Electrodeposited Na2Ni[Fe(CN)6] Thin-Film Cathodes Exposed to

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C: Energy Conversion and Storage; Energy and Charge Transport 2

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Electrodeposited NaNi[Fe(CN)] Thin Film Cathodes Exposed to Simulated Aqueous Na-Ion Battery Conditions Philipp Marzak, Jeongsik Yun, Albrecht Dorsel, Armin Kriele, Ralph Gilles, Oliver Schneider, and Aliaksandr S. Bandarenka J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b00395 • Publication Date (Web): 04 Apr 2018 Downloaded from http://pubs.acs.org on April 4, 2018

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The Journal of Physical Chemistry

Electrodeposited Na2Ni[Fe(CN)6] Thin Film Cathodes Exposed to Simulated Aqueous Na-Ion Battery Conditions

Philipp Marzak,1 Jeongsik Yun,1,2 Albrecht Dorsel,1 Armin Kriele,3 Ralph Gilles,3 Oliver Schneider,4 Aliaksandr S. Bandarenka1,2,*

1 - Physics-Department ECS, Technical University of Munich, James-Franck-Straße 1, 85748 Garching, Germany

2 - Nanosystems Initiative Munich (NIM), Schellingstraße 4, 80799 Munich, Germany

3 - Heinz Maier-Leibnitz Zentrum (MLZ), Technical University of Munich, 85748 Garching, Germany

4 - Institute for Informatics VI, Technical University of Munich, Schleißheimerstraße 90a, 85748 Garching, Germany

* Corresponding Author e-mail: [email protected] (A. S. Bandarenka)

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Abstract Na-ion batteries have recently attracted great attention regarding their application in largescale energy storage systems (ESSs). Among different types of electrode materials for those classes of batteries, so-called Prussian Blue Analogues (PBAs) are among the very attractive ones due to their comparatively simple and low-cost methods of synthesis coupled with a promising cycle performance. In this study, one of the state-of-the-art PBA battery materials, namely electrodeposited Na2Ni[Fe(CN)6] (NiHCF) thin films, were tested under simulated battery conditions in aqueous and mixed (H2O/organic) electrolytes. Prolonged stability tests in aqueous electrolytes were performed together with in-operando electrochemical AFM monitoring. It is demonstrated that degradation of this material is not associated with noticeable morphological changes (mechanical stress), but is likely caused by changes in the chemical composition of the films. Intercalation and de-intercalation reversibility of Na+ and thin film stability in aqueous electrolytes appear to be not affected negatively by changes in the pH to values below 7. However, the films showed unstable behavior in basic media (pH > 10). The increase of the content of acetonitrile, which was used as an additive to simulate the influence of antifreezes in aqueous electrolytes, appears to primarily affect the de-intercalation of Na-ions in Na2SO4-based aqueous electrolytes. 1. Introduction The vision of future energy provision is nowadays shifting towards more sustainable and eco-friendly systems such as solar and wind power systems. Provision of energy by these means still needs to be improved in terms of reliability, as they depend on fluctuating natural phenomena. In order to address this issue, it is important to elaborate efficient energy storage systems (ESSs), which can compensate the gap between generation and consumption of energy in


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different periods and that can contribute to enhance grid stability1,2. Among various ESSs, secondary battery systems are promising, especially for grid-scale applications, due to their high energy conversion efficiency, and adequate energy and power density combined with relatively simple maintenance3. Prominent examples for such secondary batteries are Li-ion batteries, since they yield the highest energy density currently commercially available4. The technology of Li-ion batteries (LIB) has already been greatly developed and successfully implemented in, for example, the market of mobile devices and electric vehicles5. For large-scale applications, however, the focus does not solely lie on providing an optimal energy density. More importantly, it is set on utilizing cost efficient, safe and environmentally friendly systems6. The huge demand for lithium in LIB for portable devices and electric mobility have resulted in a three-fold increase in the price of Li-compounds within the last decade7. Besides the Li price, the estimated demand for automobile applications is about three times larger than the current annual Li production8. Also taking into account the geographical distribution of Li-production, a further and considerable increase in Li price in the future is apparent8,9. Therefore, it is unrealistic to rely for EES for grid-scale applications on Li-based battery systems only. Commercially successful Li-ion batteries utilize organic electrolytes that provide large operational potential windows and thus a high energy density10. However, the use of flammable and toxic organic electrolytes leads to an additional increase in the cost of safety systems and maintenance if being considered for upscaling. Resulting from these economic and ecological aspects together with problems associated with accumulation of Li this technology will be questionable to address the so-called “Terawatt Challenge” by 205011. Aqueous Na-ion batteries on the contrary do not show most of the above-mentioned drawbacks and, thus, present a promising alternative to Li-ion batteries for application in grid 3 ACS Paragon Plus Environment


