The increasing concern given to the global environment and energy sustainability is driving the research and development of electrochemical energy storage devices that provide power supply with more resilience and flexibility. Currently, lithium-ion batteries dominate the power-source market for portable devices and electric vehicles due to their high energy density, high energy efficiency, and long lifetime.1,2 However, the global maldistribution of lithium resources has impeded their further widespread use.3 In particular, flammable organic electrolytes in lithium-ion batteries result in a low safety level and high fabrication/maintenance costs, both of which are unacceptable for grid-scale use.4 Therefore, aqueous rechargeable batteries that contain safe and less expensive aqueous electrolytes are an important future alternative for sustainable development.5
For the development of aqueous batteries with high energy density, exploiting Zn metal as a negative electrode is a straightforward approach because the Zn metal electrode possesses high theoretical capacities of 820 mAh/g and 5854 mAh/L. While reversible Zn metal plating/stripping is an important issue to be addressed for the development of aqueous Zn-ion batteries6, another large challenge also remains in a positive electrode. The intercalation of large hydrated Zn2+ generally causes damaging structural changes upon charge/discharge, leading to capacity degradation after cycling.7 Consequently, the concept of aqueous dual-ion batteries has recently been increasingly studied. For example, an aqueous Zn2+/Li+ dual-ion battery, which consists of a Zn metal plating/stripping negative electrode and Li+ intercalation positive electrode (e.g., LiFePO4) with an aqueous dual-ion Zn2+/Li+ electrolyte, was reported to provide an energy density of approximately 95 Wh/kg with 90% capacity retention after 80 cycles8.
In this work, we focus on protons as charge carriers in aqueous dual-ion batteries (Figure 1). Proton is the smallest and lightest cation; thus, it can be easily (de)intercalate in various structures at a fast rate.9 Moreover, barrierless H+ hopping enables fast H+ transport in an electrolyte owing to the Grotthuss mechanism, where protons are transferred through the hydrogen bond network.10 In recent years, the Grotthuss topochemistry was extended to hydrate solid-state materials; for example, Prussian blue analogs exhibit fast H+ (de)intercalation with the assistance of structural water networks.11 Similarly, the intercalated water layers in transition metal carbide nanosheets (MXenes) facilitate H+ storage and fast H+ transfer.12,13 Without structural/confined water, quinone-based organic compounds, which store H+ on carbonyl groups, exhibit both long cycle lives and large capacities.14
Among various transition metal (TM) oxides (TM= Mn, V, W, Ti, Mo)15–24 that deliver large capacities upon protonation, orthorhombic MoO3 (α-MoO3) possesses a unique bilayered structure (Figure 1) that accommodates various cations, such as Li+25,26, Na+27, Ca2+28 and Mg2+29 in organic electrolytes. In aqueous electrolytes, the intercalation of bare Zn2+30,31 and Al3+32 has also been studied. However, the H+ intercalation behavior in α-MoO3 with various aqueous electrolytes remains controversial: some reports claim a bare H+ intercalation mechanism33–37 that is consistent with HxMoO3 bronze obtained through the spillover method38, while others report water and H+ cointercalation39,40. Meanwhile, α-MoO3 suffers severe dissolution in aqueous electrolytes; therefore, the capacity decays gradually upon cycling, making it difficult to study the detailed H+ intercalation mechanism. Although many attempts have been made to solve this issue, for example, the use of highly concentrated electrolytes40,41, gel-type or quasi-solid-state electrolytes with polymer additives31, and electrode surface coating with polymers or ceramics30,42, the detailed H+ intercalation mechanism in α-MoO3 has yet to be fully understood. Herein, we provide the full comprehension of H+ intercalation in α-MoO3 as a cathode material for aqueous Zn2+/H+ batteries. Fast H+ transfer in α-MoO3 through a solid-state anhydrous Grotthuss mechanism realizes aqueous batteries with both high power and high energy densities.
Electrochemical properties of MoO3 in aqueous Zn2+/H+ electrolytes
α-MoO3 was synthesized by a previously reported hydrothermal method,40 and the synchrotron X-ray diffraction pattern confirmed the successful synthesis of a pure α-MoO3 phase (Figure S1). To study the H+ intercalation mechanism in α-MoO3, ZnCl2 was selected as an electrolyte salt. In addition to its compatibility with a Zn anode, the high solubility of ZnCl2 enabled the formation of a superconcentrated liquid structure with limited amount of free water molecules, while its Brønsted acidity generated a low pH environment with a high H+ concentration. Therefore, the electrochemical properties of α-MoO3 were evaluated using three aqueous electrolytes: conventional Zn2+ electrolyte (3 mol kg‒1 ZnCl2/H2O), superconcentrated Zn2+ electrolyte (32 mol kg‒1 ZnCl2/H2O), and superconcentrated dual-ion (Zn2+/H+) electrolytes (32 mol kg‒1 ZnCl2 + 1 mol kg‒1 P2O5/H2O). Note that P2O5 in a dual-ion electrolyte generates H+ through hydrolysis (P2O5 + 3H2O → 2H3PO4). The Raman spectra for the superconcentrated electrolytes indicate that most water molecules are coordinated to Zn2+ (Figure S2).
