Copper ions-intercalated manganese dioxide self-supporting mesoporous carbon electrode for aqueous zinc-ion batteries†

Richeng Jin, Yuan Fang, Beibei Gao, Ying Wan, Yi Zhou, Guofeng Rui, Wei Sun*, Pengpeng Qiu* and Wei Luo*
State Key Laboratory for Modification of Chemical Fibers and Polymer Materials, College of Materials Science and Engineering, Institute of Functional Materials, Donghua University, Shanghai 201620, China. E-mail: weisun@dhu.edu.cn; qiupengpeng@dhu.edu.cn; wluo@dhu.edu.cn

First published on 31st May 2024

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1 Introduction

The excessive consumption of traditional fossil fuels has led to the development of reliable and low-cost electrochemical energy storage systems (EESs). The latter are used to attain a wide application of clean energies such as nuclear and geothermal energy.^1–7 Aqueous Zn-ion batteries (AZIB) with a suitable redox potential (−0.76 V vs. standard hydrogen electrode) and large specific capacity (820 mA h g⁻¹) are considered to be promising EESs due to their advantages of low cost, safety, and non-toxicity.^8–10 However, the structural phase transition of the cathode and dissolution of the active substance limit the whole performance of ZIB.^11,12 Therefore, the exploration of a suitable cathode material is particularly important for the commercialization of AZIB.

Manganese-based oxides, vanadium-based oxides, Prussian blue analogues, and organic compounds have been extensively studied.^13–17 Among them, manganese oxide has attracted tremendous interest because of its high specific capacity of ∼308 mA h g⁻¹, large operating voltage of ∼1.35 V, and low toxicity. Compared with the other phases of MnO₂ (α-, β-, γ-, ε-, λ-MnO₂), the layered structure of δ-MnO₂ has a large layer spacing (∼0.7 nm), which makes it theoretically more suitable for the intercalation of Zn ions.^18–20 However, δ-MnO₂ still has the problems of poor kinetics of ion diffusion and unstable structure, which lead to poor magnification performance and rapid capacity decay in practical applications. To address these problems, pre-intercalation,²¹ oxygen defects,²² structural design,²³ and construction of composite materials²⁴ are usually employed to boost the electrochemical properties of the cathode. In particular, the pre-intercalation strategy is of great importance because it can increase the conductivity of MnO₂ and expand the lattice spacing to improve the ion-transport dynamics. For example, Zhang et al. prepared a PVP-intercalated layered manganese dioxide by a hydrothermal method. It showed a high performance for Zn-ion storage by accelerating electron and interlayer ion transport.²⁵ Jing et al. reported a dual-ion intercalation strategy to incorporate Mg²⁺ and K⁺ into δ-MnO₂, which showed high stability during an electrochemical process.²¹ The above studies show that the pre-intercalation of different ions affects the contribution ratio of H⁺ and Zn²⁺. However, Zn²⁺ has a strong electrostatic interaction with the base material, resulting in slow kinetics, while H⁺ has faster reaction kinetics and weak interaction with the main structure. Hence, it is a practicable scheme to increase the capacity contribution of H⁺.

Herein, we prepared a Cu²⁺-intercalated MnO₂ self-supporting electrode as the cathode of AZIB. In order to improve the intrinsic low wettability of the original carbon cloth, mesoporous structures were constructed in situ on the carbon cloth to provide a ‘cradle’ for the growth of δ-MnO₂. Then, Cu²⁺-intercalated δ-MnO₂ was grown on the substrate by a hydrothermal method. The Cu²⁺ intercalation greatly expanded the lattice spacing and strengthened the diffusion kinetics of H⁺ while maintaining the original layered structure. Benefiting from the superior structural design, the Cu–MnO₂ cathode delivered a high specific capacity and excellent cycling stability.

