system.

What's the electrolyte?

3 M Zn(CF3SO3)2 aqueous solution

The Zn symmetric and Zn–Ti asymmetric cells were assembled as CR2030 coin cells to measure the electrochemical performances. Glass fiber and a 3 M Zn(CF3SO3)2 aqueous solution were used as the separator and the electrolyte, respectively. Before cell fabrication, Ti foil, bare Zn foil and PAM/PVP-coated Zn were cut into disk-shaped electrodes. The symmetric cells consist of two bare Zn foils (or PAM/PVP-coated Zn foils) separated by a glass fiber separator. The cells underwent galvanostatic charge and discharge cycling at a current density of 0.2–10 mA cm−2 on a Neware battery testing instrument. These experiments were performed to evaluate the stripping/plating behavior and cycling stability of the bare Zn and PAM/PVP-coated Zn anodes. Furthermore, Zn–Ti asymmetric cells were fabricated to investigate the coulombic efficiency of Zn stripping/plating. The asymmetric cell was initially discharged for 10 min, and then to 2.0 V at various current densities from 0.2–20 mA cm−2. To obtain the corrosion potential (Ecorr) and corrosion current density (Icorr), Tafel plots of the Zn symmetric cells were recorded on an electrochemical workstation (CHI660E, China) at a scan rate of 10 mV s−1. Chronoamperograms (CAs) of the bare Zn and the PAM/PVP-coated Zn symmetric cells were recorded under a −200 mV overpotential. The Zn–AC hybrid supercapacitors were fabricated with AC as the cathode, bare Zn or PAM/PVP-coated Zn foil as the anode, and 3 M Zn(CF3SO3)2 aqueous solution as the electrolyte. The cathode slurry was fabricated by mixing the cathode active materials, Super P, and PTFE at a mass ratio of 80:10:10. The resulting slurry was rolled into a thin film with a thickness of about 120 μm. Finally, the film was cut into discs with a diameter of 10 mm, pressed into a Ti mesh (15 mm in diameter) and dried in a vacuum oven at 100 °C for 6 h. The electrochemical performance was evaluated by cyclic voltammetry on an electrochemical workstation. The galvanostatic charge–discharge processes of the Zn–AC hybrid supercapacitor were recorded on a Neware battery testing system.

What's the anode?

bare Zn or PAM/PVP-coated Zn foil

The Zn symmetric and Zn–Ti asymmetric cells were assembled as CR2030 coin cells to measure the electrochemical performances. Glass fiber and a 3 M Zn(CF3SO3)2 aqueous solution were used as the separator and the electrolyte, respectively. Before cell fabrication, Ti foil, bare Zn foil and PAM/PVP-coated Zn were cut into disk-shaped electrodes. The symmetric cells consist of two bare Zn foils (or PAM/PVP-coated Zn foils) separated by a glass fiber separator. The cells underwent galvanostatic charge and discharge cycling at a current density of 0.2–10 mA cm−2 on a Neware battery testing instrument. These experiments were performed to evaluate the stripping/plating behavior and cycling stability of the bare Zn and PAM/PVP-coated Zn anodes. Furthermore, Zn–Ti asymmetric cells were fabricated to investigate the coulombic efficiency of Zn stripping/plating. The asymmetric cell was initially discharged for 10 min, and then to 2.0 V at various current densities from 0.2–20 mA cm−2. To obtain the corrosion potential (Ecorr) and corrosion current density (Icorr), Tafel plots of the Zn symmetric cells were recorded on an electrochemical workstation (CHI660E, China) at a scan rate of 10 mV s−1. Chronoamperograms (CAs) of the bare Zn and the PAM/PVP-coated Zn symmetric cells were recorded under a −200 mV overpotential. The Zn–AC hybrid supercapacitors were fabricated with AC as the cathode, bare Zn or PAM/PVP-coated Zn foil as the anode, and 3 M Zn(CF3SO3)2 aqueous solution as the electrolyte. The cathode slurry was fabricated by mixing the cathode active materials, Super P, and PTFE at a mass ratio of 80:10:10. The resulting slurry was rolled into a thin film with a thickness of about 120 μm. Finally, the film was cut into discs with a diameter of 10 mm, pressed into a Ti mesh (15 mm in diameter) and dried in a vacuum oven at 100 °C for 6 h. The electrochemical performance was evaluated by cyclic voltammetry on an electrochemical workstation. The galvanostatic charge–discharge processes of the Zn–AC hybrid supercapacitor were recorded on a Neware battery testing system.

What's the electrolyte?

3 M Zn(CF3SO3)2 aqueous solution

Despite the merits of zinc anode, ARZIBs are still at a development stage due to the lack of suitable Zn-ion host materials. To date, various materials have been explored as cathodes for ARZIBs, such as manganese-based oxides, Prussian blue analogues, polyanionic compounds, and vanadium-based compounds. However, most of the reported cathode materials suffer from insufficient rate capability, low capacity, and energy density. The main reasons can be ascribed to three factors: (i) the relativity poor conductivity leads to sluggish ion/electron transport kinetics, especially in quasi-solid-state electrolytes; (ii) the use of non-active additives such as polyvinylidene fluoride and polypyrrole, which were usually employed to ameliorate the flexibility of powder-formed active materials, would inevitably increase the “dead mass” of the cathode, thus limiting the overall energy density of the full devices; (iii) the increased contact resistance between the active materials and the current collector, and the interfacial resistance between the electrode and electrolyte, arising from the repeated reversible Zn2+ (de)intercalation processes, further result in a decrease in the high rate ability. Therefore, there is an urgent demand to explore free-standing materials that can adequately accommodate Zn2+ insertion/extraction rapidly and steadily to promote the progress of ARZIBs.

What's the cathode?

manganese-based oxides, Prussian blue analogues, polyanionic compounds, and vanadium-based compounds

Despite the merits of zinc anode, ARZIBs are still at a development stage due to the lack of suitable Zn-ion host materials. To date, various materials have been explored as cathodes for ARZIBs, such as manganese-based oxides, Prussian blue analogues, polyanionic compounds, and vanadium-based compounds. However, most of the reported cathode materials suffer from insufficient rate capability, low capacity, and energy density. The main reasons can be ascribed to three factors: (i) the relativity poor conductivity leads to sluggish ion/electron transport kinetics, especially in quasi-solid-state electrolytes; (ii) the use of non-active additives such as polyvinylidene fluoride and polypyrrole, which were usually employed to ameliorate the flexibility of powder-formed active materials, would inevitably increase the “dead mass” of the cathode, thus limiting the overall energy density of the full devices; (iii) the increased contact resistance between the active materials and the current collector, and the interfacial resistance between the electrode and electrolyte, arising from the repeated reversible Zn2+ (de)intercalation processes, further result in a decrease in the high rate ability. Therefore, there is an urgent demand to explore free-standing materials that can adequately accommodate Zn2+ insertion/extraction rapidly and steadily to promote the progress of ARZIBs.