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scale energy storage systems. Na is the 4th most abundant metal to be found on Earth and is evenly distributed throughout the earth's surface in various forms12. The current production of Na and its accumulation are at least one order of magnitude greater than that of Li13. The use of aqueous electrolytes has many advantages for ESS applications. For instance, aqueous electrolytes have higher ionic conductivity and are much cheaper, safer and more environmentally benign than many organic electrolytes14-17. Moreover, aqueous electrolytes are more resistant to rapid temperature changes within the system in comparison to common organic electrolytes. To give an example, the heat capacity of aqueous electrolytes is approximately three times larger than that of alkylene carbonates like Ethylene Carbonate (EC) and Propylene Carbonate (PC)18. This is expected to reduce fluctuations of the battery performance and maintenance costs. There is another important issue for the batteries to be used for the large-scale grid-line applications: it is important to overcome the ~500-1000 charge/discharge cycle limits typical for many state-of-the-art batteries. For large-scale installations, more than 10 years of operation is required. Taking into account the daily charge and discharge of the whole battery stack, this results in the requirement of more than ~3500 charge/discharge cycles without significant degradation. For this, a deeper understanding of the degradation mechanisms during the charge and discharge cycles is inevitable. In this work, the Prussian blue analogue (PBA) Na2Ni[Fe(CN)6] (NiHCF) - one of the stateof-the-art cathode materials for aqueous Na+ batteries as introduced by Wessells et al.19,20 - was investigated as electrochemically deposited thin film electrodes. These films showed a specific capacity of ~ 80 mAhg-1 with high charge and discharge reversibility, even at a rate of 180C21. The thin films were exposed to simulated battery conditions in order to further elucidate the


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stability and performance of the chosen electrode material and the interaction with different solvents and electrolyte compositions21-24. The results strongly indicated that the degradation mechanism does not originate from mechanical or morphological alterations but rather involves chemical changes in the films. It is also demonstrated how antifreeze additives might influence the reversibility of Na+ cation (de-)intercalation in SO42- containing electrolytes by predominantly affecting the de-intercalation mechanism instead of both intercalation and de-intercalation of Na+. 2. Experimental Section Electrochemical experiments were performed in a three-electrode glass cell setup using a Bio-Logic VSP-300 potentiostat. Electrode potentials were referred to a Ag/AgCl reference electrode (SSC, 3 M KCl, SI Analytics, “B 3420+”) and a Pt wire was used as a counter electrode. Arrandee™ Au films and AT-cut Au quartz crystal wafers (Stanford Research Systems, Ti adhesive layer) were used as working electrodes. The Na2Ni[Fe(CN)6] thin films were cathodically electrodeposited from an aqueous solution containing 0.5 mM NiCl2 · 6 H2O (99.3 %, Alfa Aesar), 0.5 mM K3Fe(CN)6 (99 %, Sigma Aldrich) and 0.25 M Na2SO4 (≥99 %, Sigma Aldrich). For the deposition, the electrode potential was cycled between 0.0 V and 0.9 V vs. SSC at 50 mV/s for ~60 times. For the characterization of the films, cyclic voltammetry measurements of the NiHCF thin films were carried out by cycling the electrode potential from 0.1 V to 0.8 V vs. SSC at 50 mV/s using different aqueous electrolytes: 0.25 M Na2SO4 and 0.25 M Na2SO4 with pH adjusted to pH = 2 and 10, as well as 0.25 M Na2SO4, 0.1 M Na2SO4 and 0.2 M NaClO4 · H2O (≥98.0 %, Sigma Aldrich) each with varying mole fractions χACN ∈ {0.00; 0.06; 0.10; 0.15} of acetonitrile (ACN, 99.8 %, Sigma Aldrich). The pH value was adjusted by adding H2SO4 (ROTIPURAN®, 96 %, Carl Roth) or NaOH (>98 %, Sigma Aldrich) and controlled with a digital pH meter (OMEGA™ 5 ACS Paragon Plus Environment