2 Results and discussion

Fig. 1 shows the synthetic scheme of the Cu–MnO₂ cathode. The polyethylene oxide-b-styrene (PEO-b-PS) copolymer template was first prepared by an atom-transfer radical-polymerization method.²⁶ Gel permeation chromatography of PEO-b-PS (Fig. S1†) showed that the molecular weight was 44343.0 g mol⁻¹ and the polydispersity index was ∼1.0, indicating the high quality of PEO-b-PS. Then, mesoporous carbon layers were grown on both sides of the carbon cloths (CCs) through assembly of interfacial monomicelles, followed by a carbonization process in N₂. Finally, Cu ions-intercalated δ-MnO₂ was deposited on the surface of mesoporous carbon through a hydrothermal reaction.^27,28 As a comparison, δ-MnO₂ was prepared by a similar method without addition of the Cu precursor.

Field-emission scanning electron microscopy (FESEM) and high-resolution transmission electron microscopy (HRTEM) were used to observe the morphology and microstructure of samples. Compared with the original carbon cloth, mesoporous carbon structures were clearly deposited on the surface of the carbon cloth (Fig. S2a–d†). After hydrothermal treatment, Cu–MnO₂ nanosheets grew uniformly on the surface of the mesoporous carbon cloth with an ordered nanoarray morphology (Fig. 2a–c and S3a†). TEM images (Fig. 2d and e and S3b†) also show similar nanoarray structures with a width of 75 nm. Fig. 2e displays the high-resolution HRTEM of Cu–MnO₂, which clearly revealed a lamellar structure with a d-spacing of 0.75 nm which belongs to the (001) crystal plane of δ-MnO₂ (JCPDS no. 80-1098). The distinct diffraction rings in the SAED pattern showed that Cu²⁺ intercalation did not change the phase type of δ-MnO₂. However, compared with δ-MnO₂ (Fig. S3c†), the d-spacing of Cu–MnO₂ increased, which was caused by Cu²⁺ intercalation. Energy dispersive X-ray spectroscopy (EDX) (Fig. 2g) exhibited a uniform distribution of elements in Cu–MnO₂, indicating the average distribution of Cu²⁺ into δ-MnO₂.

The crystal structures of cathodes were examined by XRD. In the XRD patterns of δ-MnO₂, the diffraction peaks at 12°, 26.2°, and 37.5° correspond well to the (001), (002), and (−111) planes of δ-MnO₂ (JCPDS no. 80-1098), indicating the synthesis of layered manganese dioxide. Cu–MnO₂ showed a similar XRD pattern to that of δ-MnO₂, further indicating that the intercalation of Cu²⁺ did not break the phase structure of δ-MnO₂. Raman spectra of the cathodes (Fig. 3b) showed characteristic peaks at 500, 575 and 649 cm⁻¹ belonging to the Mn–O–Mn stretching vibration, Mn–O vibration along the skeleton chain, and the symmetric stretching vibration Mn–O bond, respectively.^29,30 These data suggested that these two samples had a typical crystal structure of birnessite-type. In addition, the difference in relative intensity of the Raman band between the two samples reflected the species variation between layered δ-MnO₂ layers.³¹ The N₂ sorption isotherms of Cu–MnO₂ exhibited a typical type-IV characteristic curve and a H3-type hysteresis loop, suggesting that Cu–MnO₂ had a mesoporous structure.^32,33 The Brunauer–Emmett–Teller (BET) surface area and pore size of Cu–MnO₂ were calculated to be 3.7 nm and 23.0 m² g⁻¹, respectively. In order to determine the source of the mesoporous structure of the electrode, nitrogen absorption and desorption tests were carried out on the treated carbon cloth. As shown in Fig. S2c,† the treated carbon cloth had a mesoporous structure, and its specific surface area and pore size were 3.096 m² g⁻¹ and 7.4 nm, respectively. This result showed that the mesoporous structure of the electrode was mainly derived from the accumulation of Cu–MnO₂ nanosheets rather than mesoporous carbon cloth. The elemental composition of Cu–MnO₂ was investigated by X-ray photoelectron spectroscopy (XPS). In the Mn 2p XPS spectrum of MnO₂ (Fig. 3d), dominant peaks of Mn 2p3/2 and Mn 2p1/2 were located at 642.4 and 654.3 eV, respectively. In the Mn 2p XPS spectrum of Cu–MnO₂, both Mn⁴⁺ and Mn³⁺ were present.³⁴ Among them, the amount of Mn⁴⁺ was 71.98% (Fig. S4b†) and the average chemical valence of Mn was +3.72. The valence state of Mn can be explained by the representative multiplet splitting caused by the exchange of electrons at its 3s–3d level.^35–37 As shown in Fig. 3e, the splitting energy (ΔE) of the two 3s peaks increased from 4.73 eV of δ-MnO₂ to 5.03 eV of Cu–MnO₂, indicating an increase in the amount of Mn³⁺ in order to balance the valence imbalance caused by Cu²⁺ intercalation. XPS of O 1s (Fig. 3f) showed lattice oxygen (O_lat) and O (O_Mn–OH) which belonged to Mn–OH.³⁴ We found that after pre-intercalation of Cu²⁺, O_lat was transferred in a more negative direction so that O_lat had more electrons, which was due to electron transfer between the embedded Cu²⁺ and lattice oxygen. The increase in the charge around the O atom greatly strengthened the adsorption capacity of Cu–MnO₂ for H⁺, improved the diffusion capacity of H⁺, and more H⁺ participated in the electrode reaction, thereby showing higher electrochemical performance. In Fig. S4,† characteristic peaks of Cu²⁺ were found at 934.3 eV and 954 eV, and belonged to Cu 2p 3/2 and Cu 2p 1/2, respectively. There were oscillating satellite peaks of Cu²⁺ at 942.7 eV and 962.4 eV, suggesting that Cu²⁺ had embedded within δ-MnO₂.