Before testing the electrochemical properties of α-MoO3, we evaluated the negative electrode, namely, Zn stripping and plating on a Ti current collector using the three electrolytes (Figures S3 and 2a).6 The average coulombic efficiency in the aqueous Zn2+/H+ dual-ion electrolyte is 99.0% over 200 cycles (Figure S3), largely outperforming the superconcentrated Zn2+ electrolyte (95.0% over 200 cycles) and conventional Zn2+ electrolyte (82.2% over 100 cycles). The improved zinc reversibility after the addition of P2O5 may result from the formation of a Zn3(PO4)2-based solid electrolyte interphase (SEI) layer.43 This reversible Zn stripping and plating was used as the counter electrode in this work.
Figure 2b shows the cyclic voltammetry (CV) curves measured for α-MoO3 using the three electrolytes at a scan rate of 0.5 mV s-1. To suppress the hydrogen evolution and chlorine evolution reactions on the α-MoO3 electrode, the cutoff voltages were set at 0.45 and 1.3 V for the aqueous Zn2+/H+ dual-ion electrolyte (Figure 2a) and 0.25 and 1.1 V for the Zn2+ aqueous electrolytes. In the first CV cycle (Figure 2b inset), the α-MoO3 electrode shows identical asymmetric-shaped CV curves for all three electrolytes, in which there are four cathodic and two anodic current flows. However, in the subsequent cycles, the CV curves become symmetric, showing two pairs of redox peaks. Note that the redox potentials using the aqueous Zn2+/H+ dual-ion electrolyte are centered at approximately 0.6 and 0.95 V vs. Zn/Zn2+, which shift by +0.2 V from those using Zn2+ aqueous electrolytes presumably owing to the potential shift of the Zn/Zn2+ counter electrode and/or the change in the activity of H+. The shapes of the CV curves for all three electrolytes resemble those reported for H+ (de)intercalation in α-MoO3 using a 9.5 M H3PO4 aqueous electrolyte40, suggesting dominant H+ (de)intercalation even when using the Zn2+ aqueous electrolytes. Indeed, the X-ray fluorescence (XRF) elemental analysis of the electrodes after a cathodic scan shows no evident increase in the peak intensity of Zn compared to that of the pristine electrode (Figure. S4). Considering that the aqua Zn2+ complex is a Brønsted acid to generate H+, it is most likely that H+ (de)intercalation occurs even when using the aqueous Zn2+ electrolytes. Importantly, while 3 and 32 mol kg‒1 ZnCl2 aqueous electrolytes exhibit steep redox peak degradation upon repeated CV cycling, the 32 mol kg‒1 ZnCl2 + 1 mol kg‒1 P2O5 aqueous dual-ion electrolyte exhibits stable CV curves (Figure 2b). The improved cycle stability should be ascribed to the suppression of α-MoO3 dissolution and the formation of effective SEI.43 Indeed, the X-ray photoelectron spectroscopy (XPS) analysis of the Zn metal anode after cycling in the 3 m ZnCl2 electrolyte evidences the Mo deposition from the Mo ions dissolved in the electrolyte (Figure S6).
Figure 2c shows the charge/discharge curves of the α-MoO3 electrode with galvanostatic charging followed by 3 h of potentiostatic charging in the aqueous Zn2+/H+ dual-ion electrolyte. The α-MoO3 electrode delivers a large capacity of 465 mAh g-1 at a rate of 0.5 A g-1 during the first discharge, corresponding to 2.5 H+ intercalation per formula unit of MoO3 with an average voltage of approximately 0.9 V. Note that ‘discharge’ and ‘charge’ of the α-MoO3 electrode are defined as H+ intercalation (cathodic process) and deintercalation (anodic process), respectively. Although the galvanostatic charge at 0.5 A g-1 can extract only 1.5 H+, the remaining 1.0 H+ can be extracted when applying a constant voltage of 1.3 V for 3 h (Figures S7 and S8). The diffusion coefficient determined by the potentiostatic intermittent titration technique (PITT) shows a significant 4.9-fold deceleration in H+ diffusion during the deprotonation from H1.1MoO3 to MoO3, confirming the trapped nature of ~1.0 H+ in MoO3 (Figure S9).
