A Self-Templating Redox-Mediated Synthesis of Hollow Phosphated

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A Self-Templating Redox-Mediated Synthesis of Hollow Phosphated Manganese Oxide Nanospheres as Noble-Metal-like Oxygen Electrocatalysts Tianran Zhang, Shengliang Zhang, Sheng Cao, Qiaofeng Yao, and Jim Yang Lee Chem. Mater., Just Accepted Manuscript • DOI: 10.1021/acs.chemmater.8b03681 • Publication Date (Web): 31 Oct 2018 Downloaded from http://pubs.acs.org on November 1, 2018

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Chemistry of Materials

A Self-Templating Redox-Mediated Synthesis of Hollow Phosphated Manganese Oxide Nanospheres as Noble-Metallike Oxygen Electrocatalysts Tianran Zhang, †, ‡ Shengliang Zhang, †, ‡ Sheng Cao, †, ‡ Qiaofeng Yao † and Jim Yang Lee *,†, ‡ † Department of Chemical and Biomolecular Engineering, National University of Singapore, 4 Engineering Drive 4, Singapore, 117576, Singapore. ‡ Cambridge Centre for Advanced Research and Education in Singapore, 1 Create Way, Singapore 138602, Singapore ABSTRACT: The development of low-cost high-performance electrocatalysts for the oxygen reduction reaction (ORR) and the oxygen evolution reaction (OER) to replace the use of noble metals is a challenge for the rechargeable metal-air batteries. Although manganese compounds, oxides in particular, have drawn the most interest but they rarely deliver the same performance as the noble metals in oxygen electrocatalysis; notwithstanding the enhancements introduced by nanosizing and adjuvant metal doping. Herein, we report a noble metal-like performance for manganese oxide catalysts by combining phosphate-modification with a hollow nanostructure. A simple and scalable self-templating method based on phosphate-mediated Mn redox reactions was developed for the preparation of hollow phosphated manganese oxide nanospheres at room temperature. A series of simple (h-MnOxPy) and complex phosphated manganese oxide (h-MeMnOxPy, Me = Co, Ni, Cu) hollow nanospheres can be produced more efficiently than normal hollow nanostructure construction techniques based on hard and soft templates and hydrothermal Ostwald ripening. Among the hollow phosphated manganese oxides h-MnOxP0.21 delivered the best ORR performance (half-wave potential of 0.85 V vs. RHE, similar to 20wt% Pt/C) and h-Co-MnOxP0.21 the best OER performance (1.60 V vs. RHE for 10 mA cm-2, marginally higher than 20wt% Ir/C). Small charge-discharge voltage gaps (ΔV) were shown in both alkaline (ΔV = 0.72 V at 5 mA cm-2) and neutral (ΔV = 1.28 V at 1 mA cm-2) rechargeable Zn-air batteries with the combined use of these catalysts, similar to the 20wt% Pt/C+ 20wt% Ir/C combined catalytic systems.

Introduction Low material cost and high energy density are the benefits of rechargeable metal-air batteries as a next generation power source.1-2 Among them the aqueous Zn-air batteries (ZnABs) with a high theoretical energy density of 1086 Wh kg-1, elemental abundance of zinc, and good battery safety, are one of the most studied variants.3-6 However, ZnABs presently are still fraught with issues of low energy efficiency and limited rechargeability due primarily to the sluggish kinetics of the oxygen reduction reaction (ORR) and oxygen evolution reaction (OER) at the air electrode. Effective electrocatalysis is the typical recourse. While noble metals constituted by far the largest family of ORR and OER active catalysts; their scarcity and prohibitively high cost naturally dismiss their use on any significant scale.7-9 The search for low-cost alternatives to the noble metals with sufficient performance is therefore a priority research area in the development of ZnABs.10-14 Among the alternatives developed to date, manganese compounds; and oxides in particular, have drawn the most interest because of the elemental abundance, low cost and environmental benignity of manganese.15-18 However, their low electrical conductivity and low native affinity for O2 and H2O have led to large overpotentials in the oxygen electrochemical reactions, and a notable performance gap in comparison with the noble metal catalysts.19-21

Previous research has identified increasing the surface area and tweaking the native activity of manganese oxide ORR and OER catalysts as possible remedies for their lackluster performance.22-24 In this regard, the fabrication of these catalysts as modified manganese oxides with a hollow nanostructure can be an effective solution.10, 25-28 For example, the assembly of nanosized building blocks with elevated catalytic activity into a mechanically strong hollow nanostructure can increase the electrolyte contact with the active sites and reduce the diffusional resistance of the active species. There are however very few demonstrations of hollow functionalized manganese oxide ORR and OER nanocatalysts to date; and rarely does any of them perform as well as the noble-metal catalysts.29-30 The complex synthesis process is another major drawback. Many of the published methods were based on templated synthesis (using either a hard,30-32 sacrificial,27, 33-34 or soft template35-36) and hydrothermal Ostwald ripening of manganese oxides.37-39 These processes are often harsh and laborious; and not proficient for generating manganese nanocomposites with more complex structures or compositions for activity tuning and optimization. More importantly, the oxygen electrocatalytic activities of the nanocatalysts synthesized to date as such are only moderate. This inevitably begets the question on whether there are other more effective methods for the preparation of manganese oxide nanocatalysts.