For research into the electrochemical properties of Cu–MnO₂, a coin cells configuration with a zinc anode, glass fiber film (Whatman, GF/D), and 2 M ZnSO₄ and 0.2 M MnSO₄ as the electrolyte were employed. The addition of MnSO₄ can inhibit the dissolution of Mn²⁺ during discharge processes.^38,39 The cyclic voltammetry (CV) curves of Cu–MnO₂ and δ-MnO₂ at 0.1 mV s⁻¹ are shown in Fig. 4a and S5a,† respectively. The two pairs of redox peaks of 1.56/1.38 V and 1.6/1.26 V corresponded to the deintercalation/intercalation of H⁺ and Zn²⁺, respectively. The almost overlapping CV curve after the first two cycles (Fig. 4a) illustrated the stable and reversible charge and discharge processes of the Cu–MnO₂ cathode material. At the same time, the larger CV area of Cu–MnO₂ compared with that of δ-MnO₂ also indicated that Cu–MnO₂ provided more capacity. The galvanostatic charge–discharge (GCD) voltage profiles of Cu–MnO₂ are shown in Fig. 4b, and display two charge–discharge platforms, which is consistent with CV curve. The reduction in capacity after the first cycle may have been due to the phase transition of Cu–MnO₂ caused by incomplete deintercalation/intercalation of Zn²⁺ in Cu–MnO₂. No obvious change of charge/discharge curves in subsequent cycles indicated that the Cu–MnO₂ electrode had excellent cyclic stability. The cycling performance of Cu–MnO₂ and δ-MnO₂ cathode under the rate of 1 A g⁻¹ is shown in Fig. 4c. The Cu–MnO₂ cathode stabilized at 300 mA h g⁻¹ after the initial capacity reduction, and was maintained at 280 mA h g⁻¹ after 400 cycles, with a capacity attenuation rate of 7%. By comparison, the δ-MnO₂ cathode exhibited poor cyclic performance, with 70% capacity fading after 400 cycles. It is worth noting that an increase in capacity was found in Fig. 4c. This was because as the cycle progresses, H⁺ is consumed in the electrolyte, and the pH gradually increases, triggering the deposition of additional Mn²⁺, resulting in an increase in capacity.³⁹ Fig. 4d shows the rate capability of Cu–MnO₂ and δ-MnO₂ cathodes measured at a rate of 0.3–5 A g⁻¹. As the current density increased, the specific capacity of electrodes decreased due to the limitation of electrochemical reaction kinetics. The Cu–MnO₂ cathode exhibited a discharge capacity of 617.85, 487.36, 315.14, 130.64, and 87.33 mA h g⁻¹ at a rate of 0.3, 0.5, 1, 3, and 5 A g⁻¹, respectively. After restoration to 0.3 A g⁻¹, the capacity of the Cu–MnO₂ cathode reached up to 506 mA h g⁻¹. In contrast, δ-MnO₂ exhibited much lower capacity than Cu–MnO₂ at the same rate. This finding suggested that the insertion of copper ions could modulate the electronic configuration and efficient ion/charge storage sites of δ-MnO₂ and, thus, improve the electron-transport and ion-diffusion capacities of the δ-MnO₂ cathode. The greater-than-theoretical-capacity of Cu–MnO₂ was due to the additional capacity provided by the 2 e⁻ deposition reaction of Mn²⁺ added in the electrolytic liquid.^39,40 To deeply research long-term cycling stability at high rates, Cu–MnO₂ and δ-MnO₂ cathodes were tested at a rate of 5 A g⁻¹. As shown in Fig. 4e, Cu–MnO₂ maintained a capacity of 114.23 mA h g⁻¹ after 1500 cycles of 5 A g⁻¹, which was much higher than that of δ-MnO₂ (80.4 mA h g⁻¹). At the same time, Cu–MnO₂ could maintain a stable coulombic efficiency of nearly 100% under long cycles while δ-MnO₂ showed an unstable coulombic efficiency. These results showed that Cu²⁺ could be used as the pillars in the interlayer of Cu–MnO₂ to obtain excellent electrochemical properties compared with δ-MnO₂.^21,41 For a more comprehensive study of the electrochemical properties of Cu–MnO₂, we tested its performance in different electrolytes at a current density of 1 A g⁻¹. It can be seen from Fig. S6† that Cu–MnO₂ exhibited a relatively stable long cycle performance in different electrolytes, but its specific capacity was different in different electrolytes. In an aqueous electrolyte without Mn²⁺, Cu–MnO₂ showed the highest electrochemical performance in 2 M ZnSO₄, but poor specific capacity in 3 M and 2 M Zn(CF₃SO₃)₂. This was because SO₄⁻ easily binds to H₂O and forms a water-rich layer around the cathode, which provides kinetic support for H⁺ embedding, while CF₃SO₃⁻ is hydrophobic and preferentially adsorbed at the cathode to create a low-water environment, which hinders H⁺ embedding and results in reduced capacity.⁴² In a non-aqueous point solution, DMSO electrolyte with 5% H₂O showed better electrochemical performance than that of pure DMSO. This was because the high electrostatic interaction within the ion pair (Zn²⁺/CF₃SO₃⁻) in DMSO and the high desolvation penalty at the electrode–electrolyte interface will reduce ion transport and increase the energy loss of the charge-storage reaction.⁴³ At the same time, the lack of H⁺ in DMSO stops it providing capacity contribution, so it shows poor specific capacity. The addition of H₂O causes MnO₂ to adsorb water during the discharge process and form superhydrated zinc pyroxene, which greatly promotes the diffusion of Zn²⁺ and improves the reactivity of the Zn-insertion reaction. In addition, we compared Cu–MnO₂ with other reported advanced ZIBs cathodes (Table S1†). Cu–MnO₂ showed better electrochemical properties than that of most cathodes.

