and K-Ion Batteries

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C: Energy Conversion and Storage; Energy and Charge Transport

Manganese Hexacyanomanganate as a Positive Electrode for Non-Aqueous Li-, Na- and K-Ion Batteries Viktor Renman, Dickson O. Ojwang, Cesar Pay Gòmez, Torbjörn Gustafsson, Kristina Edstrom, Gunnar Svensson, and Mario Valvo J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.9b06338 • Publication Date (Web): 26 Aug 2019 Downloaded from pubs.acs.org on August 27, 2019

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The Journal of Physical Chemistry

Manganese Hexacyanomanganate as a Positive Electrode for NonAqueous Li-, Na- and K-ion Batteries Viktor Renman1‡, Dickson O. Ojwang2‡, Cesar Pay Gómez1, Torbjörn Gustafsson1, Kristina Edström1, Gunnar Svensson2 and Mario Valvo1* 1Department

of Chemistry – Ångström Laboratory, Ångström Advanced Battery Centre, Uppsala University, Box 538, SE751 21 Uppsala, Sweden. 2Department of Materials and Environmental Chemistry, Stockholm University, SE-106 91 Stockholm, Sweden.

ABSTRACT: K2Mn[Mn(CN)6] is synthesized, characterized and evaluated as possible positive electrode material in non-aqueous Li-, Na- and K-ion batteries. This compound belongs to the rich and versatile family of hexacyanometallates displaying distinctive structural properties, which makes it interesting for ion insertion purposes. It can be viewed as a perovskite-like compound in which CN-bridged Mn(CN)6 octahedra form an open framework structure with sufficiently large diffusion channels able to accommodate a variety of insertion cations. By means of galvanostatic cycling and cyclic voltammetry tests in non-aqueous alkali metal half-cells, it is demonstrated that this material is able to reversibly host Li+, Na+ and K+ ions via electrochemical insertion/de-insertion within a wide voltage range. The general electrochemical features are similar for all these three ion insertion chemistries. An in operando Xray diffraction investigation indicates that the original monoclinic structure is transformed into a cubic one during charging (i.e., removal of cations from the host framework) and that such a process is reversible upon subsequent cell discharge and cation reuptake.

1. Introduction Heavy dependence on fossil fuels as main sources of energy production has contributed significantly to the global warming problems. With ever-increasing world population, the need for alternative cleaner and more efficient energy sources is becoming even more urgent. In this regard, solar, wind, and hydroelectric power are by far the best renewable energy sources. However, the former two are yet to be effectively integrated into the electrical power grid due to a number of limitations.1 For instance, solar and wind energy are weather dependent and, as such, highly intermittent, while pumped hydroelectric power and compressed air systems for large-scale energy storage and conversion are both site specific, require high initial investments to set up and are not optimal in terms of quick electrical response. Technological demands for improved electrical energy storage systems require designing new classes of electrode materials to reach a series of important targets (i.e., safety, cost-effectiveness, minimal carbon footprint, long cycle life, faster kinetics, high-power operation and good energy efficiency). Currently, it is extremely challenging to simultaneously meet all of these requirements, nevertheless, rechargeable batteries are receiving increasing attention in a view to achieve a progressive transition toward a fossil fuel-free society. Rechargeable batteries are expected to replace most of the current energy storage technologies because of their simple installation and higher efficiency.2 Li-ion batteries (LIBs) are the best performers for portable devices in terms of energy density and long calendar life, however, they are too expensive for stationary applications due to their high fabrication costs, limited raw material resources and safety issues.3 Conversely, redox flow batteries (e.g., vanadium redox flow battery) display

low energy densities and limited rate performances, while leadacid batteries are associated with poor cycling life and cannot withstand deep discharge.4-6 Sodium-sulphur (NaS) batteries, which are used for large-scale storage, demand operation at elevated temperatures (300-350 °C) and have low energy efficiency,7 while metal hydride batteries are currently too expensive for the power grid.8-10 The global energy situation has stimulated researchers to develop new battery technologies with better and cheaper materials, mainly because lithium exists in restricted reserves, which could lead to a supply crisis for future large scale Li-ion use.11-13 Hence, research in the field of reversible electrochemical energy storage based on potassium chemistries has lately sparked considerable interest. K-ion batteries, together with Na-ion batteries, are under scrutiny as possible candidates to enable complementary functionalities to those of Li-ion systems and advance stationary storage of electricity, especially in a perspective of materials, environmental footprint and costs. Potassium-ion batteries are particularly interesting, since the reduction potential of K in non-aqueous electrolytes is close to that of Li or even lower14 and potassium abundance in the earth crust is considerable, thus potentially supporting contained fabrication costs. Besides, K+ can be intercalated electrochemically into graphite and simultaneously exhibits favorable insertion properties with respect to the rich family of compounds known as hexacyanometallates. Prussian blue analogues (PBAs) or hexacyanometallates, i.e., compounds with the generic formula AxMy[M’(CN)6]z·nH2O (A = e.g. alkali cations, Zn2+, Al3+; M, M’ = transition metal ions), are indeed attractive candidates for use in large-scale batteries to support the electrical grid. Their structural frameworks are