Under galvanostatic charging (without a potentiostatic step), the α-MoO3 electrode in an aqueous Zn2+/H+ dual-ion electrolyte retains 98% of its initial capacity after 1000 cycles at a rate of 2 A g-1 (Figure 3a). Furthermore, 62% of the specific capacity at 1 A g-1 is available at the fast discharge rate of 16 A g‑1 (Figure 3b and Figure S10). These performance results indicate the stability of the MoO3 framework against (de)protonation as well as the fast proton diffusion therein HxMoO3 (1.0 ≤ x ≤ 2.5). In contrast, the α-MoO3 electrodes in the aqueous Zn2+ electrolytes have capacity retentions of only 24.5% and 63.8%, respectively (Figures S11 and S12). Moreover, both the capacity and cycling stability of the α-MoO3 electrode in the aqueous Zn2+/H+ dual-ion electrolyte outperform those reported previously for α-MoO3 electrodes using aqueous electrolytes, including quasi-solid-state Zn2+ electrolytes30,31 and concentrated acid electrolytes33,40.
MoO3 host-lattice response to proton intercalation
To clarify the structural evolution of the α-MoO3 electrode in the aqueous Zn2+/H+ dual-ion electrolyte, in situ X-ray diffraction (XRD) was performed during the 1st cycle (Figure 4a). The interlayer distance of α-MoO3 remains nearly constant (approximately 7.0 Å) during the entire protonation process, while it increases from 7.0 to 7.5 Å and then decreases to 7.0 Å during deprotonation. This asymmetric lattice response is consistent with the asymmetric CV and charge/discharge curves, which also highly resemble those reported in a 6 M H2SO4 electrolyte33 and in a 4.4 M H2SO4 electrolyte35. However, despite the asymmetric structural evolution, the α-MoO3 structure recovers to the pristine state after a constant voltage is applied. The structure of the fully protonated phase was clarified using ex situ synchrotron XRD and Rietveld refinement (Figure 4b and 4c). Although it is difficult to determine proton positions using X-rays, the MoO3 framework only exhibits a slight monoclinic distortion after protonation, which is consistent with a previous report on H1.68MoO344 and the in situ XRD results.
Importantly, no water cointercalation occurs in the protonated structure. When a small amount of water (one water or hydronium intercalation in 16 formula units of MoO3) is intercalated, the density functional theory (DFT) calculation predicts that the interlayer distance expands by 13.8% (H2O intercalation) and 17.2% (H3O+ intercalation) (Figure S13), which are considerably larger than those observed experimentally. Additionally, the interlayer distance after water cointercalation in a 1 M H2SO4 electrolyte has been reported to show an expansion of 11% upon protonation39, in contrast to the negligible change observed in the protonation process in our experiment. Therefore, bare H+ (de)intercalation occurs in the α-MoO3 electrode with the aqueous Zn2+/H+ electrolyte.
Solid-state anhydrous Grotthuss mechanism
The above experimental results indicate asymmetric bare-H+ (de)intercalation in the α-MoO3 electrode. To clarify the origin of this asymmetry, we conducted DFT calculations on the H+ dynamics in α-MoO3. The stable sites for H+ were determined by the structural optimization of the protonated phases. The three most stable H+ absorption sites are labeled as site A (O2-H···O2), site B (O1-H···O1/O2) and site C (O2-H···O1), where O1 and O2 are the terminating oxygen coordinated to one Mo and the bridging oxygen coordinated to two Mo, respectively (Figure S14). The remaining edge-sharing oxygen (O3) coordinated to three Mo atoms is not favorable for H+ adsorption because all O 2p orbitals (2px, 2py, and 2pz) participate in the Mo 4d-O 2p bonds. Site A is located within the MoO3 layer, while sites B and C are located in the MoO3 interlayer space.
The most favorable site for H+ changes as a function of x in HxMoO3. At low H+ concentrations during protonation (0 < x ≤ 0.5 in HxMoO3), intralayer site A is the most favorable for H+ adsorption. However, after 0.5 H+ intercalation (0.5 < x ≤ 2.5 in HxMoO3), interlayer sites B and then C become more favorable than site A, leading to the final formula of H2.5MoO3 (site A (0.5 H+) → site B (1.5 H+) → site C (0.5 H+), Figure 5a and Figure S15). Further protonation of site B (H3MoO3) is not thermodynamically favored, as it causes structural decomposition.