In order to study how Cu²⁺ improves the electrochemical performance of δ-MnO₂, the ion diffusion kinetics and energy storage mechanism of the cathode were investigated by CV, galvanostatic intermittent titration technique (GITT) and electrochemical impedance spectroscopy (EIS). As shown in Fig. 5a, the CV curves of Cu–MnO₂ at different scanning rates were measured from 0.1 to 0.5 mV s⁻¹. Two pairs of redox peaks were detected during the scan, which corresponded to the reversible insertion–extraction of Zn²⁺/H⁺. Although polarization during the intercalation/deintercalation process led to a slight peak shift, the well-maintained shape of the CV curves indicated a reversible redox reaction of the Cu–MnO₂ cathode. The variation of peak current (i) with scan rate (ν) follows an empirical power-law, which can be expressed by the following formula:⁴⁴

i = aνb	(1)

where i denotes the current, ν reflects the scan rate, and a and b are variable parameters. The fitted value of b represents the type of energy storage, when b = 0.5 indicates the diffusion-control process and b = 1 suggests the capacitive-storage process. The b values of P1, P2, P3 and P4 of Cu–MnO₂ were 0.7, 0.58, 0.51 and 0.83, respectively (Fig. 5b). These findings suggest that the electrochemical process of the Cu–MnO₂ cathode was controlled by diffusion and capacitance processes, and was principally controlled by capacitance. The contribution rates of capacitance control and diffusion control can be calculated by the following equation:⁴⁴

i(ν) = k1ν + k2ν1/2	(2)