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Scheme 1. The sequence of processes in the preparation of phosphated manganese oxide hollow nanospheres.

We present here a new strategy effective for the fabrication of manganese oxides as phosphated hollow nanospheres. The methodology is exceedingly simple and involves only of adding a permanganate compound to a mixture of manganese and dihydrogen phosphate salts at room temperature (Scheme 1). The hollow nanosphere formation process is best understood in terms of an in-situ self-templating process mediated by phosphate-assisted Mn redox reactions. The strategy is effective for the synthesis of both simple and composite phosphated manganese oxide (denoted as h-MnOxPy and h-MeMnOxPy, where x is the oxygen non-stoichiometry, y is the experimentally measured P/Mn ratio; and Me is Co, Ni, or Cu) hollow nanospheres. These phosphated manganese oxide nanocatalysts are endowed with large surface areas and an abundance of the phosphate groups; and are able to surpass the ORR and OER performance of most of the manganese oxide nanocatalysts reported to date. In particular, the Mn (III) surface-enriched h-MnOxP0.21 catalyst showed the highest ORR activity in alkaline solution; which is comparable to the performance of a benchmarking 20wt% Pt/C catalyst. The h-Me-MnOxPy catalysts are more OER-active. The h-Co-MnOxP0.21 catalyst, in particular, performed as well as a benchmarking 20wt% Ir/C catalyst for OER in alkaline solution. The phosphated manganese oxide hollow nanospheres with the best ORR and OER performance were then used in combination in alkaline rechargeable Zn-air full cell batteries. Higher activity and greater operational stability were shown in comparison with Znair full cell batteries using the Pt/C-Ir/C catalyst combination. More interestingly rechargeable ZnABs using the combined phosphated manganese oxide catalysts also outperformed the noble metals catalyzed Zn-air batteries in a neutral electrolyte.

Experimental Section Simple phosphated manganese oxide hollow nanospheres (h-MnOxP0.21). In a typical synthesis, 200 mg MnSO4 H2O (>99%, Alfa-Aldrich) and 125 mg NH4H2PO4 (>98%, Alfa-Aldrich) were dissolved in 15 ml water. 15 ml of KMnO4 (99%, AMRESCO) aqueous solution (10 mg/ml) was added slowly to the MnSO4/NH4H2PO4 mixture under vigorous stirring. Stirring was continued for 2 h after the end of addition. The solid product was recovered by centrifugation, washed with water and dried in 60 oC air overnight. h-MnOxPy with different P contents

were prepared by varying the amounts of NH4H2PO4 used (50 mg, 100 mg and 200 mg). Composite phosphated manganese oxide hollow nanospheres (h-Me-MnOxP0.21, Me= Co, Ni and Cu). For the preparation of h-Me-MnOxP0.21, freshly-made hMnOxP0.21 was re-dispersed in 30 ml of water. 5 ml of 1.18 M metal sulfate (CoSO4·7H2O, NiSO4· 6H2O and CuSO4·5H2O, Alfa-Aldrich) solution was mixed with 25% ammonia solution (0.532 ml for CoSO4 7H2O and NiSO4 6H2O and 0.355ml for CuSO4 5H2O). The mixture containing the in-situ formed Me hydroxide was then added dropwise to the h-MnOxP0.21 dispersion and stirred overnight. The solid product was recovered by centrifugation, washed with water and dried in 60 oC air overnight. Material characterizations. Powder X-ray diffraction (XRD) measurements in the 5o-80o range were conducted on a Bruker D8 Advance diffractometer using a Cu Kα source (λ= 1.5418). Field emission scanning electron microscopy (FESEM) was performed on a JEOL JSM-6700F microscope. Field emission transmission electron microscopy (FETEM) images were taken by a 200 kV JEOL 2010F microscope with an Energy Dispersive Spectrometer (EDS) attachment. X-ray photoelectron spectroscopy (XPS) was performed on a Kratos Axis Ultra DLD spectrometer. The measured binding energies were corrected by referencing the C1s peak of adventitious carbon to 284.5 eV. The Mn valence state was inferred from the multiplet splitting in the Mn 3s spectrum.40 Fourier Transform Infrared Spectroscopy (FT-IR) was carried out on a Bio-Rad FTS-3500 ARX FTIR Spectrometer. N2 adsorption/desorption isotherms were measured by a Quantachrome Instruments Autosorb-iQ surface area and pore size analyzer. Before the measurement, a sample was degassed in vacuum at 150°C for 6 h. Pore size distribution data were calculated from the N2 sorption isotherms using the quenched solid density functional theory (QSDFT) model provided by the Autosorb-iQ software package (assuming slit/cylinder geometry). Electrochemical measurements. Electrochemical measurements were based on a three-electrode system consisting of a Ag/AgCl (in 3 M KCl, aq) reference electrode, a Pt foil counter electrode (No Pt signal was detected in the aged h-MnOxP0.21 and h-Co-MnOxP0.21 catalysts after the ORR and OER stability tests (Figure S16). The use of the Pt counter electrode had therefore not