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time of 300 s and 2θ range between 10 ° and 70 °. The water content in the compound and the dehydration temperature were estimated via thermogravimetric analysis (TGA) using a TA instruments, Discovery thermobalance. The sample was placed in a Pt crucible and progressively heated from RT to 600 °C in N2 with a heating rate of 10 °C/min. The mid-infrared (4004000 cm-1) absorption spectra were recorded using a Varian 610-IR spectrometer equipped with a deuterated triglycine sulfate (DTGS) detector. Raman spectra were attained with a Renishaw InVia spectrometer using an excitation wavelength of 532 nm produced by a solid-state laser (Renishaw). The laser beam was focused on the surface of the specimens through a ×50 objective and an initial calibration was run before the measurements using a Si wafer as standard reference to obtain a characteristic peak at 520.6 cm-1. Thirty cumulative acquisitions with a measuring time of 60 s were performed between 200 and 1800 cm-1 for each spectrum to enhance the S/N ratio. A low nominal laser power of 0.5 mW was used in all the analyses together with a minimization of the beam exposure time in between subsequent scans to prevent possible degradation of the sample surface during the measurements.

closely related to ABX3 perovskite structures, except that cyanide bridges, instead of oxide ions, connect the octahedral metal centers. PBAs structures consist of a rigid threedimensional (3D) cubic network of linearly repeating -NC-MCN-M’-NC units15,16 leading to an octahedral coordination of the M and M’ cations. The void spaces or cavities formed by the network are occupied by zeolitic water molecules or alkali ions.17,18 In light of their electrochemical capability, the PBAs show a particular feature to undergo electrochemical insertion/extraction reactions with the participation of multivalent cations for charge compensation. Due to this extreme versatility in their ability to act as suitable hosts for various cations, both in aqueous and non-aqueous electrolytes, interest has blossomed among battery researchers in this large family of fascinating compounds in recent years. Prussian blue (PB) and PBAs have come to prominence as possible alternative electrode materials, owing to their particular characteristics and potentially limited environmental impact. These electrodes have a number of merits such as long cycle life, faster kinetics and higher energy efficiency than any other family of electrodes when operated in aqueous electrolytes.19,20 On the basis of these considerations, the present study aims at investigating the manganese hexacyanomanganate framework compound (Mn[Mn(CN)6]) in the context of non-aqueous rechargeable batteries in presence of Li+, Na+ and K+ ions.

2.2 Electrochemical measurements Electrochemical characterization was carried out using composite electrodes assembled in a pouch cell configuration. The active material, K2Mn[Mn(CN)6], was ball-milled with carbon black (Super P, Timcal) and polyvinylidene fluoride (PVDF, Solef 5130) in a weight ratio of 80:10:10 using Nmethyl-2-pyrrolidone (NMP, Sigma Aldrich) as solvent for the PVDF binder. The viscous slurry was spread onto a carboncoated aluminium foil using a weighted bar coater (KR – K Control Coater). After evaporating the solvent, circular electrodes were punched out and subsequently dried overnight at 120 °C in a vacuum oven inside an Ar-filled glovebox (MBraun). The electrode coatings typically had an active mass loading of ca. 3 mg cm-2. For the electrochemical tests, halfcells were assembled. The electrode stacks, which consisted of the prepared positive electrodes, a glass fiber separator soaked in electrolyte and either a Li-foil (Cypress Foote) or pin-rolled circular Na or K metal (Sigma Aldrich) disk (serving in all the cases as combined counter- and reference electrode), were vacuum-sealed within a polymer-aluminum-polymer laminate pouch cell. The electrolyte used was APF6 (A = Li+, Na+, K+) dissolved in an ethylene carbonate-diethyl carbonate (EC:DEC) 1:1 v/v mixture. For the Na and K half-cells, an additional 5 wt.% of fluoroethylene carbonate (FEC) was added to the electrolyte in order to stabilize their highly reactive metal surfaces. Galvanostatic (i.e., constant current) experiments were carried out using a Digatron BTS600 battery cycler at RT. Cyclic voltammetry (CV) analyses were also performed under analogous conditions utilizing a BioLogic MPG2 potentiostat. For the in operando XRD experiment, composite electrodes with a higher active mass loading of ~10-12 mg cm-2 were prepared by spreading a viscous slurry of the same composition as above onto a fine Al mesh. After evaporating the binder solvent under vacuum at 60 °C, the electrodes were initially pressed and then dried again at 120 °C under vacuum to expel possible residual water from the composite coating. For this experiment, a Li half-cell was utilized and a thin polyethylene separator (Solupor) was employed rather than glass fiber in order to minimize the influence of cell hardware on the resulting