The response currents generated by capacitive and diffusion-controlled behaviour were calculated using k₁ν and k₂ν^1/2. As shown in Fig. 5c, at the scanning rates of 0.1, 0.2, 0.3, 0.4 and 0.5 mV s⁻¹, the capacitance contribution was 53.48%, 70.77%, 73.89%, 79.26% and 86.34%, respectively, indicating that capacitor control was dominant in Cu–MnO₂. However, the δ-MnO₂ cathode exhibited a lower capacitance ratio, as shown in Fig. S6,† indicating that Cu²⁺ could introduce more capacitance to promote the electrochemical performance of δ-MnO₂.

The ion diffusion coefficient was measured by constant current intermittent titration (GITT) and calculated according to the following formula:⁴⁵

	(3)

τ: constant current pulse duration

m_B: mass loading of the reactant

M_B: molecular volume

V_m: molar volume

S: electrode–electrolyte interface area

ΔE_s: change in steady-state voltage

ΔE_t: change in discharge/charge voltage

Fig. 5d and S7† are the GITT curves of Cu–MnO₂ and δ-MnO₂ cathodes. The appearance of the two platforms was caused by the sequential intercalation of H⁺ and Zn²⁺.^46,47 The ion diffusion coefficient (D) is shown in Fig. 5e. The diffusion coefficient before 50% was greater than the diffusion coefficient after 50%, which corresponded to the intercalation process of H⁺ and Zn²⁺, respectively. This observation indicated that the diffusion kinetics of H⁺ and Zn²⁺ were different in the electrochemical process. Compared with δ-MnO₂, the H⁺ diffusion coefficient of the Cu–MnO₂ cathode was one order of magnitude higher, indicating that the pre-intercalation of Cu²⁺ accelerated the diffusion rate of H⁺, improved the electrochemical performance and enhanced the capacity of Cu–MnO₂. Fig. 5f shows the Nyquist plot of Cu–MnO₂ and δ-MnO₂ cathodes. The Rct (charge transfer impedance) of Cu–MnO₂ and δ-MnO₂ obtained by circuit fitting were 57.92 (Ω) and 505.04 (Ω), respectively, which further confirmed the improvement of the kinetics performance of δ-MnO₂ by Cu²⁺.