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PHH22). All solutions were produced using ultrapure water (18.2 MΩ, Evoqua, Germany) and purged with Ar gas (5.0, Westfalen AG) for about 15 minutes before experiments. In order to investigate the change of lattice parameters of NiHCF upon changes in the state of charge, ex-situ XRD measurements were performed on de-intercalated and intercalated NiHCF thin films at the Materials Science Laboratory (operated in cooperation of TU München and Helmholtz Zentrum Geesthacht) of the Heinz Maier-Leibnitz Zentrum. For the XRD measurements, thicker films were prepared by increasing the number of deposition cycles (on an Au coated quartz resonator) to 475 in order to obtain a sufficient amount of material (estimated thickness of NiHCF thin film: ~ 250 nm) and thereby increase the intensity of the XRD measurement. A PANalytical Empyrean® high-resolution powder diffractometer was used with Mo-Kα radiation source (mean wavelength λMo= 0.7107 Å, energy EMo=17.4 keV) for XRD measurements. The instrument is equipped with a 1D linear X’Celerator real time multistrip detector with an efficiency of 30 % for Mo Kα1 radiation and a pixel resolution of 0.002° in 2θ. In order to further enhance the signal of the thin NiHCF film the XRD measurements were performed at a fixed gracing incident beam angle of 3° (to achieve a small penetration depth of the X-ray beam and a long X-ray path in the film) whereas only the detector was moved along the 2θ-axis with an uncertainty of 0.008°. The beam width was set to 10 mm with an incoming beam divergence of 1/8° in order to illuminate the whole NiHCF film. Axial divergence was reduced by inserting Soller collimators with an opening of 0.04 rad at the incident and reflected beam path. Considering the x-ray optic settings and tilt of the sample an area of 12 mm in width and 16 mm in length respectively was irradiated by the x-ray beam. A 75 µm Zr filter was used to suppress Kβ radiation of Mo radiation below 0.01 Kβ/Kα. A total measurement time of 66 h (1.1 min per step) for each measurement was required to achieve reliable data. Sample spinning at 4


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rpm was applied to improve statistics. Rietveld refinement procedures adapted for 2θ scans were applied for the crystal structure analysis using HighScore Plus™ software. In-operando electrochemical AFM (EC-AFM) measurements were performed in a Veeco MMTMEC fluid cell using a Veeco MultiMode V AFM with a NanoScope 3D controller and a NanoScope Universal Bipotentiostat. A gold wire was used as a counter electrode and a Ag/AgCl electrode consisting of a Ag wire coated with AgCl that reaches into the fluid cell was used as a quasi-reference electrode. Utilization of the quasi Ag/AgCl reference electrode in the EC-AFM fluid cell results in a shift of 0.3 V vs. SSC = 0 V vs. quasi Ag/AgCl for CVs performed in the EC-AFM setup. 0.25 M Na2SO4 electrolytes for the measurement were deoxygenated and then injected into the fluid cell for the measurements. X-ray photoelectron spectroscopy (XPS) was performed in an ultra-high vacuum (UHV) steel chamber utilizing a SPECS XPS spectrometer (SPECS, Germany). The used XR50 X-ray tube employed an Al anode (12 kV, 200 W). The characteristic Kα line of the X-ray radiation spectrum with a photon energy of 1486.61 eV irradiated the sample on an area of 1.4 mm × 4 mm. The kinetic energies of the electrons were detected with a semi-spherical electron energy analyzer (PHOIBOS 150 2D CCD) operating with a pass energy of 20 eV. 3. Results and Discussion 3.1. Ex-situ XRD Measurements Significant expansion and shrinkage of the cathode upon charge and discharge is regularly observed for commonly used intercalation materials in Li-ion batteries. These lead to severe mechanical degradation of the material26-28. In contrast, one might expect that no significant expansion or shrinkage will take place upon (de-)intercalation of Na+ from or into NiHCF