2. Experimental 2.1 Materials synthesis and characterization Mn(NO3)2·4H2O and KCN salts were used as purchased from Sigma-Aldrich. K2Mn[Mn(CN)6]·nH2O was synthesized following a precipitation procedure earlier reported by Qureshi and Sharpe.21 Briefly, distilled water was deoxygenated through boiling and flushing with nitrogen gas for 1 h. Solutions of 0.25 M Mn(NO3)2·4H2O and 1.25 M KCN were then separately prepared using the deoxygenated water. This was followed by dropwise addition of Mn(NO3)2·4H2O solution to that of KCN under constant stirring at room temperature (RT). A gray precipitate immediately formed and progressively turned its color to yellow, to blue and then to light green within 5 minutes. The precipitate was centrifuged after 10 minutes, washed several times with air-free ethanol and finally freeze-dried overnight. All syntheses were performed in a dark container placed in a glove box. It should be noted that the compound is susceptible to oxidative hydrolysis. If kept under ambient conditions, in air, the green powder gradually turns brown over the course of several days, indicating an early degradation. The material was thus carefully stored in an Ar-filled glovebox after its synthesis to prevent this detrimental process. A JEOL JSM-7401F scanning electron microscope was used to visualize the sample morphology and particle size. It was operated at an accelerating voltage of 1-2 kV and with a working distance of 3 mm. Energy dispersive X-ray spectroscopy (EDS) analyses were performed on a HITACHI TM3000 electron microscope to determine the cation compositions and obtain a semi-quantitative estimation of the C and N contents. X-ray powder diffraction (XRPD) patterns were collected using a PANalytical X’pert PRO X-ray diffractometer operating in a Bragg-Brentano geometry at 40 kV and 40 mA with CuKα1 (λ = 1.5406 Å) radiation. The diffraction patterns were collected with a step size of 0.03°, step

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diffractograms. This in operando XRD experiment was carried out in transmission mode utilizing a Stoe Stadi P diffractometer equipped with a Dectris Mythen 1K position sensitive detector, which covers a 2θ range of 19° in stationary mode. The cell was cycled at a rate of C/5 and a diffraction pattern was collected every ten minutes. Ex-situ XRD samples were prepared by cycling cells to a desired potential and then disassembling the cells inside an Ar-filled glovebox to retrieve the electrodes. After rinsing the electrodes with dimethyl carbonate (DMC) to remove excess electrolyte salt, the sample electrodes were enclosed between two layers of polyimide adhesive tape (Kapton) in order to provide protection against atmospheric conditions during the short duration of X-ray data collection.

atom

site

X

y

z

occ.

Biso(Å2)

Mn1

2a

0.0000

0.0000

0.0000

1

1.3(3)

Mn2

2d

0.0000

0.5000

0.5000

1

1.2(2)

N1

4e

0.451(2)

0.253(3)

0.856(3)

1

1.8(2)

C1

4e

0.536(2)

0.641(4)

0.270(4)

1

1.8(2)

N2

4e

0.052(2)

0.342(3)

0.751(3)

1

1.8(2)