To further investigate the ion-transfer behaviour and energy-storage mechanism of the Cu–MnO₂ cathode in AZIB, the crystal configuration, valence state and micromorphology of the cathode were studied by ex situ XRD, XPS and SEM. Fig. 6a and b show the XRD patterns of Cu–MnO₂ in different states during one cycle. When fully discharged (0.8 V), there were characteristic peaks at 34.6° and 58.69° which were attributed to Zn₂MnO₄ (JCPDS no. 19-1459). Diffraction peaks were observed at 28.14° and 32.8°, which corresponded to Zn₄SO₄(OH)₆·H₂O (ZSH) (JCPDS no. 39-0690), while the characteristic peak of MnOOH (JCPDS no. 24-0713) was observed at 21.13°. The formation of MnOOH and ZSH during discharge indicated that H⁺ intercalated into δ-MnO₂. With more H⁺ being consumed, a large amount of OH⁻ was produced in the electrolyte, and excess OH⁻, Zn²⁺ and SO₄²⁻ reacted to obtain ZSH. The formation of Zn₂MnO₄ was due to the insertion of Zn²⁺ into the interlayer of δ-MnO₂. The shift of the characteristic peaks of δ-MnO₂ from 37.39° and 12.52° to 35.86° and 10.93° confirmed the behaviour of ion intercalation. In the process of charging, the characteristic peaks of MnOOH and ZSH gradually disappeared, which suggested the good cyclic reversibility of H⁺. The characteristic peaks of Zn₂MnO₄ could still be observed when the charge was up to 1.8 V, indicating the poor reaction kinetics of Zn²⁺. The process stated above suggested the reaction mechanism of H⁺ and Zn²⁺ intercalation. The ex situ SEM results (Fig. 6c and f) showed that ZSH nanosheets were formed when discharged to 0.8 V, and disappeared when charged to 1.8 V, revealing the good reversibility of the microstructure. Fig. 6d and e show the ex situ XPS spectra of Mn 2p and O 1s at full charge (1.8 V) and complete discharge (0.8 V). When fully discharged (0.8 V), the peak area of Mn³⁺ increased significantly, further demonstrating the embedding of Zn²⁺ and H⁺, and leading to the reduction of Mn⁴⁺ to Mn³⁺. When fully charged (1.8 V), the content of Mn⁴⁺ increased, indicating that the deintercalation of ions re-oxidized Mn³⁺ to Mn⁴⁺. The XPS spectrum of Zn 2p is shown in Fig. S8.† The peak of Zn 2p decreased when charging from 0.8 to 1.8 V, which indicated the deintercalation behaviour of Zn²⁺. The observation of the peak of Zn at 1.8 V suggested that the Zn²⁺ mentioned above was not completely deintercalated, indicating the low diffusion kinetics of Zn²⁺. In addition, Fig. 6d reveals the change of O valence state during charge and discharge. When discharging to 0.8 V, the peak area of O_H2O was increased significantly due to the formation of ZSH, while the peak area of O_lat decreased due to the intercalation of ions. The opposite process during charging indicated that H⁺ and Zn²⁺ left the lattice of Cu–MnO₂ and the large chunks of ZSH disappeared, which showed that Cu–MnO₂ had good reversibility. The above results showed that the Cu–MnO₂ cathode exhibited good cyclic reversibility, and it was a feasible strategy to improve the embedding and removal ability of H⁺ by intercalation of Cu²⁺. In addition, we studied the structural changes during the long-term cycling of the battery. As shown in Fig. S10a,† an obvious characteristic peak of ZSH appeared in the XRD curve of the electrode after 1500 cycles, indicating that irreversible phase transition occurred during the long cycle. As mentioned above, H⁺ in the electrolyte was consumed during the electrochemical cycle, which led to an increase in pH and triggered ZSH deposition. The deposition of a large amount of ZSH would hinder the contact between Cu–MnO₂ and the electrolyte, resulting in a decrease in capacity. Fig. S10b† shows the SEM image of the electrode after 1500 C, and many ZSH patches in the figure confirm this interpretation.

3 Conclusions

Cu²⁺-intercalated δ-MnO₂ (Cu–MnO₂) was grown in situ on a mesoporous CCs substrate by a hydrothermal method as the cathode of AZIB, and its electrochemical properties were investigated. The designed Cu–MnO₂ cathodes showed excellent specific capacity, rate performance and cycle stability in AZIB. Cu–MnO₂ exhibited a reversible capacity of 280 mA h g⁻¹ at the rate of 1 A g⁻¹, maintained 93% capacity after 400 cycles, and had a high specific capacity of 114.23 mA h g⁻¹ at 5 A g⁻¹ after 1500 cycles. The combined use of CV, GITT, EIS, ex situ XRD and XPS showed that the intercalated Cu²⁺ formed a strong ionic bond with the oxygen atom, thus improving the stability of the layered structure and enhancing the H⁺ transfer kinetics.

4 Experimental section

4.1 Chemicals and materials

PEO-5000, 2-bromoisobutyryl bromide, and N,N,N′,N′,N′-pentamethyldiethylenetriamine (PMDETA) were purchased from Aladdin. Tetrahydrofuran (THF) (>99%), pyridine (>99%), cuprous bromide (CuBr), ammonium hydroxide (NH₄OH, 28–30 wt%), dopamine hydrochloride, and copper(II) chloride dihydrate (CuCl₂·2H₂O) were obtained from Greagent. Potassium permanganate (KMnO₄) and ethanol were sourced from Sinopharm Group Chemical Reagents. All reagents were of analytical grade and did not require further purification. Deionized water was used for all experiments.