What's the anode?

zinc

Despite the merits of zinc anode, ARZIBs are still at a development stage due to the lack of suitable Zn-ion host materials. To date, various materials have been explored as cathodes for ARZIBs, such as manganese-based oxides, Prussian blue analogues, polyanionic compounds, and vanadium-based compounds. However, most of the reported cathode materials suffer from insufficient rate capability, low capacity, and energy density. The main reasons can be ascribed to three factors: (i) the relativity poor conductivity leads to sluggish ion/electron transport kinetics, especially in quasi-solid-state electrolytes; (ii) the use of non-active additives such as polyvinylidene fluoride and polypyrrole, which were usually employed to ameliorate the flexibility of powder-formed active materials, would inevitably increase the “dead mass” of the cathode, thus limiting the overall energy density of the full devices; (iii) the increased contact resistance between the active materials and the current collector, and the interfacial resistance between the electrode and electrolyte, arising from the repeated reversible Zn2+ (de)intercalation processes, further result in a decrease in the high rate ability. Therefore, there is an urgent demand to explore free-standing materials that can adequately accommodate Zn2+ insertion/extraction rapidly and steadily to promote the progress of ARZIBs.

What's the electrolyte?

The electrochemical performance of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors was tested by cyclic voltammetry and galvanostatic charge–discharge tests within the voltage range of 0–2.0 V (Fig. 5). The CV curves of the hybrid ion capacitors at various scan rates from 10 mV s−1 to 200 mV s−1 maintain a quasi-rectangular shape, suggesting an EDLC-type behavior (Fig. 5a). Fig. 5b shows the voltage curves of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors at various current densities from 0.5–30 A g−1. The symmetric and linear shape of the curves in the graph can be observed corresponding to the results of the CV test. At a current density of 1 A g−1, the bare Zn–AC hybrid ion supercapacitors failed after 10 cycles, while the PAM/PVP-coated Zn–AC hybrid ion supercapacitors show a good cycling stability (Fig. S10†). This superior stability can be attributed to the critical role of the PAM/PVP layer coated on the zinc anode acting as an inhibitor on the O2/water interphase and a distributor of ions during the zinc deposition. The PAM/PVP-coated Zn–AC hybrid ion supercapacitors show an excellent rate capability of 336, 292, 264, 232, 220, 195, 190, 175, and 165 F g−1 at a current density of 0.5, 1, 2, 5, 8, 10, 15, 20, 25, and 30 A g−1 (Fig. 5c and S11†). The galvanostatic charge–discharge curves and long cycling stability of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors at a current density of 15 A g−1 are shown in Fig. 5d and e. The specific capacitance of the devices slightly increases at first due to activation, then remains quite stable, and ultimately gives a high retention of 100% of the initial specific capacitance after 6000 cycles. These results further demonstrate the excellent electrochemical performance of the Zn–AC hybrid ion supercapacitors with the PAM/PVP layer. Notably, the PAM/PVP-coated Zn–AC hybrid ion supercapacitors have a high energy density of 118 W h−1 kg−1 and a power density of 17.9 kW kg−1 (based on the mass of AC materials, Fig. S12†), where the energy density is superior to those of the bare Zn–AC hybrid ion supercapacitors (Fig. S13 and Table S1†).

What's the anode?

zinc

The electrochemical performance of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors was tested by cyclic voltammetry and galvanostatic charge–discharge tests within the voltage range of 0–2.0 V (Fig. 5). The CV curves of the hybrid ion capacitors at various scan rates from 10 mV s−1 to 200 mV s−1 maintain a quasi-rectangular shape, suggesting an EDLC-type behavior (Fig. 5a). Fig. 5b shows the voltage curves of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors at various current densities from 0.5–30 A g−1. The symmetric and linear shape of the curves in the graph can be observed corresponding to the results of the CV test. At a current density of 1 A g−1, the bare Zn–AC hybrid ion supercapacitors failed after 10 cycles, while the PAM/PVP-coated Zn–AC hybrid ion supercapacitors show a good cycling stability (Fig. S10†). This superior stability can be attributed to the critical role of the PAM/PVP layer coated on the zinc anode acting as an inhibitor on the O2/water interphase and a distributor of ions during the zinc deposition. The PAM/PVP-coated Zn–AC hybrid ion supercapacitors show an excellent rate capability of 336, 292, 264, 232, 220, 195, 190, 175, and 165 F g−1 at a current density of 0.5, 1, 2, 5, 8, 10, 15, 20, 25, and 30 A g−1 (Fig. 5c and S11†). The galvanostatic charge–discharge curves and long cycling stability of the PAM/PVP-coated Zn–AC hybrid ion supercapacitors at a current density of 15 A g−1 are shown in Fig. 5d and e. The specific capacitance of the devices slightly increases at first due to activation, then remains quite stable, and ultimately gives a high retention of 100% of the initial specific capacitance after 6000 cycles. These results further demonstrate the excellent electrochemical performance of the Zn–AC hybrid ion supercapacitors with the PAM/PVP layer. Notably, the PAM/PVP-coated Zn–AC hybrid ion supercapacitors have a high energy density of 118 W h−1 kg−1 and a power density of 17.9 kW kg−1 (based on the mass of AC materials, Fig. S12†), where the energy density is superior to those of the bare Zn–AC hybrid ion supercapacitors (Fig. S13 and Table S1†).

What's the electrolyte?

The intermediate LiPSs generated from charge/discharge processes can spontaneously dissolve into the electrolyte and freely migrate between the cathode and the anode through the separator, resulting in the loss of active materials, passivation of both the electrodes, and unavoidable self-discharge/recharge. So far two typical strategies covering physical confinement and chemical adsorption have been proven and generally accepted as an effective solution to address this issue. Physical confinement often relies on the high surface area or pore structure of the host materials, while chemical adsorption depends on the strong interaction between the host materials and LiPS species. Accordingly, rationally creating metal sulfides with a desired surface area, well-designed porosity, enhanced surface polar/selectivity, and tailored crystalline form has been achieved by a variety of approaches, e.g. solvothermal/hydrothermal methods, ball-milling technology, coprecipitation method, and innovative heat treatment technology.

What's the cathode?

The intermediate LiPSs generated from charge/discharge processes can spontaneously dissolve into the electrolyte and freely migrate between the cathode and the anode through the separator, resulting in the loss of active materials, passivation of both the electrodes, and unavoidable self-discharge/recharge. So far two typical strategies covering physical confinement and chemical adsorption have been proven and generally accepted as an effective solution to address this issue. Physical confinement often relies on the high surface area or pore structure of the host materials, while chemical adsorption depends on the strong interaction between the host materials and LiPS species. Accordingly, rationally creating metal sulfides with a desired surface area, well-designed porosity, enhanced surface polar/selectivity, and tailored crystalline form has been achieved by a variety of approaches, e.g. solvothermal/hydrothermal methods, ball-milling technology, coprecipitation method, and innovative heat treatment technology.