On the other hand, at the early stage of deprotonation (2.5 > x ≥ 2.25 in HxMoO3), the DFT calculations suggest that 0.25 H+ is extracted from interlayer site C. Then, for 2.25 > x ≥ 1.5 in HxMoO3, 0.5 H+ at site A is predominantly deintercalated in addition to the simultaneous partial deintercalation of H+ at site B. After completing H+ deintercalation from site A, the H+ at site B is deintercalated (1.5 > x ≥ 0 in HxMoO3). In parallel, the remaining 0.25 H+ at site C is deintercalated at the end of deprotonation. Notably, the deprotonation order (site C (0.25 H+) → site A (0.5 H+) → site B (0.75 H+) → site B + site C (1.0H+)) (Figure 5b) does not follow the protonation order (site A (0.5 H+) → site B (1.5 H+) → site C (0.5 H+)), which explains the asymmetric charge/discharge profile. Experimentally, complete deprotonation is possible only when a constant voltage is applied (Figure 2c). The sluggish deprotonation at the end of charging may arise from the slow kinetics of the remaining H+ at sites B and C. The calculated voltage profile agrees well with the experimental results, confirming the validity of the asymmetric (de)protonation processes (Figure S16).
To unveil the origin of the facile (de)protonation in the 1.0 < x ≤ 2.5 range of HxMoO3 and the sluggish deprotonation in the 1.0 ≥ x ≥ 0 range of HxMoO3, climbing image nudged elastic band (CI-NEB) calculations were conducted, and the energy profile of H+ diffusion in MoO3 was visualized (Figure 5c). At the beginning of protonation, H+ rotates and hops consecutively between site A with a low energy barrier of 0.13 eV. The dense, zigzagged O2 (bridging oxygen) array provides a 1D channel for fast proton transport within the MoO3 intralayer space, where the short O2-O2 distance (2.63 Å) facilitates proton hopping. Upon the further protonation of H0.5MoO3, the interlayer diffusion channel containing site B also has a low energy barrier of 0.26 eV with an O1-O1/O2 distance of 2.71 Å (Figure S17), confirming the high mobility of H+ at site B.
As shown in Figure 5d, the deprotonation from site C of H2.5MoO3 also involves a small energy barrier of 0.29 eV through the interlayer diffusion channel containing sites B and C (O-O distance of 2.86 Å). The H+ transfer in the dense oxide-ion array is described as the solid-state anhydrous Grotthuss mechanism, whose kinetics highly rely on the distance between two adjacent lattice oxide ions. Figure S18 summarizes the statistics of O-O distances in all possible deprotonation channels of the charged structures. As expected, during deprotonation from H2.5MoO3 to HMoO3, short O-O distances of < 2.9 Å exist to enable solid-state anhydrous Grotthuss H+ transfer. However, after extracting 0.25 H+ from site C in H2.5MoO3, the O-O distances in the diffusion channel consisting of sites B and C increase drastically (> 3.1 Å); meanwhile, the energy barrier therein for H+ diffusion becomes remarkably high (0.98 eV, Figure S19). Therefore, further deprotonation from site C is unfavorable, instead, protons in site A are deintercalated, resulting in asymmetric (de)protonation (Figure 5b). Indeed, after deprotonation to HMoO3, all the long-range 1D channels are disrupted so that fast Grotthuss H+ transfer is no longer applicable (Figure S20). The remaining H+ can be removed only under long relaxation times, such as potentiostatic charging and PITT (Figure S21 and S22). However, except for the trapped H+ at site C, HxMoO3 (1.0 ≤ x ≤ 2.5) exhibits fast H+ transport through the diffusion channels upon charge/discharge, providing a remarkably high capacity and high-rate capability, as demonstrated in Figure 3.
To summarize, coupled with a zinc metal anode and an aqueous dual-ion Zn2+/H+ (32 m ZnCl2 + 1 m P2O5) electrolyte, the MoO3-Zn battery delivers a large energy density of 413 Wh kg-1 upon discharge at a power density of 0.90 kW kg-1 as well as a peak power density of 10.52 kW kg-1 at an energy density of 217 Wh kg-1 per weight of MoO3; these results are more than double that of a similar MoO3-Zn battery with a ZnCl2-based electrolyte39 (energy density of 198 Wh kg-1 at a power density of 0.28 kW kg-1 and power density of 6.7 kW kg-1 at an energy density of 104.5 Wh kg-1). Moreover, with the aid of the solid-state anhydrous Grotthuss mechanism, this prototype cell successfully outperforms most aqueous zinc-ion batteries and proton batteries (Figure S23). Contrary to conventional intercalation chemistry, which requires a porous host that accommodates ion diffusion and storage, the solid-state anhydrous Grotthuss mechanism demonstrated in this work enables fast H+ transfer and accumulated H+ storage in dense oxide-ion arrays. Therefore, further exploration for other host materials capable of H+ intercalation based on the solid-state anhydrous Grotthuss mechanism will be an important challenge for not only fabrications of high-power aqueous H+ batteries but also other solid-state ionics applications using protons.