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considering that the crystal structure of NiHCF offers intercalation sites with a radius of ~1.6 Å10 and sodium ions have a radius of only ~1.02 Å29. The sharp, well defined NiHCF peaks in the samples’ characteristic XRD diffractograms (Fig. 1A and 1B) show that these films are crystalline. The XRD peaks matched the literature and ICSD data base well19,30,31. Despite using a fixed grazing incident angle of 3° diffraction peaks of the subjacent gold electrode and quartz substrate overlay most of the NiHCF peaks. Only the (200) NiHCF peak at 8.2° was appropriate for analysis because of its high intensity. The XRD data have been reduced by masking the predominant SiO2 substrate peak and by background correction (Fig. 1A and 1B). The (111) and (200) Au-peaks were used to refine the sample displacement and detector offset32. The peak related to Ti(100), which originates from the Ti adhesion layer between Au and quartz, can be seen in the figure as well. Fig. 1C shows good match of both Au-peaks for the de-intercalated and intercalated sample in contrast to the shift of the (200) NiHCF peak that is due to Na+ intercalation (Fig. 1D). The calculated lattice parameters (a, b, c) of the NiHCF thin films results in 10.01 Å for the deintercalated and 10.08 Å for the Na+ intercalated film, respectively. The fitting of the XRD data was performed using a cubic phase Fm-3m Na-NiHCF space group from the 4091 ICSD data base. The volume expansion after Na+ intercalation is in the order of 2.2 %. The calculated lattice parameter of the NiHCF films differs from values reported in literature (in the order of 10.2 Å)19,29 which can be explained both by the underlying Au-layer which has a strong preferred orientation in (111) direction and the sample preparation by means of cathodic electrodeposition in contrast to co-precipitation as performed by Wessells et al.19 Table 1 shows the peak parameters obtained by the Rietveld refinement.


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Figure 1. Reduced X-ray diffractogram of (A) de-intercalated and (B) intercalated NiHCF thin films. (C) Whereas the XRD peaks correlated to the Au substrate match well for both samples, (D) the (200) NiHCF peak is shifted upon Na+ (de-)intercalation.



Pos. [°2θ] NiHCF (200)

Au (111)

Au (200)

Ti (100)

d-spacing [Å]

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Height [cts]

FWHM [°2θ]

Unit cell

de-intercalated

8.1263

5.0059

13985

0.0933

10.009

intercalated

8.1725

4.9777

16994

0.0867

10.081

de-intercalated

17.7109

2.3041

22648

0.2485

intercalated

17.7314

2.3015

31193

0.2509

de-intercalated

20.4600

1.9972

7517

0.2787

intercalated

20.4853

1.9947

8492

0.2574

4.006

de-intercalated

15.758

2.5875

5108

0.1574

4.8002

intercalated

15.776

2.5846

7053

0.1639

4.8001

Table 1. The peak parameters obtained by the Rietveld refinement

3.2. In-situ Electrochemical AFM Stability Measurements In-operando electrochemical AFM measurements were carried out in order to examine the morphology of the NiHCF electrode during sodium (de-)intercalation during cycling and accelerated aging and to confirm the absence of severe volume changes. Fig. 2A shows the CVs of a NiHCF thin film in a 0.25 M Na2SO4 electrolyte recorded in the EC-AFM. Degradation of the thin films was accelerated by cycling to potentials much higher (1.1 V vs. quasi Ag/AgCl) than necessary for de-intercalation, in a range where additional electrode reactions are apparent in the CV. The anodic and cathodic peaks recorded around 0.15 V vs. quasi Ag/AgCl correspond to the de-intercalation and intercalation of Na+, respectively. A drastic current drop of these peaks after 45 cycles (blue line) compared to the second cycle (black line) demonstrates the artificial acceleration of degradation. The first EC-AFM images of a NiHCF thin film recorded at constant potentials in the deintercalated state (0.55 V vs. quasi Ag/AgCl) and the intercalated state (-0.2 V vs. quasi Ag/AgCl)


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Figure 2. EC-AFM characterization of Na2Ni[Fe(CN)6] thin films. (A) Cyclic voltammograms of sodium intercalation and de-intercalation. Rather positive upper potential limits (1.1 V vs. quasi Ag/AgCl) were applied in order to accelerate the degradation. AFM images of the Na2Ni[Fe(CN)6] thin film (B) at constant potential after Na+ de-intercalation (0.55 V vs. quasi Ag/AgCl), (C) at constant potential after Na+ intercalation (-0.2 V vs. quasi Ag/AgCl), (D) during CV starting at the 2nd cycle and (E) during CV during the last two cycles. A morphological change of the film was neither observed during the Na+ intercalation and de-intercalation nor after degradation of the film.