3. Results and discussion 3.1 Structure and morphology The phase purity of the target compound was confirmed by means of powder X-ray diffraction (Fig. 1). The crystal structure of K2Mn[Mn(CN)6] was refined according to the Rietveld method.22 A monoclinic structural model provided by Her et al.23 (S. G. P21/n) resulted in a good refinement fit without considerable modification. The refined cell parameters and volume are a = 10.1660(37) Å, b = 7.3879(22) Å, c = 6.9832(26) Å, β = 90.1685(48) °, and V = 524.473(31) Å3. The atomic coordinates and parameters of the refined structural model are displayed in Table 1. In hexacyanometallates, the metal atoms coordinated to the C and N ends of the cyanide groups are typically found in low-spin and high-spin states, respectively. From a structural viewpoint, Mn2+ is situated in two distinct crystallographic positions (0, 0, 0) and (0, ½, ½). The Mn2+ ion sitting at (0, 0, 0) is in a low-spin state, as it is octahedrally bonded to six C atoms with one short Mn-C distance of 1.948(2) Å and two long ones of 2.015(1) Å and 2.057(2) Å. On the other hand, the Mn2+ ion sitting at (0, ½, ½) is in a high-spin state, as it is octahedrally bonded to six N atoms with Mn-N atomic distances of 2.172(2) Å, 2.178(1) Å and 2.268(4) Å. The difference in bond lengths for both the lowspin and high-spin Mn2+ ions is ~ 0.1 Å. This value is of the same magnitude as those reported in literature and may well be credited to slightly distorted MnC6 and MnN6 octahedra.23,24 The average Mn-C-N and Mn-N-C bond angles are 165.37(5) ° and 143.65(5) °, respectively. These values deviate significantly from linearity (i.e., 180 °). Similar values have also been reported previously for A2Mn[Mn(CN)6] (A = K, Rb).23 The linearity of the PBAs framework arrangement is dependent on the ionicity of the M’-CN and M-NC bonding, viz., if they are covalent then the M’-CN-M angles should be linear. However, the bond angles may sensibly deviate from linearity when cations with strong electrostatic interactions, such as K+, enter PBA structures with less covalent (more ionic) M species, e.g., MnII.25 This case directly applies to the unusual bond angles observed here for K2Mn[Mn(CN)6].

C2

4e

0.961(3)

0.774(4)

0.151(4)

1

1.8(2)

N3

4e

0.204(2)

0.454(2)

0.380(2)

1

1.8(2)

C3

4e

0.698(2)

0.487(2)

0.562(2)

1

1.8(2)

K1

4e

0.750(1)

0.9315(3)

0.4780(5)

0.99(1)

3.3(1)

Rp = 4.4%, Rwp = 5.9%, χ2 = 3.9.

Figure 1. The Rietveld refinement of the as-prepared K2Mn[Mn(CN)6]. The results of the structural refinement (Rietveld) on the as-prepared powder. Red circles indicate experimental data, the black line represents the diffraction pattern generated from the structural model employed here. The blue line is the difference curve between the observed and calculated XRD patterns, respectively, and the green ticks refer to characteristic Bragg positions. The inset provides a schematic visualization of the crystal structure along the a-axis displaying the open structural framework in which conjoined Mn(CN)6 units form open windows for the mobile cations.

The scanning electron microscopy (SEM) images shown in Fig. 2 reveal particles with sizes spanning approximately from 400 nm to 1 μm. The individual particles also appear to have distinct morphologies that can tentatively be described as distorted cubes, somehow reflecting the monoclinic crystal structure of the compound and its low amount of defects or vacancies. For instance, the precipitation of K2Mn[Mn(CN)6] in aqueous solution is somewhat slower compared to that of the PBA-copper hexacyanoferrate (CuHCF), thereby leading to the formation of significantly larger primary particles. Altogether, particle sizes in the sub-micrometer range are beneficial in terms of short diffusion paths for mobile ions through this structure and exposure of a larger electrochemically active surface area.

Table 1. Synoptic table of fractional atomic coordinates and parameters of the refined structural model for K2Mn[Mn(CN)6 in space group no. 14 setting P21/n. The lattice parameters of the monoclinic unit cell are a =10.1660(37) Å, b = 7.3879(22) Å, c = 6.9832(26) Å, and β = 90.1685(48) °.



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region of the spectrum in Fig. 3a. This corresponds to weakly hydrogen-bonded water. Accordingly, such a very low intensity of the ν(OH) band suggests that only a very small amount of water is present in the structure, even though its initial preparation occurred in an aqueous medium. The Raman spectrum in Fig. 3b shows the characteristic ν(CN) vibrations for K2Mn[Mn(CN)6]·0.1H2O. The peaks appear within the same spectral region of the main IR absorption and highlight two major features at 2073 cm-1 and 2097 cm-1, which may be related respectively to Eg and A1g modes for ν(CN), in analogy with earlier Raman spectra for K4[Mn(CN)6]·3H2O.27,30 The shift towards higher wavenumbers for the peaks of K2Mn[Mn(CN)6]·0.1H2O, compared to those of K4[Mn(CN)6]·3H2O, is likely due to the fact that here both C and N are bonded to Mn in K2Mn[Mn(CN)6]. This gives rise also to asymmetric ν(CN) stretching modes due to symmetry-breaking Jahn-Teller distortion. The Eg mode at 2073 cm-1 appears rather broad compared to the sharper and more intense A1g vibration. A moderate overlap of the Eg and A1g modes can be noted as well. The faint feature around 2120 cm-1 in Fig. 3b can likely be associated with a very small amount of H2O, which could form water-Mn-cyanide moieties.31

Figure 2. SEM micrographs at different magnifications of the asprepared K2Mn[Mn(CN)6] powder. The particle size ranges approximately from 400 nm to 1 µm and the morphology of the individual crystallites displays features akin to those of distorted cubes.