What's the anode?

The intermediate LiPSs generated from charge/discharge processes can spontaneously dissolve into the electrolyte and freely migrate between the cathode and the anode through the separator, resulting in the loss of active materials, passivation of both the electrodes, and unavoidable self-discharge/recharge. So far two typical strategies covering physical confinement and chemical adsorption have been proven and generally accepted as an effective solution to address this issue. Physical confinement often relies on the high surface area or pore structure of the host materials, while chemical adsorption depends on the strong interaction between the host materials and LiPS species. Accordingly, rationally creating metal sulfides with a desired surface area, well-designed porosity, enhanced surface polar/selectivity, and tailored crystalline form has been achieved by a variety of approaches, e.g. solvothermal/hydrothermal methods, ball-milling technology, coprecipitation method, and innovative heat treatment technology.

What's the electrolyte?

Bulk heterojunction (BHJ) polymer solar cells (PSCs) that convert absorbed sunlight into electrical energy have been in the spotlight due to their potential in large-scale and cost-effective roll-to-roll fabrication. In the past decade, massive endeavors concerning the development and interfacial engineering of novel light-harvesting materials have been made to improve the power conversion efficiency (PCE). Nowadays, the highest PCE exceeds 16% for single-junction BHJ PSCs. Interface layers, positioned between the photoactive layer and the anode/cathode electrode, play vital roles in governing the performance of PSCs. Interface layers are usually utilized to tailor the work function of the electrode for charge carrier collection maximization, modify the interface to alter the photoactive layer micromorphology, as well as minimize carrier recombination at the interfaces between the active layer and the transport layer. For a hole transport layer (HTL), pivotal parameters consist of high conductivity, high transparency, solution processability, favorable stability, etc. PEDOT:PSS, as the state-of-the-art hole transport material in organic optoelectronic devices, possesses outstanding edges of high transmittance in the visible spectrum and superior film-forming ability and thermal stability. However, PEDOT:PSS solution is strongly acidic and hygroscopic, which can etch the indium tin oxide (ITO) anode. Moreover, it has been verified that the spin-casted PEDOT:PSS thin film often exhibits inhomogeneous morphology and poor electrical conductivity, leading to poor efficiency and stability of devices. The development of a PEDOT:PSS HTL that possesses interconnected conductive domains and enhanced electrical conductivity is crucial to PSCs.

What's the cathode?

Bulk heterojunction (BHJ) polymer solar cells (PSCs) that convert absorbed sunlight into electrical energy have been in the spotlight due to their potential in large-scale and cost-effective roll-to-roll fabrication. In the past decade, massive endeavors concerning the development and interfacial engineering of novel light-harvesting materials have been made to improve the power conversion efficiency (PCE). Nowadays, the highest PCE exceeds 16% for single-junction BHJ PSCs. Interface layers, positioned between the photoactive layer and the anode/cathode electrode, play vital roles in governing the performance of PSCs. Interface layers are usually utilized to tailor the work function of the electrode for charge carrier collection maximization, modify the interface to alter the photoactive layer micromorphology, as well as minimize carrier recombination at the interfaces between the active layer and the transport layer. For a hole transport layer (HTL), pivotal parameters consist of high conductivity, high transparency, solution processability, favorable stability, etc. PEDOT:PSS, as the state-of-the-art hole transport material in organic optoelectronic devices, possesses outstanding edges of high transmittance in the visible spectrum and superior film-forming ability and thermal stability. However, PEDOT:PSS solution is strongly acidic and hygroscopic, which can etch the indium tin oxide (ITO) anode. Moreover, it has been verified that the spin-casted PEDOT:PSS thin film often exhibits inhomogeneous morphology and poor electrical conductivity, leading to poor efficiency and stability of devices. The development of a PEDOT:PSS HTL that possesses interconnected conductive domains and enhanced electrical conductivity is crucial to PSCs.

What's the anode?

indium tin oxide (ITO)

Bulk heterojunction (BHJ) polymer solar cells (PSCs) that convert absorbed sunlight into electrical energy have been in the spotlight due to their potential in large-scale and cost-effective roll-to-roll fabrication. In the past decade, massive endeavors concerning the development and interfacial engineering of novel light-harvesting materials have been made to improve the power conversion efficiency (PCE). Nowadays, the highest PCE exceeds 16% for single-junction BHJ PSCs. Interface layers, positioned between the photoactive layer and the anode/cathode electrode, play vital roles in governing the performance of PSCs. Interface layers are usually utilized to tailor the work function of the electrode for charge carrier collection maximization, modify the interface to alter the photoactive layer micromorphology, as well as minimize carrier recombination at the interfaces between the active layer and the transport layer. For a hole transport layer (HTL), pivotal parameters consist of high conductivity, high transparency, solution processability, favorable stability, etc. PEDOT:PSS, as the state-of-the-art hole transport material in organic optoelectronic devices, possesses outstanding edges of high transmittance in the visible spectrum and superior film-forming ability and thermal stability. However, PEDOT:PSS solution is strongly acidic and hygroscopic, which can etch the indium tin oxide (ITO) anode. Moreover, it has been verified that the spin-casted PEDOT:PSS thin film often exhibits inhomogeneous morphology and poor electrical conductivity, leading to poor efficiency and stability of devices. The development of a PEDOT:PSS HTL that possesses interconnected conductive domains and enhanced electrical conductivity is crucial to PSCs.

What's the electrolyte?

LiMn2O4 (LMO) was used as a cathode to evaluate the electrochemical performance of Al2O3-coated LTO in the WiSE (Fig. S5†). The mass ratio of LMO:LTO was set as 2.5:1 to accommodate the low CE of LTO during the initial several cycles. 1C was used instead of a high rate to demonstrate the stability of the electrolyte in the full cell. Such LTO/LMO full cells delivered a voltage plateau at ∼2.4 V during discharge. The discharging capacity based on LTO mass was 145 mA h g−1. In the first cycle, a coulombic efficiency (CE) of 84.5% was delivered, indicating that a relatively small amount of electrolyte was consumed to form an additional LiF-rich SEI on the Al2O3-coated LTO anode. In comparison, pairing the uncoated LTO and LMO delivered a CE of 50% (Fig. S6†), further confirming the effect of the Al2O3 coating in suppressing the side reaction. As we reported previously, the reduction of salt anions bis(trifluoromethane sulfonyl)imide (TFSI) occurs between 1.9 V and 2.9 V. Although reduction of TFSI anions is expected to occur if an Al2O3 coating is absent (Fig. S7†), formation of a complete SEI needs a long time (i.e., few cycles in galvanostatic charge/discharge cycles). The lithiation potential of pristine LTO resides beyond the cathodic limit of the WiSE, so a significant HER will occur before lithiation. The persistent evolution of gas undoubtedly prevents complete formation of the SEI. For the initial cycles, when a robust SEI has not been constructed, protection of the Al2O3 surface serves as a key barrier to ensure that SEI chemistry occurs, and that the SEI ingredient formed from the reduction of the TFSI anion adheres to the anode surface. After the most challenging period in the initial cycles, a dense and complete SEI will come into shape (Fig. S7†), eventually providing long-term protection and allowing LTO to deliver a reversible capacity.