are shown in Fig. 2B and 2C, respectively. By comparing the two AFM images no significant morphological change in grain size or structure can be detected. Slight differences most likely result from noise and drift during the measurements. These findings are in good accordance with the results of XRD measurements. Fig. 2D and 2E show the EC-AFM images of the NiHCF thin films in a 0.25 M Na2SO4 electrolyte taken immediately after start of the second CV cycle and of the 45th cycle, respectively. Here the potential was cycled continuously between -0.2 V and 1.1 V vs. quasi Ag/AgCl during image acquisition, one CV cycle taking 52 s and one image taking 128 s to acquire. Despite the significant decrease in the peak current density after repeated cycles, morphology or grain size changes of the films are barely visible. Despite continuous potential scanning, these images match each other very well. Therefore, morphological changes of the NiHCF thin films caused by a volume expansion and shrinkage upon intercalation and de11 ACS Paragon Plus Environment


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intercalation of Na+ that might result in cracking of the material and mechanical degradation are apparently not the source for the loss of the cathodes’ electrochemical performance during repeated cycling. Supposedly, the degradation of NiHCF thin films upon repeated cycling instead results from chemical degradation of the electrode material. 3.3. Electrochemical Experiments on Cycling Stability Cyclic voltammetry performed using NiHCF thin films in pH-modified 0.25 M Na2SO4 electrolytes yields further evidence on the underlying mechanism of the electrode’s chemical degradation. As shown in Fig. 3A the stability of the film in the pH = 2 electrolyte does not significantly differ from the unmodified 0.25 M Na2SO4 electrolyte over 10 cycles. The subsequently recorded voltammograms match another very well. In contrast, in the electrolyte with pH 10 one sees alterations in the peak positions and shapes already within the first 5 cycles. These changes become more pronounced during the course of the recorded following 20 cycles (see Fig. 3B). Especially the peak separation increases, indicating a less reversible (de-)intercalation process and impaired reaction kinetics connected to the chemical decomposition of the film. These findings are in accordance with literature where the electrode material has been reported to show an increased stability in acidic media, whereas it tends to decompose in basic media33-36. Measurements in 0.1 M NaOH aqueous electrolyte further indicate that NiHCF gradually reacts to give NiOx during charging and discharging in aqueous media when interaction with OHtakes place. Cyclic voltammograms in such an aqueous sodium hydroxide solution are shown in Fig. 3C. One can clearly notice the evolution towards a typical NiOx characteristic voltammogram in this medium32,33. Additionally, visual characterization and XPS analysis of a NiHCF thin film electrode before and after cyclic voltammetry in 0.1 M NaOH show that the 12 ACS Paragon Plus Environment

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film decomposes (Fig. 3D and 3E). First, within the course of this experiment the color of the thin film electrodes changes

Figure 3. NiHCF film stability evaluation. Electrochemical behavior of the Na2Ni[Fe(CN)6] thin films in (A) 0.25 M Na2SO4 aqueous electrolytes without additive (pH = 7, black line) and at pH 2 (red dashed line) (10 cycles each), (B) in 0.25 M Na2SO4 aqueous electrolyte at pH 10 (cycles 1, 5, 10, 15 and 20) and (C) in 0.1 M NaOH. Fig. 3C also depicts one cycle of an initial voltammogram of the thin film electrode in 0.25 M Na2SO4 aqueous electrolyte (dashed black line), which evolves to the voltammogram of NiOx. (D,E) Photographs and XPS analysis of the NiHCF thin films before (left, goldish electrode and solid black lines) and after (right, blackish electrode and dashed blue lines) cyclic voltammetry in 0.1 M NaOH.

from a goldish color - which is characteristic for NiHCF thin films deposited on gold - to a characteristic blackish color of an oxidized NiOx film when the cyclic voltammetry is stopped after the oxidation half wave. Second, as seen in the XPS analysis, the reaction takes place to such an extent, that the Fe centers of the coordination complex and Na-ions are displaced 13 ACS Paragon Plus Environment