3.2 Vibrational spectroscopy analyses Infrared and Raman analyses were employed to ascertain the presence of characteristic vibrational frequencies for the (CN)group and water molecules of the as-prepared powder sample, as well as to gain insight into possible spectral features related to the slightly distorted MnC6 and MnN6 octahedra. The infrared absorption frequencies associated with the ν(CN) stretching mode in PBAs are typically located around 2100 cm1. The CN stretching frequency can be used to determine the oxidation state, coordination number and electronegativity of the metals bound to the CN ligands.26 The ν(CN) stretching mode of [MnII(CN)6]4- in PBAs was previously reported to occur in a wavenumber range 2 in AxMn[Mn(CN)6] (A = Li+, K+). By comparison with the K half-cell in Fig. 5e this feature seems rather pronounced, which indicates that the structural framework can accommodate an excess of smaller Li+ ions.

Figure 4. Thermogravimetric (TG) curve of the as-prepared K2Mn[Mn(CN)6] heated in N2 atmosphere with a heating rate of 10 °C/min. Note the elevated thermal stability of this compound.

3.4 Electrochemical properties The general electrochemical features obtained upon testing this active material in various cells with non-aqueous electrolytes are displayed in Fig. 5. Here, the as-synthesized Kcontaining manganese hexacyanomanganate composite electrodes are tested directly in alkali metal half-cells containing APF6 (A = Li+, Na+, K+) salts in the electrolyte. Hence, in the case of Li and Na, these are technically hybrid cells since small amounts of K will always be present. However, the residual amounts of K from the pristine active material is estimated to be relatively low in relation to the excessive amount of Li and Na which comes from both the metal anode and the electrolyte. Galvanostatic and cyclic voltammetric experiments demonstrate the manganese hexacyanomanganate (MnHCM) structural framework’s ability to reversibly host Li+, Na+ and K+ ions as a positive electrode in non-aqueous cells. In broad terms, the different cell chemistries exhibit almost the same features, except in the low voltage region, where some additional capacity due to alkali ion insertion is observed for the

3.4.2 Na half-cell cycling While it has previously been demonstrated that up to 3 Na+ per formula unit can be inserted into the structural framework, if pushed to a low voltage of 1.6 V vs. Na+/Na,34 the observed effect of this phenomenon is less pronounced here, at least in the first cycle. In fact, the extent of the lowest plateau in the Na cell grows as the number of cycles increases. It has also been shown that similar hexacyanometallate frameworks exhibit a selectivity toward K+ compared to Na+ and Li+.35 One conceivable explanation for this particular behavior observed here could be that residual amounts of K+ are preferentially cycled in the host material, thus simply blocking further insertion of Na+ at these low potentials in the first cycle(s). The gradual increase of this electrochemical feature might be related to a successive dilution of K+ in the framework and electrolyte.



Here, the amount of electrolyte used was in excess and an oversized, well-soaked separator was used as well. It is estimated that the ratio of Na+/K+ in the system is ~20, hence the electrochemical behavior is likely to be dominated by Na, eventually.