What's the cathode?

LiMn2O4 (LMO)

LiMn2O4 (LMO) was used as a cathode to evaluate the electrochemical performance of Al2O3-coated LTO in the WiSE (Fig. S5†). The mass ratio of LMO:LTO was set as 2.5:1 to accommodate the low CE of LTO during the initial several cycles. 1C was used instead of a high rate to demonstrate the stability of the electrolyte in the full cell. Such LTO/LMO full cells delivered a voltage plateau at ∼2.4 V during discharge. The discharging capacity based on LTO mass was 145 mA h g−1. In the first cycle, a coulombic efficiency (CE) of 84.5% was delivered, indicating that a relatively small amount of electrolyte was consumed to form an additional LiF-rich SEI on the Al2O3-coated LTO anode. In comparison, pairing the uncoated LTO and LMO delivered a CE of 50% (Fig. S6†), further confirming the effect of the Al2O3 coating in suppressing the side reaction. As we reported previously, the reduction of salt anions bis(trifluoromethane sulfonyl)imide (TFSI) occurs between 1.9 V and 2.9 V. Although reduction of TFSI anions is expected to occur if an Al2O3 coating is absent (Fig. S7†), formation of a complete SEI needs a long time (i.e., few cycles in galvanostatic charge/discharge cycles). The lithiation potential of pristine LTO resides beyond the cathodic limit of the WiSE, so a significant HER will occur before lithiation. The persistent evolution of gas undoubtedly prevents complete formation of the SEI. For the initial cycles, when a robust SEI has not been constructed, protection of the Al2O3 surface serves as a key barrier to ensure that SEI chemistry occurs, and that the SEI ingredient formed from the reduction of the TFSI anion adheres to the anode surface. After the most challenging period in the initial cycles, a dense and complete SEI will come into shape (Fig. S7†), eventually providing long-term protection and allowing LTO to deliver a reversible capacity.

What's the anode?

Al2O3-coated LTO

LiMn2O4 (LMO) was used as a cathode to evaluate the electrochemical performance of Al2O3-coated LTO in the WiSE (Fig. S5†). The mass ratio of LMO:LTO was set as 2.5:1 to accommodate the low CE of LTO during the initial several cycles. 1C was used instead of a high rate to demonstrate the stability of the electrolyte in the full cell. Such LTO/LMO full cells delivered a voltage plateau at ∼2.4 V during discharge. The discharging capacity based on LTO mass was 145 mA h g−1. In the first cycle, a coulombic efficiency (CE) of 84.5% was delivered, indicating that a relatively small amount of electrolyte was consumed to form an additional LiF-rich SEI on the Al2O3-coated LTO anode. In comparison, pairing the uncoated LTO and LMO delivered a CE of 50% (Fig. S6†), further confirming the effect of the Al2O3 coating in suppressing the side reaction. As we reported previously, the reduction of salt anions bis(trifluoromethane sulfonyl)imide (TFSI) occurs between 1.9 V and 2.9 V. Although reduction of TFSI anions is expected to occur if an Al2O3 coating is absent (Fig. S7†), formation of a complete SEI needs a long time (i.e., few cycles in galvanostatic charge/discharge cycles). The lithiation potential of pristine LTO resides beyond the cathodic limit of the WiSE, so a significant HER will occur before lithiation. The persistent evolution of gas undoubtedly prevents complete formation of the SEI. For the initial cycles, when a robust SEI has not been constructed, protection of the Al2O3 surface serves as a key barrier to ensure that SEI chemistry occurs, and that the SEI ingredient formed from the reduction of the TFSI anion adheres to the anode surface. After the most challenging period in the initial cycles, a dense and complete SEI will come into shape (Fig. S7†), eventually providing long-term protection and allowing LTO to deliver a reversible capacity.

What's the electrolyte?

3Mg/Mg2Sn electrodes were subjected to galvanostatic C/D at various rates in Mg(HMDS)2/MgCl2. In addition to the unprecedented high capacity, the rate performance of 3Mg/Mg2Sn was also impressive. The stepwise increase in C/D rates resulted in a continuous decrease in reversible capacities, but the degree of reduction was not abrupt (Fig. 5A). For example, a reversible capacity of 805 mA h g−1 was delivered at 100 mA g−1 and as much as 87 and 53% (700 and 430 mA h g−1) of it was retained at 500 and 1500 mA g−1, respectively. Accordingly, the corresponding C/D profiles also showed a relatively small increase in overpotentials with an increase in the C/D rates (Fig. 5B). When compared with the plateau voltage at 100 mA g−1, the additional overpotential required at 1500 mA g−1 was only ca. ±50 mV. It should be stressed here that such excellence has never been reported for either alloy- or intercalation-type electrodes in MIBs. Most anode materials show a capacity of less than 400 mA h g−1 at a C/D rate lower than 100 mA g−1, which significantly decreases with increasing C/D rates (Table 1). A few materials can deliver reasonable capacities at high rates, but no performance has been comparable to that of 3Mg/Mg2Sn.

What's the anode?

3Mg/Mg2Sn electrodes were subjected to galvanostatic C/D at various rates in Mg(HMDS)2/MgCl2. In addition to the unprecedented high capacity, the rate performance of 3Mg/Mg2Sn was also impressive. The stepwise increase in C/D rates resulted in a continuous decrease in reversible capacities, but the degree of reduction was not abrupt (Fig. 5A). For example, a reversible capacity of 805 mA h g−1 was delivered at 100 mA g−1 and as much as 87 and 53% (700 and 430 mA h g−1) of it was retained at 500 and 1500 mA g−1, respectively. Accordingly, the corresponding C/D profiles also showed a relatively small increase in overpotentials with an increase in the C/D rates (Fig. 5B). When compared with the plateau voltage at 100 mA g−1, the additional overpotential required at 1500 mA g−1 was only ca. ±50 mV. It should be stressed here that such excellence has never been reported for either alloy- or intercalation-type electrodes in MIBs. Most anode materials show a capacity of less than 400 mA h g−1 at a C/D rate lower than 100 mA g−1, which significantly decreases with increasing C/D rates (Table 1). A few materials can deliver reasonable capacities at high rates, but no performance has been comparable to that of 3Mg/Mg2Sn.

What's the electrolyte?