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interfered with the measurements); and an electrolyte of aqueous 0.1 M KOH solution saturated with high-purity O2. The measured potentials were converted to the reversible hydrogen electrode (RHE) scale (ERHE =EAg/AgCl +0.21+0.059×pH V). The working electrode was fabricated by dispensing a catalyst ink onto a glassy carbon (GC, 5 mm in diameter) electrode. The catalyst ink was prepared as follows: 4 mg of h-MnOxPy (or h-Me-MnOxPy ) and 1 mg of carbon powder (Ketjen Black) were ultrasonically dispersed in a mixture of Nafion solution (20 μl, SigmaAldrich), water (250 μl) and ethanol (750 μl) for 30 min. 12 μl of the suspension was drop cast onto the polished GC electrode and dried overnight. The effective catalyst loading prepared as such was 0.30 mg cm-2. In comparison tests, the GC electrode was loaded with a 20wt% Pt/C or 20wt% Ir/C commercial catalyst to the same level (0.30 mg cm-2). Electrochemical measurements were carried out on an Autolab type III potentiostat/galvanostat using a rotating disk electrode (RDE). The electron transfer number (n) was calculated by the Koutecky-Levich (K-L) equation: 11 𝟏 𝒊

𝟏

𝟏

𝒌

𝒅

𝟏

= 𝒊 + 𝒊 = 𝒏𝑭𝑨𝒌𝒄𝟎 ―

𝟏 ―𝟏/𝟔 𝟎 𝟏/𝟐 𝟎.𝟔𝟐𝒏𝑭𝑨𝑫𝟐/𝟑 𝑪 𝛚 𝑶𝟐 𝛎

(1)

where i, ik , and id are the measured, kinetic, and diffusion-limited current, respectively. n is the overall electron transfer number, F is the Faraday constant, A is the geometric electrode area (cm-2), k is the rate constant for oxygen reduction, C0 is the saturated oxygen concentration in 0.1 M KOH (1.14*10 -6 mol cm -3), DO2 is the diffusion coefficient of oxygen (1.73*10-5 cm2 s-1), ν is the solution kinetic viscosity (0.01 cm2 s -1), and ω is the rotation rate (rad s-1) of the electrode.40 In order to correct for the capacitance effect in the measurements, the polarization curves of samples in O2-saturated solution were subtracted from the polarization curves in N2saturated solution. In-situ UV-vis transmission spectra as a function of the applied potential were recorded by an AvaSpec-UV/Vis spectrometer attached to the Autolab electrochemical workstation in a two-electrode setup. The spectrometer was connected to the light source by fiber-optics. The working electrode was fabricated as follows: h-MnOxP0.21 (or h-Co-MnOxP0.21) was ultrasonically dispersed in ethanol. The solution was then centrifuged at 5000 rpm for 10 min and the supernatant was drop cast onto a fluorine doped tin oxide (FTO) coated glass and dried at room temperature. A Zn foil was used as the reference electrode. The measured potentials were converted to the reversible hydrogen electrode (RHE) scale (ERHE =EZn/Zn2+ -0.76 +0.059×pH V). For cyclic voltammetry, the potential was scanned from 0 to 0.8V vs. RHE for ORR and from 0.8 to 2.4 V vs. RHE for OER in 0.1 M KOH at 5 mV s-1. Primary and rechargeable Zn-air batteries. Primary Zn-air batteries were configured with a Zn foil anode, a 6 M KOH aqueous electrolyte and a catalyst-loaded carbon paper air cathode. Rechargeable Zn-air batteries were configured with a Zn foil anode, a catalyst-loaded carbon

paper air cathode and alkaline or neutral electrolyte. The alkaline electrolyte for the latter consisted of 6 M KOH aqueous solution and 0.1 M zinc acetate. The neutral electrolyte was formulated from a 2 M NH4Cl and 0.1 M ZnCl2 base solution. The solution pH was adjusted to ~7 with ammonia solution. The air electrode was fabricated as follows: 2 mg of catalyst and 0.5 mg carbon (Ketjen carbon black) were ultrasonically dispersed in 1 ml ethanol containing 10 μl of Nafion solution. The as-formed slurry was drop cast on a carbon paper (1 cm x 1 cm) and dried in air. The effective catalyst loading prepared as such was 2 mg cm-2. For the primary Zn-air batteries, h-MnOxP0.21 was the catalyst for the air electrode. For the rechargeable Znair batteries, a mixture of h-MnOxP0.21 and h-Co-MnOxP0.21 (1:1 ratio by weight) was used. The primary Zn-air batteries were discharged at 10 mA cm-2 current density. The alkaline rechargeable Zn-air batteries were continuously discharged/charged at 5 mA cm−2 for 1 h per cycle. The neutral rechargeable Zn-air batteries were continuously discharged/charged at 1 mA cm−2 for 10 h per cycle.