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completely. The Ni- and O-peaks, however, remain very pronounced, indicating that the film completely converts to NiOx. One can thus draw the conclusion, that the chemical degradation of NiHCF thin films upon repeated charging and discharging in aqueous electrolytes bases on a reaction mechanism involving OH-. In general, the cathode material is operated at very anodic potentials as referred to the hydrogen evolution reaction (0 V vs. SHE = -0.21 V vs SSC and the mean value of the half wave potentials for Na+ (de-)intercalation from and into NiHCF is 0.42 V vs. SSC). In this potential regime, the adsorption of OH- is favored. It has been shown in research on metal and metal oxide based catalysts, that lattice strain additionally affects the adsorption strength of OH- onto the materials surface and thereby influences the rate of reactions involving this adsorbate37-42. Concluding from experiments with alkaline electrolytes, demonstrating the decomposition of NiHCF to NiOx, it appears that an interaction between Ni2+ in the film and OH- is preferred. The change in lattice parameters in the order of 0.7 % upon (de-)intercalation of Na-ions from or into NiHCF, as identified by means of XRD, will not result in a mechanical degradation of the films. However, similar to the findings in case of metal and metal oxide based catalysts, the nevertheless connected volume expansion or shrinkage will result in a change of the lattice stress, hence change the adsorption strength of OH- and will, as a result, influence the reaction rate of NiHCF with OH-. The following reaction mechanism can be suggested based on the findings presented so far: 1.

OH- adsorbs onto the surface of the NiHCF electrode, presumably at sites

exposing nickel to the electrolyte. Upon intercalation of Na+, which is driven by the reduction of Fe(CN)63- to Fe(CN)64-, the surface binding of OH- is strengthened so that the likelihood of a reaction of nickel with this adsorbate is increased.


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Figure 4. Proposed chemical decomposition of electrodeposited Na2Ni[Fe(CN)6] thin films upon (de-) intercalation of Na+ from and into the electrode material. In order to simplify the graphics, Na+ cations are not depicted in the shown section of the electrodes crystal lattice. They usually are positioned at the interstitial sites of the cubic lattice structure. Red lines or arrows symbolize which species are interacting in general. (A) OH- adsorbs onto the surface of the NiHCF electrode, presumably at sites exposing nickel to the electrolyte (blue balls, highlighted in green). In the case of intercalation of Na+, which is driven by the reduction of Fe(CN)63- to Fe(CN)64-, the surface binding of OH- is strengthened so that the likelihood of a reaction of nickel with this adsorbate is further increased. (B) As a result of the reaction of nickel with OH-, the Fe(CN)63-/4- is extracted from the film completely. Alternatively, by exchanging CNligands of this complex with OH-, the complex is destabilized and the iron centers (red ball, highlighted in green) are extracted. (C) By reducing the number of active centers (iron centers), which are the source of the driving force for (de-)intercalation of Na-ions, the electrode becomes increasingly inactive. The reaction product of nickel with OH- still remains on the samples surface as a residue. Importantly, this process is expected to take place stochastically and locally, so that initially defects in the crystal structure are introduced instead of the electrode’s crystallography being altered.



2.

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As a result of the reaction of nickel with OH-, “Fe(CN)63-/4-” is either extracted

from the film completely, or by exchanging CN- ligands of this complex with OH-, the complex is destabilized and the iron centers are extracted. By reducing the number of active centers (iron centers), which are the source of the driving force for (de-)intercalation of Na-ions, the electrode becomes increasingly inactive. The reaction product of nickel with OH- still remains on the samples surface as a residue. A visualization of this suggested decomposition mechanism is depicted in Fig. 4. It is important to point out, that this degradation mechanism will most likely take place locally and stochastically. This is, even by changing the local atomic composition of the NiHCF through this mechanism, the general crystal structure will not immediately be altered completely. These changes rather appear as local defects. As such, Fig. 4 represents only a small section of the electrode’s crystal structure and does not imply an immediate and complete change of the crystallography. If one compares the different experiments in pH-modified and pure 0.25 M Na2SO4 electrolytes as well as aqueous 0.1 M NaOH with a focus on the decomposition reaction, one might observe that the kinetics of the decomposition reaction increases with the OHconcentration in the solution. One can thus conclude, that Na2Ni[Fe(CN)6] thin film electrodes will behave more stable in aqueous media the less they are exposed to OH-. However, even if this process is very slow when the activity of OH- is greatly reduced, it will gradually take place. 3.4. What is the influence of an antifreeze compound? Acetonitrile (ACN) is chosen as a reference solvent for modeling the influence of antifreeze-like organic additives for multiple reasons: importantly, it shows electrochemical stability within the investigated potential window. Furthermore, it is mixable with water without solubility limits and the interactions of water with ACN have been investigated thoroughly43-45. 16 ACS Paragon Plus Environment