Electrochemical impedance spectroscopy (EIS) curves are presented in Fig. S8.† A larger slope, which is associated with the Warburg impedance (Ws), was observed in the low-frequency region of the curve of the Fe-intercalated ML Ti3C2Tx electrode. Such an occurrence is related to the migration of electrolyte ions within the electrode. It is inferred that the faster kinetics of EMIm+ were derived from the increased space caused by Fe pre-intercalation. Ex situ Ti 2p and Fe 2p XPS spectra are presented in Fig. S9;† no obvious shifts were observed. This finding indicates that the charge storage mechanism of the Fe pre-intercalated ML Ti3C2Tx anode is based on the intercalation/de-intercalation behavior of EMIm+ instead of a redox process. Therefore, the larger real interlayer space contributed to the enhanced electrochemical performance. The scheme in Fig. 5d shows the charge storage mechanism of this DIB. Upon charging, EMIm+ cations were inserted into Fe pre-intercalated ML Ti3C2Tx, whereas PF6− anions were intercalated into the graphite cathode. Charge–discharge profiles of the full DIB are presented in Fig. S10†. The DIB provided an energy density of 76 W h kg−1 at a power density of 360 W kg−1. The cycling stabilities and capacity retentions of the electrodes were examined via performing GCD tests at the same current density. After 50 cycles, the DIB retained 94% of its initial capacity (Fig. 5e), indicating its improved stability compared with a DIB using the ML Ti3C2Tx anode (71%). Ragone plots are presented in the inset of Fig. 5e. The full DIB displayed a good rate capacity, with energy densities of 76, 70, and 64 W h kg−1 at power densities of 360, 900, and 1800 W kg−1, respectively. However, when the power density was higher than 3600 W kg−1, the energy density was reduced. The dual-ion intercalation mechanism of this DIB provides a new route for multi-ion storage.

What's the cathode?

graphite

Electrochemical impedance spectroscopy (EIS) curves are presented in Fig. S8.† A larger slope, which is associated with the Warburg impedance (Ws), was observed in the low-frequency region of the curve of the Fe-intercalated ML Ti3C2Tx electrode. Such an occurrence is related to the migration of electrolyte ions within the electrode. It is inferred that the faster kinetics of EMIm+ were derived from the increased space caused by Fe pre-intercalation. Ex situ Ti 2p and Fe 2p XPS spectra are presented in Fig. S9;† no obvious shifts were observed. This finding indicates that the charge storage mechanism of the Fe pre-intercalated ML Ti3C2Tx anode is based on the intercalation/de-intercalation behavior of EMIm+ instead of a redox process. Therefore, the larger real interlayer space contributed to the enhanced electrochemical performance. The scheme in Fig. 5d shows the charge storage mechanism of this DIB. Upon charging, EMIm+ cations were inserted into Fe pre-intercalated ML Ti3C2Tx, whereas PF6− anions were intercalated into the graphite cathode. Charge–discharge profiles of the full DIB are presented in Fig. S10†. The DIB provided an energy density of 76 W h kg−1 at a power density of 360 W kg−1. The cycling stabilities and capacity retentions of the electrodes were examined via performing GCD tests at the same current density. After 50 cycles, the DIB retained 94% of its initial capacity (Fig. 5e), indicating its improved stability compared with a DIB using the ML Ti3C2Tx anode (71%). Ragone plots are presented in the inset of Fig. 5e. The full DIB displayed a good rate capacity, with energy densities of 76, 70, and 64 W h kg−1 at power densities of 360, 900, and 1800 W kg−1, respectively. However, when the power density was higher than 3600 W kg−1, the energy density was reduced. The dual-ion intercalation mechanism of this DIB provides a new route for multi-ion storage.

What's the anode?

ML Ti3C2Tx

Electrochemical impedance spectroscopy (EIS) curves are presented in Fig. S8.† A larger slope, which is associated with the Warburg impedance (Ws), was observed in the low-frequency region of the curve of the Fe-intercalated ML Ti3C2Tx electrode. Such an occurrence is related to the migration of electrolyte ions within the electrode. It is inferred that the faster kinetics of EMIm+ were derived from the increased space caused by Fe pre-intercalation. Ex situ Ti 2p and Fe 2p XPS spectra are presented in Fig. S9;† no obvious shifts were observed. This finding indicates that the charge storage mechanism of the Fe pre-intercalated ML Ti3C2Tx anode is based on the intercalation/de-intercalation behavior of EMIm+ instead of a redox process. Therefore, the larger real interlayer space contributed to the enhanced electrochemical performance. The scheme in Fig. 5d shows the charge storage mechanism of this DIB. Upon charging, EMIm+ cations were inserted into Fe pre-intercalated ML Ti3C2Tx, whereas PF6− anions were intercalated into the graphite cathode. Charge–discharge profiles of the full DIB are presented in Fig. S10†. The DIB provided an energy density of 76 W h kg−1 at a power density of 360 W kg−1. The cycling stabilities and capacity retentions of the electrodes were examined via performing GCD tests at the same current density. After 50 cycles, the DIB retained 94% of its initial capacity (Fig. 5e), indicating its improved stability compared with a DIB using the ML Ti3C2Tx anode (71%). Ragone plots are presented in the inset of Fig. 5e. The full DIB displayed a good rate capacity, with energy densities of 76, 70, and 64 W h kg−1 at power densities of 360, 900, and 1800 W kg−1, respectively. However, when the power density was higher than 3600 W kg−1, the energy density was reduced. The dual-ion intercalation mechanism of this DIB provides a new route for multi-ion storage.

What's the anode?

ML Ti3C2Tx

From the linear relationship of Ip and ν0.5 (Fig. S13†), (cathodic peak at 2.3 V), (cathodic peak at 1.9 V), and (anodic peak at 2.4 V) were obtained. The DLi+ values of the Co5.47Nx/S cathode were 219.9, 14.8, and 31.5 × 10−15 cm2 s−1 for peak A, B, and C, respectively. Importantly, the DLi+ of peak A of Co5.47Nx/S was up to four fold that of the Co5.47N/S cathode (48.6, 18.1, and 18.1 × 10−15 cm2 s−1 for peak A, B, and C, respectively), as plotted in Fig. 6d. The adsorption efficacy of cobalt nitride is attributed to its strong chemical interaction with LiPSs through both Co–S and N–Li bonds. The strong Li–N bonding in Co5.47N without nitrogen vacancies impedes direct electron transfer to the LiPSs and delays Li+ diffusion, resulting in sluggish reaction kinetics. Compared with Co5.47N, Co5.47Nx with nitrogen vacancies has a lower N content, which leads to decreased N–Li bonding and accelerated Li+ transportation. The cycling stability of Co5.47Nx and Co5.47N was evaluated at a current density of 0.5C, and the results are shown in Fig. S14.† After 200 cycles, the Co5.47Nx electrode showed a capacity retention of 85%, which was higher than the 76% for the Co5.47N electrode without nitrogen vacancies. These results suggest that Co5.47Nx facilitates Li+ transportation through its hierarchical structure and nitrogen vacancies, resulting in promising cycling and rate capacity.