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It is worth remembering that the various cell chemistries exhibit slightly different properties in terms of voltage hysteresis (K > Na > Li) between the insertion (discharge) and extraction (charge) processes (see Fig. 6d-f). It is known that the high reactivity of a freshly exposed Na metal surface immediately triggers the formation of a Solid Electrolyte Interphase (SEI) upon contact with common electrolytes, and that the resulting SEI layer has disadvantageous electrochemical properties in terms of interfacial resistances as compared to an analogous Li metal-based counterpart.36,37 This, in turn, gives rise to larger overpotentials which cannot be ascribed to variations in ionic conductivities of alkali-based electrolytes alone. 3.4.4 Slow cycling vs. fast cycling Cycling the half-cells at different C-rates yields significant changes in the electrochemical behavior in terms of voltage polarization, accessible capacity as well as capacity retention as a function of cycle number. At a relatively slow rate of C/5, the capacity retention for all three systems (i.e., Li+, Na+ and K+) appears to decrease in a similar fashion, as indicated by a rapid capacity fading during the 100 cycles reported in Fig. 6a-c. The initial discharge capacity of the Li half-cell cycled at C/5 is 174 mAh g-1, while after 100 cycles it drops dramatically to 25 mAh g-1, implying that only ≈ 14% of the capacity is retained. Clearly, the long-term stability of this material appears rather poor. There are a number of possible explanations for this, which are most likely associated with well-known problems typically observed in Mn-based systems. The Mn3+ species is known to be rather unstable and has a tendency to disproportionate to Mn2+ and Mn4+.38 Overall, such a rapid fading suggests that the crystal structure of MnHCM might be unstable upon prolonged electrochemical cycling. Interestingly, if the Li half-cell is cycled five times faster at a rate of 1 C, the capacity retention over 100 cycles is significantly improved. With an initial discharge capacity of 120 mAh g-1 it drops to 79 mAh g-1 after 100 cycles, that is ≈ 66% of the capacity is retained. On the other hand, only about 69% of the initial capacity is retained when going from C/5 to 1 C and the voltage polarization between charge and discharge increases by ≈ 0.21 V as well. Judging by the appearance of the voltage profiles at different cycling rates, it is evident that the prominence of the plateau at ≈ 3.8 V vs. Li+/Li during discharge is drastically diminished for the profile obtained at 1 C as compared to the one in correspondence to C/5. This is likely related to the sluggish rate performance of N-coordinated Mn2+/3+ redox couple, a phenomenon which has been observed in structurally and chemically similar PBAs.34 In other words, when the cell is cycled fast, the utilization of the N-coordinated Mn2+/3+ redox couple appears to be lower, thereby circumventing structural instabilities associated with its presence. For the Na half-cell cycled at C/5, the capacity fades rapidly in a manner similar to that of its Li counterpart. An initially promising discharge capacity of 129 mAh g-1 drops to a mere value of 13 mAh g-1 after 100 cycles. At a rate of 1 C, the electrochemical signature of the N-coordinated Mn2+/3+ redox activity appears modest, in line with what is observed in the Li half- cell as well. Furthermore, the voltage polarization between charge and discharge increases by ≈ 0.22 V. After 100 cycles at

Figure 5. Qualitative comparisons of the galvanostatic (chronopotentiometric) cycling curves of the five initial cycles of the MnHCM framework tested in (a) Li half-cell, (c) Na half-cell and (e) K half-cell, respectively. Note that the gravimetric capacities in the galvanostatic experiments, as well as the currents in the voltammetric measurements, have been normalized with respect to the mass of the initial K2Mn[Mn(CN)6] in all three cases. The cyclic voltammograms obtained at a scan rate of 0.1 mVs-1 of the respective cell chemistries are displayed in (b, d, f) to highlight conveniently their characteristic redox peak features.

3.4.3 K half-cell cycling The most obvious difference in the cycling behavior of MnHCM in a K half-cell is manifested at low voltages in Fig. 5e. In the previous experiments, it was demonstrated that an excess (in relation to the pristine material) of both Li+ and Na+ ions can be inserted into the framework at low potentials. In the pure K-system, this behavior is practically absent. The simplest explanation for this would be that the larger K+ ions occupying the structural voids merely block additional insertion. However, it cannot be ruled out that higher overpotentials are possibly required for an additional insertion of K+ into the structure. Again, the overall trend for this compound is a gradual weakening of the electrochemical signatures related to the redox activity of N-coordinated Mn2+/Mn3+.