Results and Discussion The synthesis of simple phosphated manganese oxide hollow nanospheres (h-MnOxPy) was carried out at room temperature; using a facile wet chemistry method which involved the mixing of MnSO4 and NH4H2PO4 in water; followed by the addition of KMnO4 in excess (more details in Experimental Section). h-MnOxPy with different shellthicknesses could be obtained simply by varying the MnSO4 to NH4H2PO4 ratio (by weight). The elemental compositions of variously synthesized h-MnOxPy as analyzed by energy dispersive spectroscopy (EDS) are given in Table S1. In the case of a low NH4H2PO4 content (MnSO4/NH4H2PO4 = 4, h-MnOxP0.08), the product was hollow nanospheres with a shell thickness of ~40 nm (Figure S1a). Increasing the NH4H2PO4 content decreased the shell thickness to ~20 nm (h-MnOxP0.16 and hMnOxP0.21, Figure S1b and 1c). A further increase in the NH4H2PO4 content (MnSO4/NH4H2PO4 = 1.2, h-MnOxP0.59, Figure S1d) led to the formation of a yolk-shell structure with a prominent internal cavity. The MnSO4/NH4H2PO4 ratio of 1.6 (h-MnOxP0.21) yielded the catalyst with the best ORR performance. As such h-MnOxP0.21 was examined in more detail as shown in the following. The broken spheres in the scanning electron microscopy (SEM) image in Figure 1a expose the hollow interior of the ~150 nm nanospheres that aggregated to form h-MnOxP0.21. An ensuing transmission electron microscopy (TEM) examination (Figure 1b) measured the shell to be ~ 20 nm thick. Further analysis by high-resolution TEM (Figure 1c) generated even more structural detail – that the nanospheres were constituted by very small (~10 nm) interconnected amorphous nanoparticles. The homogeneity of composition was confirmed by EDS elemental analysis which showed the uniform distribution of the Mn, O, P and K elements throughout the hollow nanospheres (Figure 1d).



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Figure 1. Microstructural characterizations of h-MnOxP0.21. (a) SEM image, (b)-(c) TEM images, (d) EDS elemental maps. The scale bars are all 100 nm, (e) Mn 3s XPS spectrum, (f) N2 adsorption/desorption isotherms and the pore distribution plot (inset); and (g) a digital photo of the amount of h-MnOxP0.21 that could be produced in a single batch.

The broad powder X-ray diffraction pattern (XRD, Figure S2) of h-MnOxP0.21 suggests a low crystallinity phase. The overall profile is nonetheless similar to that of cryptomelane MnO2 but with some differences in the peak intensity due to the presence of phosphates in the structure. The hollow nanospheres may therefore be referred to as phosphate-modified manganese oxide; or phosphated manganese oxide for short. The Mn oxidation states in h-MnOxP0.21 were analyzed by X-ray photoelectron spectroscopy (XPS). The two peaks in the Mn 3s spectrum (Figure 1e) with a peak separation of 5.39 eV are fairly typical of Mn (III).41 The existence of Mn (III) was also corroborated by cyclic voltammetry in Ar-saturated 0.1 M KOH solution (Figure S3).42 The phosphate presence in hMnOxP0.21 as PO43- is confirmed by the ~133 eV peak in the P 2p XPS spectrum. (Figure S4). The surface area and pore structure of h-MnOxP0.21 were determined by the Brunauer–Emmett–Teller (BET) method based on the measured N2 adsorption/desorption isotherms (Figure 1f). A surface area of ~80 m2 g-1 was calculated and the analysis also indicated the presence of both micropores and mesopores in h-MnOxP0.21. The use of low-cost common laboratory chemicals and room temperature reactions enabled the synthesis to be easily scaled up for gram-level production in a single batch, limited only by the size of the batch reactor used (Figure 1g).

A time-course study of the reactions was used to gain some insights of the hollow nanosphere formation process. The morphologies of the solid products sampled from different reaction stages (Figure 2b-2e) can best be rationalized by a redox reaction mediated self-templating mechanism (Figure 2a). At the start of the reaction with only a minute amount (3.5 ml) of KMnO4 solution (10 mg ml-1) present in the MnSO4/NH4H2PO4 mixture, porous nanoparticles (Figure 2b) were formed. XRD (Figure S5a, black line) and Fourier transform infrared spectroscopy (FT-IR, Figure S5b, black line) analyses suggested that they were amorphous ammonium manganese phosphatehydrogen phosphate. Since the Mn oxidation state in the porous nanoparticles was ~2+ (Figure S5c, black line), the nanoparticles were readily oxidized by KMnO4 to form a thin phosphated manganese oxide layer on the nanoparticle surface (Figure 2c), which acted as a physical barrier to reduce the direct contact between KMnO4 and the porous nanoparticles; and consequently the reactions between them. Since the MnO4- anions are larger than the Mn (II) cations in size, the outward diffusion of the Mn cations was faster than the inward diffusion of the MnO4ions. Therefore, the interior of the porous nanoparticles was gradually depleted; developing a void space between the core and the shell (Figure 2d and Figure 2e).


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Figure 2. (a) Schematics of the proposed mechanism of hollow structure formation. (b)-(e) TEM images of the products formed at different stages of the reaction by adding (b) 3.5 ml, (c) 5 ml, (d) 7 ml and (e) 15 ml of KMnO4 solution (10 mg ml-1) to the MnSO4/NH4H2PO4 mixture.