What's the cathode?

Co5.47Nx/S

From the linear relationship of Ip and ν0.5 (Fig. S13†), (cathodic peak at 2.3 V), (cathodic peak at 1.9 V), and (anodic peak at 2.4 V) were obtained. The DLi+ values of the Co5.47Nx/S cathode were 219.9, 14.8, and 31.5 × 10−15 cm2 s−1 for peak A, B, and C, respectively. Importantly, the DLi+ of peak A of Co5.47Nx/S was up to four fold that of the Co5.47N/S cathode (48.6, 18.1, and 18.1 × 10−15 cm2 s−1 for peak A, B, and C, respectively), as plotted in Fig. 6d. The adsorption efficacy of cobalt nitride is attributed to its strong chemical interaction with LiPSs through both Co–S and N–Li bonds. The strong Li–N bonding in Co5.47N without nitrogen vacancies impedes direct electron transfer to the LiPSs and delays Li+ diffusion, resulting in sluggish reaction kinetics. Compared with Co5.47N, Co5.47Nx with nitrogen vacancies has a lower N content, which leads to decreased N–Li bonding and accelerated Li+ transportation. The cycling stability of Co5.47Nx and Co5.47N was evaluated at a current density of 0.5C, and the results are shown in Fig. S14.† After 200 cycles, the Co5.47Nx electrode showed a capacity retention of 85%, which was higher than the 76% for the Co5.47N electrode without nitrogen vacancies. These results suggest that Co5.47Nx facilitates Li+ transportation through its hierarchical structure and nitrogen vacancies, resulting in promising cycling and rate capacity.

What's the cathode?

Co5.47N/S

In this work, a mechanism is proposed to explain the phenomenon of columnar lithium metal deposition. An electrolyte additive, such as HF, is selectively reduced at high potential vs. Li/Li+ to form uniformly distributed crystalline LiF particles with preferred crystallographic texture which are then encased in an amorphous matrix of solvent reduction products, as evidenced by systematic electrochemical analysis and in situ X-ray surface scattering. The LiF quantity and distribution are directly tunable by choosing an appropriate additive concentration and electrochemical cycling rate. Interfaces between LiF and the amorphous phase act as fast lithium-ion diffusion pathways, promoting a thin and more uniform SEI which leads to a very high lithium metal nucleation density relative to additive-free systems followed by nearly-isotropic and eventually vertical columnar growth, observed in real-time using operando small angle X-ray scattering. Understanding this process step-by-step provides new insights into the role of electrolyte additives and provides new information for the rational design of such additives. As HF is damaging to cathodes, current collectors, and other cell components, new additives which decompose at high potentials vs. lithium and direct the formation of a columnar morphology should be developed for practical use in lithium metal batteries.

What's the cathode?

To summarize, multiple organic-based active materials showed promising characteristics in potassium-based batteries. Their capacities and potentials are typically less attractive than those of the state-of-the-art inorganic analogs, but the demonstrated rate and cycle capabilities make them highly competitive. There are still plenty of opportunities to improve the characteristics of the organic-based compounds using the truly unlimited diversity of molecular structures and tunability of their properties via functional group substitution. New advanced cathode and anode materials can be developed via rational molecular design, morphology optimization, engineering conductive fillers, and tuning the electrolyte composition.

What's the cathode?

To summarize, multiple organic-based active materials showed promising characteristics in potassium-based batteries. Their capacities and potentials are typically less attractive than those of the state-of-the-art inorganic analogs, but the demonstrated rate and cycle capabilities make them highly competitive. There are still plenty of opportunities to improve the characteristics of the organic-based compounds using the truly unlimited diversity of molecular structures and tunability of their properties via functional group substitution. New advanced cathode and anode materials can be developed via rational molecular design, morphology optimization, engineering conductive fillers, and tuning the electrolyte composition.

What's the anode?

To summarize, multiple organic-based active materials showed promising characteristics in potassium-based batteries. Their capacities and potentials are typically less attractive than those of the state-of-the-art inorganic analogs, but the demonstrated rate and cycle capabilities make them highly competitive. There are still plenty of opportunities to improve the characteristics of the organic-based compounds using the truly unlimited diversity of molecular structures and tunability of their properties via functional group substitution. New advanced cathode and anode materials can be developed via rational molecular design, morphology optimization, engineering conductive fillers, and tuning the electrolyte composition.

What's the electrolyte?

After structural comparison, the activity of oxygen and manganese was further probed to reveal the charge compensation in electrochemistry. Operando DEMS was performed to monitor the released anionic species during the charging process. The operando DEMS data was collected with active material loading of around 8 mg (Fig. 3). In the initial charge process for typical LMO, carbon dioxide gas was evolved at the beginning, accompanied by a large amount of oxygen gas release with maximum carbon dioxide evolution. The total gas evolution was summed to be 59.3 mmol CO2 per mol of active material and 70.2 mmol O2 per mol of active material. The CO2 gas may result from the electrolyte decomposition, the reaction between the electrolyte and O2, and impurity Li2CO3 decomposition. In contrast, the T-LMO electrode delivered a minimal amount of carbon dioxide and oxygen gas release. The cumulative CO2 and O2 detected from T-LMO during the initial charge process was 10.5 mmol and 9.1 mmol per mol of active material, respectively. The reduced gas release in the treated sample indicates the improved electrochemical stability of the T-LMO material when operating at high voltage in the battery. The different outgassing behaviors of LMO and T-LMO under identical cycling conditions imply that the two materials have different oxygen activities.

What's the electrolyte?