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The Journal of Physical Chemistry 3.5 Structural transitions during electrochemical cycling 3.5.1 In operando XRD In operando XRD was used to monitor structural changes upon repeated removal and insertion of mobile cations into the MnHCM host framework. A Li half-cell was chosen for the in operando experiments. The reason is two-fold, firstly because the thin Li-foil provides higher quality XRD data than thicker pin-rolled Na or K, and secondly in order to elucidate the additional feature observed at low voltages in the Li cells. The results are summarized in Fig. 7. During the first charge, i.e., upon electrochemical extraction of cations from the pristine material, a decrease in intensity of existing peaks along with the appearance of new ones at lower angles is consistent with a series of phase transitions from the original monoclinic phase to a new one of higher symmetry. For a number of related structures in the family of PBAs, an expansion of the unit cell upon removal of cations from the host framework is usually observed.40,41 This arguably counterintuitive behavior has been ascribed to a lengthening of the M’-N bond as an electron is removed from the M-C-N-M’ π bonding system. Here, the shifting of the diffraction peaks to lower angles indeed suggests that an expansion of the unit cell takes place. MM’(CN)6 structures (M, M’ = transition metals) are often found to be cubic, yet structures with lower symmetries such as tetragonal, rhombohedral or even orthorhombic exist.23,24,34 Hence, the transition from a monoclinic phase to a cubic one can be qualitatively justified by a decrease of the monoclinic angle β and, thereby, an overall expansion of the unit cell. When discharging the Li half-cell, the diffraction peaks gradually revert to their original positions, suggesting a short-term structural stability and reversibility. Upon deep discharging (i.e., by inserting Li+ beyond the initial state of potential), the diffraction peaks become smeared out and are slightly shifted to higher angles. This suggests that in the fully discharged, deeply lithiated state, the excess amount of Li+ inserted into the host material involves a stronger distortion of the monoclinic structural framework. Upon charging the Li half-cell in the second electrochemical cycle, the main electrochemical features in the voltage profiles reappear and the characteristic peaks corresponding to the monoclinic phase regain their sharpness, thus indicating the structure’s ability to reversibly accommodate a high cation (i.e., Li+) content without destroying the framework, at least in the short-term. After additional removal of alkali cations from the structure, the tetragonal and cubic fingerprints also reappear, thereby exhibiting a similar behavior as in the first cycle, in accordance with the voltage profile of the first cycle. Here, it appears as the first voltage plateau centered at 3.2 V vs. Li+/Li corresponds to the onset of a phase transition from the original monoclinic K2Mn[Mn(CN)6] to a partially depotassiated tetragonal phase. A LeBail fit of the newly formed phase could be matched with a space group I-4m2 and a tetragonal unit cell with cell parameters a ≈ 10.5 Å and c ≈ 7.1 Å.42,43 Upon further electrochemical removal of K+ at the end of charging, the structure changes to a cubic one (Fm-3m) with a ≈ 10.7 Å. Such a cubic structure and a value are commonly reported for PBAs, e.g., CuHCF.40

a faster rate, the gravimetric discharge capacity is 40 mAh g-1, which corresponds to a retention of ≈ 50%. The K half-cell behaves in a similar fashion to both the Li and Na half-cells. At a cycling rate of C/5, the initial discharge capacity is 101 mAh g-1. This number drops to 10 mAh g-1 after 100 cycles, corresponding to a capacity retention of only ≈ 10%. If dissolution of Mn is the main cause of such stability issues,39 it makes sense to observe a higher capacity retention with fast cycling compared to slow cycling, since the detrimental effects of such a process have less time to fully manifest. In this respect, it appears as if the accessibility of the N-coordinated Mn2+/Mn3+ redox couple plays a crucial role here. Its electrochemical sluggishness leads to an underutilization of the redox-active material when the cells are cycled fast. A granted penalty in the accessible capacity is suffered as a result. On the other hand, the long-term cycling stability seems to benefit from the indirect exclusion of some of the Mn3+, thereby providing more hints at the structural instability of this compound. Overall, it can be concluded that the cycling rate of the various half-cells is of high importance with respect to the long-term electrochemical properties of these Mn-based hexacyanometallates.

Figure 6. The capacity retention of MnHCM in three different half-cell types cycled at a slower (C/5) and a faster (1 C) rate are displayed in (a, b, c) for Li+, Na+ and K+ ions, respectively. The galvanostatic voltage profiles of the corresponding initial cycles for these half-cells are shown in (d, e, f), respectively.



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more reliable unit cell parameter refinements. For a fair comparison of the samples at different states of charge, electrodes with the same mass loadings were chosen for each experiment and their XRD measurements were conducted under identical experimental conditions. The intensities in the diffractograms in Fig. 8 have therefore not been normalized. A structureless LeBail fit of a pristine electrode (Fig. 8a) confirms that K2Mn[Mn(CN)6] does not undergo decomposition during the electrode preparation procedure. In the fully charged (i.e., fully oxidized) state, a LeBail fit matches well with a cubic structure in space group Fm-3m with a unit cell parameter a = 10.696 Å, see Fig. 8b.

Figure 8. Ex situ XRD patterns of: (a) a pristine (uncycled) electrode on an Al current collector; (b) an identical electrode which has been fully charged to 4.5 V vs. Li+/Li; (c) an identical electrode which has been fully discharged to 1.5 V vs. Li+/Li (i.e., after one full electrochemical cycle). The experimental data points are indicated as red circles. The black lines correspond to the calculated patterns obtained in structureless LeBail fits. The blue lines indicate the difference between the experimental and calculated data. The grey circles denote the characteristic diffractions generated by the Al current collector.