In a way, the in-situ formed starting porous nanoparticles self-templated the development of the final hollow nanosphere morphology; via the following categorical reactions and the accompanying Kirkendall effect.43 𝑴𝒏𝑺𝑶𝟒(𝒂𝒒) + 𝑵𝑯𝟒𝑯𝟐𝑷𝑶𝟒(𝒂𝒒) + 𝑲𝑴𝒏𝑶𝟒(𝒂𝒒)→ (𝑵𝑯𝟒)𝒙𝑴𝒏𝑰𝑰(𝑷𝑶𝟒)𝒚(𝑯𝟐𝑷𝑶𝟒)𝒛(𝒔) + 𝑯𝟐𝑶(𝒍) + 𝑲𝟐𝑺𝑶𝟒(𝒂𝒒) (2)

(𝑵𝑯𝟒)𝒙𝑴𝒏𝑰𝑰(𝑷𝑶𝟒)𝒚(𝑯𝟐𝑷𝑶𝟒)𝒛(𝒔) + 𝑲𝑴𝒏𝑶𝟒(𝒂𝒒)→ 𝑲𝒙𝑴𝒏𝑰𝑰𝑰𝑶𝒚(𝑷𝑶𝟒)𝒛(𝒔) + 𝑯𝟐𝑶 + (𝑵𝑯𝟒)𝟑𝑷𝑶𝟒(𝒂𝒒)

(3)

There were three noteworthy observations in the synthesis of these hollow nanospheres. First, the use of the dihydrogen phosphate salt (H2PO4-) was essential. Without the H2PO4-, solid manganese oxide nanoparticles (Figure S6) were formed when the KMnO4 solution was added to the MnSO4 solution. We believe that H2PO4- were coordinated to Mn (II) cation in the MnSO4 and NH4H2PO4 mixture in water. This was confirmed by the formation of a white precipitate upon ethanol addition, which is a characteristic of Mn(H2PO4)2 (soluble in water but not in ethanol). On the other hand, using a mono hydrogen phosphate salt (HPO42-) would result in the precipitation of white manganese hydroxide because of the (weak) alkalinity of HPO42-. The coordinated H2PO4- moderated the reaction between MnO4- and Mn (II) cation to form the self-template. Second, the product morphology, the formation of the hollow structure in particular, was not affected by the cations in the dihydrogen phosphate salt (i.e. NaH2PO4 and KH2PO4 would produce the same outcome as NH4H2PO4) (Figure S7). Third, the Mn redox reaction was facile enabling an ultrafast synthesis at room temperature - The hollow structure could be established in 2 min or less when sufficient KMnO4 was used. The preparation of composite phosphated manganese oxide hollow nanospheres (h-Me-MnOxPy, where Me= Co, Ni, Cu) is a simple extension of the above procedure. This was done by adding a solution of a second metal (Me) salt

in ammonia water to a freshly-prepared h-MnOxPy dispersion. The Me hydroxide in the former reacted with h-MnOxPy in the latter probably through an exchange reaction in the solid state to form Me-doped manganese hydroxyphosphate which was experimentally detected. The final product was likely a mixture of a Me-doped manganese hydroxyphosphate and unreacted h-MnOxPy and is denoted as h-Me-MnOxPy for convenience. We used the h-MnOxP0.21 as an example to prepare a series of h-MeMnOxP0.21 for further investigations. The hollow structure of h-MnOxP0.21 was retained in the h-Me-MnOxP0.21 as evidenced from the SEM and TEM images (Figure 3a-3i). However, the shell of h-Me-MnOxP0.21 now consisted mostly of nanosheets, instead of nanoparticles as in the case of h-MnOxP0.21 (Figure 3g-3i). EDS elemental mapping suggests the uniform distribution of the second metal element (Me) in the shell area of the hollow structure (Figure 3j-3l and Table S2). h-Co-MnOxP0.21 among the h-Me-MnOxP0.21 showed the best catalytic performance; and consequently its microstructures were examined in greater detail. XRD determined h-Co-MnOxP0.21 to be as poorly crystallized as its h-MnOxP0.21 precursor (Figure S8a). The broad peak around 2θ = 10-20o and a small peak at 2θ = 35o are the new features in the XRD pattern to suggest the presence of a cobalt hydroxyphosphate Co2(OH)(PO4) phase (JCPDS : 83-246). A large surface area of 324.85 m2 g-1 indicative of the preponderance of micropores and mesoporous was calculated from the N2 adsorption/desorption isotherms (Figure S8b). The large surface area of h-Co-MnOxP0.21 suggests an increase in the proportion of the micropores, which could be related to the observation of shell morphology changes. The presence of Co2+ was confirmed by the Co 2p XPS spectrum of h-Co-MnOxP0.21 (Figure S8c). The Mn oxidation state in h-Co-MnOxP0.21 was still 3+ as shown in the Mn 3s and Mn 2p XPS spectrum (Figure S8d and Figure S8e). The state of the phosphate group (PO43-) in h-Co-MnOxP0.21 was also unchanged as measured by the P 2p XPS spectrum (Figure S8f).



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doped mesoporous nanocarbon (NPMC-1000, E1/2 = 0.85V 13 vs. RHE), nitrogen doped carbon nanotube (N-CNT, E1/2 = 0.84V vs. RHE)45, Co3O4/N-rmGO (E1/2 = 0.83V vs. RHE)11 and others (Table S3). Low material cost and ease of preparation without advanced chemical procedures are its strength relative to the other Pt/C alternatives.

Figure 3. SEM images (a-c), TEM images (d-i) and EDS elemental maps (j-l) for h-Co-MnOxP0.21 (a, d, g, j), h-NiMnOxP0.21 (b, e, h, k) and h-Cu-MnOxP0.21 (c, f, i, l). All scale bars in the EDS elemental maps are 100 nm.