The normalized voltage characteristics for the 1st and 2nd cycles at C/20 (standard CC protocol) for Li[Li0.2Mn0.6Ni0.1Co0.1]O2 are presented in Fig. 2. During the 1st charge, we observed a rapid increase of the potential up to 4.4 V followed by a pseudo-plateau, while during the 1st discharge, the potential declines down to 3.5 V and a subsequent long pseudo-plateau was registered. The observed substantial hysteresis between charge and discharge potentials suggests: (i) sluggish kinetics of the occurring process and (ii) structural reorganization taking place. Comparing the 1st and 2nd cycles, we noticed that the 2nd charge characteristics do not show a large pseudo-plateau between 4.4 and 4.8 V but rather a monotonic S-type shape. Interestingly the potential of the 2nd discharge curve is slightly higher than that of the 1st discharge. In turn, the lower potential for the 2nd charge characteristics (in relation to the 1st one) leads to much lower hysteresis for the 2nd cycle compared with that for the initial one. This transition from stair-case to S-type potential characteristics between the 1st and 2nd cycles is commonly observed for other Li-rich NMC compounds; however the exact mechanism is still elusive and will be investigated further in this study. Comparing the discharge capacities at various current densities (Fig. 3a), the significant discharge capacity fade from approximately 350 to 290 mA h g−1 during the first two cycles is observed. Furthermore, the specific charge only slightly responds to the increase of the current loads and each rate increase (C/20, C/10, C/5, C/2, and 1C) is followed by a minor decrease of the practical specific charge until reaching 150 mA h g−1 at a 1C rate. Going back to the C/20 rate leads to an increase of the discharge capacity slightly below 250 mA h g−1. A similar observation was made in other reports within the literature where the first discharge capacity is above 300 mA h g−1 and after a few cycles it drops to ca. 250 mA h g−1 comparable to that in our paper. We suspect that most probably the anomalously high discharge capacity is due to the nanosize of primary particles which have a very high total surface area and as such enhance the electrolyte decomposition, increasing the observed specific charge. It is also worth noticing that after rate capability tests, during the second cyclic tests at C/20 the material works with a reversible capacity of 250 mA h g−1, suggesting most probably that the observed very high initial discharge capacity is due to the parasitic reactions. Voltage profiles for the second cycles registered at various current loads are shown in Fig. 3b and exhibit almost the same S-type character, confirming that reorganization taking place between the 1st and 2nd cycles persists in the following tests.

What's the electrolyte?

The electrochemical performance of LTO electrodes coated with various materials was evaluated by linear sweep voltammetry (Fig. 3). Hydrogen evolution began at ∼1.8 V vs. Li on the pristine LTO surface, which was higher than its lithiation potential (1.55 V). The HER process (rather than lithiation of LTO) dominated the cathodic reaction during the scan. The carbon coating enhanced the electronic conductivity of LTO, thus accelerating the HER (as evidenced by the higher currents and positively shifted HER potential). The TiO2 coating also positively shifted the cathodic limit due to its high catalytic activity. In contrast, the ZnO coating and Al2O3 costing negatively shifted the cathodic limit potential by 0.1 V. In addition, the HER currents on the ZnO-coated and Al2O3-coated LTO electrode were much lower compared with those on pristine LTO. The capability of HER suppression as quantified by the onset potential of the HER should increase in the order TiO2 < carbon < LTO < ZnO < Al2O3, which was consistent with the prediction in Fig. 1. The Al2O3 coating not only suppressed the HER catalytic activity but also acted as a kinetic barrier to slow down electron transfer from the electrode bulk to protons in the electrolyte. Al2O3 coating shifted the HER potential to <1.5 V, so Li+ intercalation was enabled before the HER, as evidenced by a sharp lithiation peak at 1.55 V (Fig. 4a). Using a similar approach, we also evaluated the effect of surface coating on the oxygen evolution reaction (OER). Al2O3 coating and TiO2 coating on LiNi0.5Mn1.5O4 reduced the side reactions on this high-voltage (4.8 V) cathode material only slightly (Fig. S3†).

What's the cathode?

LiNi0.5Mn1.5O4

The electrochemical performance of LTO electrodes coated with various materials was evaluated by linear sweep voltammetry (Fig. 3). Hydrogen evolution began at ∼1.8 V vs. Li on the pristine LTO surface, which was higher than its lithiation potential (1.55 V). The HER process (rather than lithiation of LTO) dominated the cathodic reaction during the scan. The carbon coating enhanced the electronic conductivity of LTO, thus accelerating the HER (as evidenced by the higher currents and positively shifted HER potential). The TiO2 coating also positively shifted the cathodic limit due to its high catalytic activity. In contrast, the ZnO coating and Al2O3 costing negatively shifted the cathodic limit potential by 0.1 V. In addition, the HER currents on the ZnO-coated and Al2O3-coated LTO electrode were much lower compared with those on pristine LTO. The capability of HER suppression as quantified by the onset potential of the HER should increase in the order TiO2 < carbon < LTO < ZnO < Al2O3, which was consistent with the prediction in Fig. 1. The Al2O3 coating not only suppressed the HER catalytic activity but also acted as a kinetic barrier to slow down electron transfer from the electrode bulk to protons in the electrolyte. Al2O3 coating shifted the HER potential to <1.5 V, so Li+ intercalation was enabled before the HER, as evidenced by a sharp lithiation peak at 1.55 V (Fig. 4a). Using a similar approach, we also evaluated the effect of surface coating on the oxygen evolution reaction (OER). Al2O3 coating and TiO2 coating on LiNi0.5Mn1.5O4 reduced the side reactions on this high-voltage (4.8 V) cathode material only slightly (Fig. S3†).

What's the electrolyte?

As ideal self-template precursor, metal–organic frameworks (MOFs) with regular compositions can be converted into porosity-tunable materials with a uniform distribution of metal-based components and carbon. The nano-sized metal-based components sustaining the framework of MOFs are able to curtail the distance of ion diffusion and create more electrolyte-accessible sites. Porous carbon, derived from decomposition of the organic linkers, can efficiently alleviate the volume variation of the metal-based components during the ion insertion/extraction process, and it can also reinforce electron transmission. In addition, the incorporation of heteroatoms can be easily realized by the modification of organic linkers in MOFs. Recently, some studies have reported that MOF-derived nanocomposites are promising anode materials for LIHCs and SIHCs, and they can facilitate boosting the pseudocapacitive-dominant charge storage as well as maintain the structure stability.

What's the anode?

MOF-derived nanocomposites

As ideal self-template precursor, metal–organic frameworks (MOFs) with regular compositions can be converted into porosity-tunable materials with a uniform distribution of metal-based components and carbon. The nano-sized metal-based components sustaining the framework of MOFs are able to curtail the distance of ion diffusion and create more electrolyte-accessible sites. Porous carbon, derived from decomposition of the organic linkers, can efficiently alleviate the volume variation of the metal-based components during the ion insertion/extraction process, and it can also reinforce electron transmission. In addition, the incorporation of heteroatoms can be easily realized by the modification of organic linkers in MOFs. Recently, some studies have reported that MOF-derived nanocomposites are promising anode materials for LIHCs and SIHCs, and they can facilitate boosting the pseudocapacitive-dominant charge storage as well as maintain the structure stability.

What's the electrolyte?

Based on the above discussion, a driving mechanism of the EC procedures in the shoulder-by-shoulder structure was proposed (Fig. 2d) which is similar to that of a planar-arrangement supercapacitor. When both WO3 films were transparent during the first process, in essence, the conductive ITO layers were likely acting as two split capacitance plates. Once a working voltage was applied, Li ions were driven by the edge effect of the electric field of the capacitance plates and intercalated into the WO3 film at the negative potential, namely PI. Because the edge effect of the electric field was intense in the middle of the two EC films, more Li ions were injected into this area than into other areas, showing an uneven color distribution. These concentration differences resulted in the self-diffusion of Li ions in the colored film, namely PIII. However, when one EC unit was colored in the subsequent processes, it was likely acting as a charged capacitor. Once a reverse driving voltage was employed, the charge transfer procedure (namely PII) preferentially occurred. Electrons were transferred through the outer loop, while Li ions were transferred through the electrolyte, and they simultaneously intercalated into another EC film that was at a negative potential and colored it. The ions (electrons and Li ions) in the area of the colored film that was next to the middle of the two EC units were first transported into its counterpart and placed in the area that was also next to the seam because the electric field in the seam area was very intense. It is reasonable that PII is much quicker than PI because discharging a capacitor is a rapid process. When PII finished, the EC film was partly colored. Next, PI and PIII both occurred, and the partly colored film became completely colored. When the working voltage was turned off, the colored film became evenly colored, driven by PIII.