Figure 7. (a) In operando X-ray diffraction patterns displaying the structural evolutions upon repeated removal and insertion of cations (i.e., Li+) into the MnHCM host framework. The galvanostatic voltage profile of a Li half-cell with K2Mn[Mn(CN)6] cycled at C/5 for two full charge-discharge cycles (shown on the left) is accompanied by a contour plot (right, top-view, intensity axis is perpendicular to the plane of the figure) of the XRD patterns collected at 10 minutes intervals. In this inverse grey scale representation, black indicates the diffraction peaks. Note that the structural transitions of the MnHCM host framework are schematically indicated by horizontal dotted lines. (b) Separate selection of XRD patterns taken in correspondence of various cell voltages. The latter are shown here to further clarify how the structural transitions are manifested at various states of charge.

In the fully discharged (i.e., reduced) state, the diffraction pattern (Fig. 8c) appears strikingly similar to that of the pristine electrode. Hence, a LeBail fit in space group P21/n and unit cell parameters a = 10.178 Å, b = 7.393 Å, c = 6.998 Å and β = 90.185 ° provided a satisfactory fitting procedure. As complementary results to the previous in operando XRD measurements, these ex situ XRD analyses of the samples probed at different states of charge clearly confirm that the structural framework of MnHCM changes between a monoclinic structure and a cubic one during electrochemical cycling in a Li half-cell. This kind of structural evolution has been observed for Na+ or K+ ions during the charging processes in related half-cells.44,45

3.5.2 Ex situ XRD analysis A series of ex situ XRD measurements were performed to confirm the findings from the in operando experiment with the Li half-cell. The results of these additional measurements are displayed in Fig. 8a-c. Note that the characteristic diffraction peaks arising from the Al current collector (marked with grey circles in Fig. 8) were included in all the experiments. While this gives rise to strong peaks in each of the diffractograms, it can also be used as an internal standard, which is useful in terms of aligning the sample in a correct position, as well as providing


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The Journal of Physical Chemistry * [email protected]

If the initial monoclinic unit cell is expressed in terms of an equivalent cubic one, the cell parameter becomes √(b2+c2) ≈ 10.17 Å. Viewed in this manner, the overall expansion of the crystal structure, as it transforms from the monoclinic phase to the cubic one, is ≈ +16%. Finally, it should be noted that it was not possible to isolate the intermediate phase (suspected to be tetragonal) from the ex situ measurements alone. It appears as this intermediate phase is rather elusive, in the sense that it can only be observed during the electrochemical measurement, (i.e., in operando) and not after the cycling has been interrupted. Further studies are required to shed light on this peculiarity, while this finding highlights once more the potentialities and crucial advantages of in operando XRD as powerful means to probe subtle structural changes almost in real time. 4.

Author Contributions ‡V. R. and D. O. O. contributed equally to this work.

ACKNOWLEDGMENTS Viktor Renman and Dickson O. Ojwang acknowledge the ‘‘Consortium for Crystal Chemistry, C3 within the Röntgen Ångström cluster, Swedish Research Council VR, diary number: 2011-6512 for financial support. M. Valvo gratefully acknowledges the contribution from the ÅForsk Foundation (grant no. 18-317) and the Swedish Energy Agency (grant no. 2017013531). StandUp for Energy is also gratefully acknowledged for financial support.

Conclusions

REFERENCES

K2Mn[Mn(CN)6] particles were synthesized through a facile precipitation procedure under anaerobic conditions at room temperature. The material in question is primarily composed of earth abundant non-toxic elements, which are favorable in a wider economic and environmental perspective for possible application to electrochemical energy storage. The Mn[Mn(CN)6] framework can reversibly host Li+, Na+ and K+, as demonstrated by a series of electrochemical tests carried out in various alkali metal half-cells. By comparing the results of these different half-cells cycling at different rates, it becomes evident that their voltage polarization increases in the order Li+ < Na+ < K+. This may be partly attributed to the size dependency of cations insertion into PBA structures. Accordingly, the size of the solvated cations increases in the sequence Cs+ < Rb+ < K+ < Na+ < Li+.46 In terms of cycling stability and capacity retention, the Li half-cells appeared to display the best performances, although comparing these different types of cells is not straightforward. In fact, one should be careful in drawing too categorical conclusions, since the highly reactive Na and K metal anodes are likely to be detrimental with respect to long cycle life and voltage polarization, too. Via an advanced in operando XRD approach, it was possible to reveal that the structural framework undergoes a series of transitions during the electrochemical cycling in Li half-cells, namely from a monoclinic structure to a cubic one through an intermediate phase with apparent tetragonal features. These findings were also confirmed by means of complementary ex situ XRD analyses, which clearly demonstrated the phase transitions. However, such ex situ analyses could not outline conclusively the intermediate phase, which was clearly detected during the in operando measurements, instead.

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Supporting Information. Description of sample handling during various analyses. Plot of magnified IR spectral region.

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Corresponding Author

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and K-Ion Batteries

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