4.2 Synthesis of Cu–MnO₂ and δ-MnO₂

Cu–MnO₂ and δ-MnO₂ was prepared by a one-step hydrothermal method. KMnO₄ and CuCl₂·2H₂O at a molar ratio of 10:1 were dissolved in 10 ml of water under vigorous stirring for 10 min. One piece of carbon cloth (2 × 2) was placed into the above solution, transferred into a Teflon™-lined stainless-steel autoclave and heated at 180 °C for 2 h. After cooling down to room temperature, the obtained sample was washed with DI water thrice, and dried using a vacuum oven at 60 °C for 12 h. δ-MnO₂ was synthesized using a similar procedure without CuCl₂·2H₂O.

4.3 Materials characterization

The crystal structure of the sample was characterized by X-ray diffraction (XRD) using a D/Max-2550 PC diffractometer (Rigaku). The microstructure of electrodes was characterized by scanning electron microscopy (SEM) using the MAIA3 system (TESCAN) and transmission electron microscopy (TEM) employing a 2100F system (JEM). HAADF-STEM and EDS elemental mapping were done using a Talos F200S setup operating at 200 kV. Raman spectra were acquired using a reflex Raman microscope with laser excitation at 532 nm. N₂ sorption was carried out with an Autosorb-iQ system (Quantachrome) to measure the Brunauer–Emmett–Teller (BET) surface area and pore-size distribution. The BET method was employed for the determination of the specific surface area of catalysts. The Barrett–Joyner–Halenda (BJH) model was employed to analyze the pore-size distribution and pore volume. Electrode surface valence and chemical environment were investigated using an 250Xi setup (Escalab) equipped with Al Kα radiation.

4.4 Electrochemical measurements

Cu–MnO₂ was prepared on carbon cloth by in situ growth and used as the cathode (area: 0.25 cm²) without conductive additives or adhesives. The mass loading of electrodes was 1.5 mg cm⁻². CR2032 coin cells were assembled in air. The configuration was with a zinc anode (diameter: 14 mm), glass fiber film (GF/D; Whatman), and 2 M ZnSO₄ and 0.2 M MnSO₄ as the electrolyte to estimate the properties of cathodic active materials. The galvanostatic discharge/charge (GCD) curve and GITT curve were created on the NEWARE battery test system, and the test voltage was 0.8–1.8 V. The specific capacity of the battery was calculated using this formula:

where Q_s represents the specific capacity (mA h g⁻¹), I represents the current (mA), m denotes the mass of the active substance (g) and Δt represents the discharge time.

The discharge GITT profiles were obtained at 0.1 A g⁻¹ for 1200 s followed with a rest for 1 h. The CV curve and EIS curve were measured by an electrochemical workstation (CHI600E). The frequency range of EIS was 100 kHz to 0.01 Hz. All tests were performed at room temperature.

Author contributions

Conceptualization: P. Q., R. J. and Y. F. Methodology: P. Q., R. J., W. S. and Y. F. Validation: P. Q. and R. J. Data analysis, R. J. Investigation: R. J. Resources, P. Q. and W. L. Data curation: R. J., Y. Z., G. R., Y. W. and B. G. Writing (preparation of original draft): R. J., W. S. and P. Q. Writing (review and editing): P. Q., W. S. and R. J. Visualization: R. J. Supervision: P. Q. and W. L. Project administration: P. Q. and W. L. Funding acquisition: P. Q. and W. L.

Conflicts of interest

There are no conflicts of interest to declare.

Acknowledgements

This work was supported by the National Natural Science Foundation of China (52225204, 52173233, 52002059 and 52202085), Innovation Program of Shanghai Municipal Education Commission (2021-01-07-00-03-E00109), Natural Science Foundation of Shanghai (23ZR1479200), “Shuguang Program” supported by the Shanghai Education Development Foundation and Shanghai Municipal Education Commission (20SG33), DHU Distinguished Young Professor Program (LZA2022001 and LZB2023002) and the Fundamental Research Funds for Central Universities (2232023G-07, 2232024Y-01).

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Footnote

† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d4im00042k

This journal is © Institute of Process Engineering of CAS 2024