What's the electrolyte?

To confirm the improved energy density, a hybrid supercapacitor was fabricated based on a solid-state 400-KOH-Ti3C2 anode and an active catholyte containing Mn2+ in 2 M H2SO4 electrolyte. Before assembling the hybrid supercapacitor, the CF as a current collector was pretreated by electrochemical predeposition with 3 mA h cm−2 MnO2 on the surface. The electrochemical performance of the as-assembled hybrid device was systematically investigated to demonstrate its unique advantages. As shown in Fig. 4a, in situ potential detection was performed, in which the Ag/AgCl reference electrode was introduced into the two-electrode system to explore in situ the potential variation of the respective electrode. This hybrid supercapacitor achieves a wide voltage window up to 1.7 V (red line), benefiting from the potential difference between the solid 400-KOH-Ti3C2 anode (dark cyan line) and active catholyte containing Mn2+ (black line). It is worth mentioning that such a wide operating voltage window is superior to that of most recently reported hybrid supercapacitors (see Table S2†) and even comparable to that of the “water in salt” electrolyte-based hybrid capacitors.Fig. 4b displays the CV curves of the hybrid supercapacitor at various scan rates. All the CV curves exhibit an approximate rectangular shape, which suggests that the operation mechanism in this hybrid device is bonding/debonding-induced pseudocapacitance in nature. Additionally, it is found that the addition of Mn2+ does not significantly affect the reaction kinetics of the 400-KOH-Ti3C2 anode (see Fig. S16† and the corresponding discussions). On the other hand, the near linear symmetric triangular charge/discharge curves at various current densities also confirm the pseudocapacitance characteristics of the hybrid device (Fig. 4c), which agrees well with the CV results. The corresponding specific capacitance was calculated based on the charge–discharge curves (see detailed calculations in the ESI†). Impressively, the hybrid device delivers an appreciable specific capacitance of 312.8 F g−1 at 1 A g−1 and 131.2 F g−1 at 80 A g−1, showing good rate performance (Fig. S17†). The energy and power densities of the present hybrid supercapacitor were also calculated on the basis of the total mass of 400-KOH-Ti3C2 and consumed MnO2, which is given as a Ragone plot profile (Fig. 4d). As displayed in Fig. 4d, the hybrid supercapacitor exhibits a maximum energy density of 43.4 W h kg−1 at a power density of 488.7 W kg−1. Even at a high power density of 40 kW kg−1, the energy density still remains 18.2 W h kg−1. It is worth noting that the energy density achieved here is markedly superior to those of most recently reported hybrid supercapacitors, as listed in Table S3.† These facts indicate that benefiting from both the high operating voltage and specific capacity of the active catholyte containing Mn2+, the energy density of the hybrid supercapacitor is significantly enhanced. In addition, this hybrid supercapacitor exhibits prominent cycling stability with a high-capacitance retention of 75% over 20000 cycles at 4 A g−1 (Fig. 4e). The charge/discharge curves at different cycles are shown in Fig. S18† to illustrate the high coulombic efficiency of the hybrid supercapacitor.

What's the anode?

400-KOH-Ti3C2

To confirm the improved energy density, a hybrid supercapacitor was fabricated based on a solid-state 400-KOH-Ti3C2 anode and an active catholyte containing Mn2+ in 2 M H2SO4 electrolyte. Before assembling the hybrid supercapacitor, the CF as a current collector was pretreated by electrochemical predeposition with 3 mA h cm−2 MnO2 on the surface. The electrochemical performance of the as-assembled hybrid device was systematically investigated to demonstrate its unique advantages. As shown in Fig. 4a, in situ potential detection was performed, in which the Ag/AgCl reference electrode was introduced into the two-electrode system to explore in situ the potential variation of the respective electrode. This hybrid supercapacitor achieves a wide voltage window up to 1.7 V (red line), benefiting from the potential difference between the solid 400-KOH-Ti3C2 anode (dark cyan line) and active catholyte containing Mn2+ (black line). It is worth mentioning that such a wide operating voltage window is superior to that of most recently reported hybrid supercapacitors (see Table S2†) and even comparable to that of the “water in salt” electrolyte-based hybrid capacitors.Fig. 4b displays the CV curves of the hybrid supercapacitor at various scan rates. All the CV curves exhibit an approximate rectangular shape, which suggests that the operation mechanism in this hybrid device is bonding/debonding-induced pseudocapacitance in nature. Additionally, it is found that the addition of Mn2+ does not significantly affect the reaction kinetics of the 400-KOH-Ti3C2 anode (see Fig. S16† and the corresponding discussions). On the other hand, the near linear symmetric triangular charge/discharge curves at various current densities also confirm the pseudocapacitance characteristics of the hybrid device (Fig. 4c), which agrees well with the CV results. The corresponding specific capacitance was calculated based on the charge–discharge curves (see detailed calculations in the ESI†). Impressively, the hybrid device delivers an appreciable specific capacitance of 312.8 F g−1 at 1 A g−1 and 131.2 F g−1 at 80 A g−1, showing good rate performance (Fig. S17†). The energy and power densities of the present hybrid supercapacitor were also calculated on the basis of the total mass of 400-KOH-Ti3C2 and consumed MnO2, which is given as a Ragone plot profile (Fig. 4d). As displayed in Fig. 4d, the hybrid supercapacitor exhibits a maximum energy density of 43.4 W h kg−1 at a power density of 488.7 W kg−1. Even at a high power density of 40 kW kg−1, the energy density still remains 18.2 W h kg−1. It is worth noting that the energy density achieved here is markedly superior to those of most recently reported hybrid supercapacitors, as listed in Table S3.† These facts indicate that benefiting from both the high operating voltage and specific capacity of the active catholyte containing Mn2+, the energy density of the hybrid supercapacitor is significantly enhanced. In addition, this hybrid supercapacitor exhibits prominent cycling stability with a high-capacitance retention of 75% over 20000 cycles at 4 A g−1 (Fig. 4e). The charge/discharge curves at different cycles are shown in Fig. S18† to illustrate the high coulombic efficiency of the hybrid supercapacitor.

What's the electrolyte?

Mn2+ in 2 M H2SO4