The performance of h-MnOxPy and h-Me-MnOxPy in oxygen electrocatalysis in alkaline solution (0.1 M KOH) was first screened by the rotating disk electrode (RDE) method. The screening identified h-MnOxP0.21 and h-CoMnOxP0.21 as the best-in-class for ORR and OER respectively (Figure S9 and Figure S10). These two catalysts were then subjected to additional electrochemical measurements. Figure 4a shows the linear sweep voltammograms (LSV) of ORR on h-MnOxP0.21 and h-CoMnOxP0.21, in comparison with 20wt% Pt/C and Ketjen carbon black. The ORR performance of h-MnOxP0.21 exceeded those of h-Co-MnOxP0.21 and carbon black in terms of half-wave potential (E1/2) and limiting current density. Indeed, the E1/2 of h-MnOxP0.21 at 0.85 V vs. RHE is as good as the benchmarking 20wt% Pt/C ORR catalyst. The 3.99 electron transfer number (n) calculated from the Koutecky-Levich (K-L) plot (Figure 4b and Figure S11) is significantly higher than those of h-Co-MnOxP0.21 (n = 3.34) and carbon (n = 2.58); and is Pt-like (n = 4.01). h-MnOxP0.21 also demonstrated good catalytic stability in ORR. Relative to Pt/C, h-MnOxP0.21 was more able to retain a larger fraction of its initial current after 20 h at 0.85 V vs. RHE (Figure 4c). The integration of catalytically active nanoparticles into a hollow nanosphere structure is therefore an effective mitigating measure against the loss of activity caused by the unabated aggregation of dispersed nanoparticles. Overall h-MnOxP0.21 surpasses most of the recent manganese oxide-based catalysts in ORR such as cCoMn2/C (E1/2 = 0.83 V vs. RHE),44 MnxOy/NC (E1/2 = 0.81 V vs. RHE),21 α-MnO2-SF/C (E1/2 = 0.76 V vs. RHE)16. It also compares favorably to the best of the non-noble metal ORR electrocatalysts reported to date, such as N and P co-

h-Co-MnOxP0.21, on the other hand, was more OER active than h-MnOxP0.21. Figure 4d shows that h-Co-MnOxP0.21 could deliver 10 mA cm-2 of current density at 1.60 V vs. RHE, which is ~160 mV lower than that required by hMnOxP0.21. Though the onset potential and the Tafel slope (Figure 4e) of h-Co-MnOxP0.21 in the low overpotential region were still higher than the 20wt% Ir/C benchmark, the potential required to support a current density higher than 10 mA cm-2 (E10mA cm-2) was interestingly lower. More importantly h-Co-MnOxP0.21 was more stable than 20wt% Ir/C in OER – the increase in the OER potential at 10 mA cm-2 after 45 cycles was smaller for h-Co-MnOxP0.21 than for 20wt% Ir/C (Figure 4f). However, there was a notable catalyst activity decrease after an extended (24 h) OER stability test, due to structural changes caused by Co oxidation, and Mn and PO43- leaching (Figure S12). The OER performance of h-Co-MnOxP0.21 is comparable to the best of non-noble metal OER electrocatalysts, such as ZIF derived carbon framework (E10mA cm-2 = 1.6o V),10 (Pr0.5Ba0.5)CoO3 (E10mA cm-2 = 1.56 V),46 CoSe2 nanosheet (E10mA cm-2 = 1.55 V) 47 CoMn LDH (E10mA cm-2 = 1.55V)48 (Table S4). These measurements verified the value of h-CoMnOxP0.21 as a low-cost, easily manufactured substitute for the noble metal OER catalysts. The large surface area of a hollow nanostructure with a porous shell (Figure 1f and Figure S8b) and the presence of PO43- groups are the most visible features that correspond well with the good ORR performance of h-MnOxP0.21 and the good OER performance of h-Co-MnOxP0.21 - a large surface area would provide more active sites for catalysis; and the phosphate groups could increase the ORR/OER performance by facilitating O2 and OH- adsorption on the active sites and ensuing proton transfers via the protoncoupled electron transfer (PCET) mechanism. 49-51 For more insights in-situ UV-vis spectroscopy over the course of ORR (for h-MnOxP0.21, 0.8 to 0.0 V vs. RHE) and OER (for h-Co-MnOxP0.21, 0.8 to 2.4 V vs. RHE) was used to follow the changes to h-MnOxP0.21 and h-Co-MnOxP0.21 during oxygen electrocatalysis (see the Experimental Section for more details). For h-MnOxP0.21 in ORR, difference spectra were generated by subtracting the spectrum at the potential of interest from the spectrum measured at 0.8 V vs. RHE (the open-circuit potential). Continuing scanning in the cathodic direction (Figure 4g) found the emergence of a broad peak in the 500-750 nm spectral region. The difference peak intensity grew with the increasingly cathodic potential and finally stabilized at the lowest potential in the scan (0 V vs RHE). This broad peak agrees well with the d-d transition of phosphate-coordinated Mn (III).52-53 The difference peak intensity increase also trends similarly as that of the ORR current density increase, suggesting Mn (III) as the most likely ORR active site.