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Water can form hydrogen bonds with ACN, even though they might not be as strong as H2O-H2O hydrogen bonds. (De-)intercalation experiments with NiHCF have already been performed utilizing electrolytes based on pure ACN41. It is thus safe to assume that NiHCF is sufficiently stable towards ACN and that the system behaves reversible. Cyclic voltammetry performed with electrolytes based on a mixture of water with acetonitrile have been performed with two different solutes, Na2SO4 and NaClO4. The mole fraction of acetonitrile was varied as well and chosen to be χACN ∈ {0.00; 0.06; 0.10; 0.15}. Integration of the cyclic voltammogram in Fig. 5A leads to the charge vs. electrode potential plot in Fig. 5B, which helps in visualizing the differences in half wave potentials, peak positions and peak separation. For a 0.25 M Na2SO4 solution one can clearly identify a correlation between an increasing mole fraction of χACN and a less reversible (de)intercalation process as measured by the difference in half wave potentials ∆E1/2 (refer to Fig. 5A - 5D). Error bars included in these graphs result from the experimental procedure. Cyclic voltammetry of each NiHCF thin film was first performed in pure 0.25 M Na2SO4 solution for characterization in a reference system and then repeated in one electrolyte containing a varied mole fraction χACN. The measurement for χACN = 0.15 was repeated several times. The experimental procedure is overall very reproducible. A decrease of electrolyte conductivity would surely increase the electrochemical cell’s uncompensated resistance, increasing the peak separation, but would result in a symmetrical shift of the peaks. In contrast to this, mainly the half wave potential for the oxidation half wave (de-intercalation) is shifted towards more positive potentials. This rather indicates that primarily the de-intercalation of Na-ions from NiHCF is affected negatively by adding more acetonitrile. The same trend is observed for a lower



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concentration of 0.1 M Na2SO4. The reversibility of Na+ (de-)intercalation from and into NiHCF is, however, in general worse due to the decreased concentration as reported in literature21. A similar experiment performed in 0.2 M NaClO4 electrolyte with different mole fractions χACN shows different results. Here one cannot observe a comparable decrease in reversibility between χACN = 0.00 and χACN = 0.15. How the mole fraction of ACN affects reversibility in these different scenarios is compared by means of half wave potentials in Fig. 6A to 6D. The difference in the dependence of reversibility on χACN between Na2SO4 and NaClO4 electrolytes might result from the different solubility of the two salts in Acetonitrile. While the solubility of NaClO4 is sufficient in ACN41, it is very low for Na2SO4, if this salt is soluble at all. The resulting competition between Na2SO4 and ACN for water molecules might negatively affect de-intercalation of Na-ions. The hydration shell of Na+, which is rejected during intercalation31, is impaired in its formation upon de-intercalation and rehydration of Na-ions. First, this is because the ACN molecules hinder the formation of the hydrogen-bond network in the solvation shell. Second, since Na2SO4 is insoluble in ACN, a solvation shell composed of both H2O and ACN is strongly unfavored and local separation of water and ACN enforced.


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Figure 5. Electrochemical characterization of the Na2Ni[Fe(CN)6] thin films in aqueous electrolytes modified with acetonitrile. (A) Cyclic voltammograms of the films in a 0.25 M Na2SO4 aqueous electrolyte with varying mole fraction of acetonitrile (ACN) in the solvent: χACN ∈ {0.00; 0.06; 0.10; 0.15} and (B) corresponding charge and discharge curves obtained by integrating the CVs in (A). (C) The half wave potentials for intercalation and de-intercalation as a function of mole fraction of ACN. Solid triangles denote mean values with standard deviation where a statistic is available. Open symbols denote single measurements. (D) Potential shifts between Na+ intercalation and de-intercalation at the half wave potentials (∆E1/2) as a function of the mole fraction of ACN.



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Figure 6. Electrochemical characterization of the Na2Ni[Fe(CN)6] thin films in aqueous electrolytes enhanced with Acetonitrile. (A) Cyclic voltammograms of the films in 0.1 M Na2SO4 and 0.2 M NaClO4 aqueous electrolyte with varying mole fraction of Acetonitrile (ACN): χACN ∈ {0.00; 0.15} and (B) corresponding charge and discharge curves obtained by integrating the CVs in (A). (C) The half wave potentials for intercalation and de-intercalation as a function of mole fraction of ACN. Solid triangles denote mean values with standard deviation where a statistic is available. Open symbols denote single measurements. (D) Potential shifts between Na+ intercalation and (de-)intercalation at the half wave potential (∆E1/2) as a function of the mole fraction of ACN.


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4. Summary and Conclusions The perspective goal of this work was to gain insights into the mechanisms governing the long-term stability of cathodically electrodeposited Na2Ni[Fe(CN)6] thin film electrodes for aqueous sodium ion batteries and to further elucidate the materials performance under conditions relevant for field applications. These are steps on the path to identify crucial cell parameters and characteristics that need to be considered when building a functional full device that can be applied in real world scenarios. First, the degradation mechanism upon repeated charging and discharging of the electrode material, that governs the long-term stability, has been shown not to result from mechanical degradation. The results of two independent methods as XRD analysis to determine lattice parameters in dependence of the state of charge and in-operando EC-AFM providing direct images of Na2Ni[Fe(CN)6] thin film are in good agreement, that volume expansion and shrinkage of NiHCF upon Na-ion (de-)intercalation is marginal in contrast to common Li-ion battery cathode materials. Electrochemical experiments with pH modified 0.25 M Na2SO4 aqueous electrolytes and in 0.1 M NaOH together with XPS analysis instead support the hypothesis that the degradation mechanism in aqueous media is of a chemical nature. Supposedly, the presence of OH- in the electrolyte plays a crucial role, as it appears that Ni2+ in the electrode reacts with OH- to NiOx. Thereby, the iron centers - that represent the source of the driving force for (de-)intercalation - are extracted from the material and the cathode loses performance as it continuously degrades during operation. With increasing OH- concentration this reaction might be accelerated, which is in agreement with findings in literature that NiHCF behaves less stable towards basic media. Second, the focus was set onto the influence of acetonitrile (ACN) as a model antifreeze-like organic additive. It is apparent that such an additive can have a noticeable



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negative influence on the performance and reversibility of the electrode operation. It is indicated that the solubility of the electrolyte salt plays a crucial role. While experiments in aqueous NaClO4 solution do not reveal a significant decrease of reversibility with increasing amount of ACN, primarily the de-intercalation process in Na2SO4 appears to be increasingly impaired when the mole fraction χACN of ACN present in the electrolyte rises. It is believed, that the rebuilding of a new hydration shell around Na+ upon de-intercalation is impaired by the presence of ACN, especially because Na2SO4 is, in contrast to NaClO4, insoluble in ACN. One might think of generally operating NiHCF and comparable PBAs in acidic media in order to slow down the degradation of such cathode materials through the suggested decomposition mechanism and thus increasing the lifetime of battery systems. However, it is questionable if this can be realized in the design of a full cell device where anode materials used are potentially sensitive to acidic media. As an alternative, one might identify additives and electrolyte compositions that hinder adsorption of OH- onto the electrodes surface or even the autoprotolysis in the first place46,47. However, care must be taken not to affect the performance of the battery materials negatively. Experiments comprising ACN as a model additive demonstrate, that the interplay of all electrolyte components first needs to be understood before one can rationally design the battery system to fulfill the requirements for grid scale ESS applications. Acknowledgements Financial support from the cluster of excellence Nanosystems Initiative Munich (NIM) is gratefully acknowledged. Jeongsik Yun is thankful for the financial support from Nagelschneider Stiftung. Conflict of interests The Authors declare no conflict of interests. 22 ACS Paragon Plus Environment

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TOC Graphic.


Electrodeposited Na2Ni[Fe(CN)6] Thin-Film Cathodes Exposed to

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