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Figure 4. Oxygen electrocatalytic performance of h-MnOxP0.21 and h-Co-MnOxP0.21 in 0.1 M KOH solution in comparison with the commercial 20wt% Pt/C and 20wt% Ir/C catalyst benchmarks. (a) ORR linear sweep voltammograms (LSV) at 5 mV s−1 and 1600 rpm in O2 saturated electrolytes (b) The corresponding K-L plots at 0.70 V vs. RHE with error bars. (c) Chronoamperograms (% retention of initial current vs. time) at 0.85 V vs. RHE. (d) OER polarization curves at 5 mV s−1 and 1600 rpm in O2 saturated electrolyte. (e) The corresponding Tafel plots. (f) Chronoamperograms (% retention of initial current vs. time) at 10 mA cm-2. (g) In-situ UV-vis difference spectra of h-MnOxP0.21 at different ORR potentials (vs. RHE) and the potential dependence of the current density and absorbance difference at 650 nm plots. (h) In-situ UV-vis difference spectra of h-Co-MnOxP0.21 at different OER potentials (vs. RHE) and the potential dependence of the current density vs absorbance difference at 480 nm plots.

For the h-Co-MnOxP0.21 catalyst in OER, similar difference spectra were obtained by subtracting the spectrum of interest from the spectrum at 0.8 V vs. RHE. As shown in Figure 4h, a broad peak centering around ~410 nm started to appear at 1.0 V which corresponds well with the presence of Co (II).54 The difference peak became more intense and red-shifted to ~480 nm along with the increase in applied potential, suggesting the oxidation of Co (II). Cyclic voltammetry measurements (Figure S13) also confirmed the oxidation of Co (II) to Co (III) and the oxidation of Co (III) to Co (IV), as two anodic peaks at 1.6 V and 2.1 V vs. RHE, respectively. On the other hand, the current at potentials higher than ~2.2V vs RHE in Figure S13 was due to the OER. The stable UV-vis spectra after the OER onset suggests an unvarying Co oxidation state during the OER (Figure 4h, right panel). It may therefore be deduced that Co (IV) is the likely active site for OER on hCo-MnOxP0.21.

Encouraged by the favorable ORR and OER performance, we evaluated the h-MnOxP0.21 and h-CoMnOxP0.21 in the full-cell configuration of Zn–air batteries (Figure 5a). The h-MnOxP0.21 catalyst with good ORR performance was first benchmarked in a primary alkaline (6 M KOH) Zn-air battery. The measured discharge polarization curves (Figure 5b) show that the h-MnOxP0.21 cell could deliver a higher peak power density (120 mW cm2) than the Zn-air battery with a 20wt% Pt/C catalyst in its air electrode (~105 mW cm-2). The galvanostatic discharge curves (Figure 5c) also show that the h-MnOxP0.21 cell could exceed the Pt/C cell in specific capacity (~230 mAh cm-2 vs ~210 mAh cm-2) measured at the standard test current density of 10 mA cm−2. The good application performance of the primary Zn–air battery with the h-MnOxP0.21 catalyst ascertains the effectiveness of h-MnOxP0.21 for ORR catalysis in alkaline solution under practical application conditions.



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Figure 5. (a) The configuration of the Zn-air test cell. (b) Discharge polarization curves and (c) discharge plots of primary Zn-air batteries with a h-MnOxP0.21 or a 20wt% Pt/C catalyst in the air cathode at 10 mA cm-2. (d) discharge/charge polarization curves and (e) discharge–charge cycling of alkaline rechargeable Zn-air batteries using h-MnOxP0.21+h-CoMnOxP0.21 (red) or 20wt% Pt/C+20wt% Ir/C (blue) catalyst combinations. (f) Cyclic voltammograms and (g) Discharge– charge cycling of rechargeable Zn-air batteries in the neutral electrolyte using the h-MnOxP0.21+h-Co-MnOxP0.21 (red) or the 20wt% Pt/C+20wt% Ir/C (blue) combined catalyst system in the air electrodes.

For the rechargeable alkaline Zn-air test batteries, a mixture of h-MnOxP0.21 and h-Co-MnOxP0.21 (in a 1:1 ratio by weight) was used as the combined catalyst system. Figure 5d shows the discharge/charge polarization plots of the hMnOxP0.21 + h-Co-MnOxP0.21 cell in a 6 M KOH + 0.1 M ZnCl2 electrolyte, in comparison with a 20wt% Pt/C + 20wt% Ir/C cell in the same electrolyte. The performance of the two cells was almost identical. The h-MnOxP0.21 + h-CoMnOxP0.21 cell could even surpass the 20wt% Pt/C + 20wt% Ir/C cell at high current density. In typical discharge/charge cycling, the initial charge-discharge voltage gap of the h-MnOxP0.21 + h-Co-MnOxP0.21 cell (0.72 V at 5 mA cm-2 and 1 h/cycle, Figure 5e) was the same as that of the 20wt% Pt/C+ 20wt% Ir/C cell. The former was however more stable than the latter in extended cycling. After cycling for 90 h, the voltage gap of the h-MnOxP0.21 + h-Co-MnOxP0.21 cell was increased to 0.85 V, smaller than the 0.91V from the 20wt% Pt/C+ 20wt% Ir/C cell. The hollow structures of h-MnOxP0.21 and h-Co-MnOxP0.21 were still detectable after cycling but with a notable increase of the nanosheet presence in the shell (Figure S14a, b). We hypothesize that this was due to the oxidation of Co2+ and

the inclusion of Zn2+ in h-MnOxP0.21 + h-Co-MnOxP0.21 (Figure S14c-e). All of these measurements corroborate the usability of the h-MnOxP0.21 + h-Co-MnOxP0.21 combined catalyst system in actual rechargeable alkaline Zn-air batteries. Their performance matches well with the best of non-noble metal and metal-free air cathodes in the current rechargeable alkaline Zn-air batteries (Table S5), but with the added benefit of being low cost and easy fabrication. The h-MnOxP0.21 + h-Co-MnOxP0.21 combined catalyst system was also evaluated in a neutral rechargeable Zn-air test battery. Neutral rechargeable Zn-air batteries are not only the most environmentally benign and the safest at the point of use; but are also able to eradicate the corrosion and electrolyte carbonation issues of alkaline batteries. 5556 Neutral rechargeable Zn-air batteries are however rarely reported because of their low energy efficiency (large discharge-charge voltage gap) relative to the alkaline batteries, mandating the use of only high-activity catalysts. A mixture of 2 M NH4Cl and 0.1 M ZnCl2 was used as the electrolyte because of its good ionic conductivity.55 Its pH value was adjusted to 7 by ammonia solution. Standard cyclic voltammetry measurement was carried out using the


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full cell system with a Zn foil anode and an air electrode cathode. Figure 5f showed that the h-MnOxP0.21 + h-CoMnOxP0.21 system could deliver a respectable ORR and OER performance in the neutral electrolyte, and was only marginally lower than the 20wt% Pt/C+ 20wt% Ir/C system. Figure 5g shows that the initial charge-discharge voltage gap of the h-MnOxP0.21 + h-Co-MnOxP0.21 cell in (1.45 V) was higher than that of the 20wt% Pt/C+ 20wt% Ir/C cell (1.15 V, measured at a current density of 1 mA cm2 and 10 h/cycle). Interestingly, the performance of the hMnOxP0.21 + h-Co-MnOxP0.21 cell improved after a few cycles; with charge-discharge voltage gap narrowed down to 1.28V after 700 h of operation, which was smaller than the charge-discharge voltage gap of the 20wt% Pt/C+ 20wt% Ir/C cell (1.38 V). The increase in cell performance with cycling is often indicative of an activation process. The only material difference between fresh and cycled hMnOxP0.21 + h-Co-MnOxP0.21 catalysts was the presence of Zn and some loss of Co (Figure S15b). The causality between these observations has yet to be confirmed in further work. In addition, catalyst characterizations after the cycling tests indicated the persistence of the hollow structures of h-MnOxP0.21 + h-Co-MnOxP0.21; but with the increased presence of large nanosheets (Figure S15a); similar to the case of the alkaline electrolyte. These measurements are sufficient to demonstrate h-MnOxP0.21 + h-Co-MnOxP0.21 as one of the best catalyst systems for neutral rechargeable Zn-air batteries.

Conclusions In summary, a facile room-temperature wet chemistry procedure was used to prepare phosphated manganese oxides hollow nanospheres with noble metal-like ORR/OER performances. The preparation was based on insitu formed self-templates and phosphate-mediated redox reactions between Mn (II) cation and MnO4-. This method is effective, easily scalable; and can be used to synthesize both simple and composite phosphated manganese oxide hollow nanospheres. The as-synthesized h-MnOxP0.21 and h-Co-MnOxP0.21 showed good catalytic activities and stability in ORR and OER respectively, similar to the 20wt% Pt/C (for ORR) and 20wt% Ir/C (for OER) at a small fraction of the cost of the latter. In particular, the ORR halfwave potential of h-MnOxP0.21 was only 0.85 V vs. RHE and the OER potential of h-Co-MnOxP0.21 to support 10 mA cm-2 of current density was 1.60 V vs. RHE. Consequently, these phosphated manganese oxide hollow nanospheres were used as the combined catalyst system for the air electrode of demonstrative alkaline and neutral rechargeable Zn-air full cells. Small charge-discharge voltage gaps (ΔV) were shown for both alkaline (ΔV = 0.72 V at 5 mA cm-2) and neutral (ΔV = 1.28 V at 1 mA cm-2) rechargeable Zn-air batteries with the combined use of these catalysts, similar to the combined use of 20wt% Pt/C+ 20wt% Ir/C catalysts The methods shown here may also have general utility for the preparation of high-performance hollow nanomaterials for electrocatalysis and energy conversion; an area that is definitely worth pursuing.

ASSOCIATED CONTENT Supporting Information. : This material is available free of charge via the Internet at http://pubs.acs.org. TEM images, XRD pattern and elemental analysis of hMnOxPy; characterizations of h-Co-MnOxP0.21; electrochemical measurements of h-MnOxPy and h-CoMnOxP0.21.

AUTHOR INFORMATION Corresponding Author *E-mail: [email protected]

Author Contributions The manuscript was written through contributions of all authors.

Notes

The authors declare no competing financial interest.

ACKNOWLEDGMENT This work was supported by grants from the National Research Foundation (NRF), Prime Minister’s office, Singapore, under the Campus for Research Excellence and Technological Enterprise (CREATE) program.

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A Self-Templating Redox-Mediated Synthesis of Hollow Phosphated

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