Kinetics and Degradation Processes of CuO as Conversion Electrode

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Kinetics and Degradation Processes of CuO as Conversion Electrode for Sodium-Ion Batteries: An Electrochemical Study Combined with Pressure Monitoring and DEMS Franziska Klein, Ricardo Pinedo, Balázs B. Berkes, Jürgen Janek, and Philipp Adelhelm J. Phys. Chem. C, Just Accepted Manuscript • Publication Date (Web): 23 Mar 2017 Downloaded from http://pubs.acs.org on March 23, 2017

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The Journal of Physical Chemistry

Kinetics and Degradation Processes of CuO as Conversion Electrode for Sodium-Ion Batteries: An Electrochemical Study Combined with Pressure Monitoring and DEMS Franziska Klein1, Ricardo Pinedo1, Balázs B. Berkes2, Jürgen Janek1,2 and Philipp Adelhelm*3,4 1

Justus-Liebig-University Giessen, Institute of Physical Chemistry, Heinrich-Buff-Ring 17, 35392 Giessen, Germany

2

Karlsruhe Institute of Technology, Battery and Electrochemistry Laboratory, Hermann-von-Helmholtz-Platz 1, 76344 Eggenstein-Leopoldshafen, Germany

3

Friedrich-Schiller-University Jena, Institute for Technical Chemistry and Environmental Chemistry, Philosophenweg 7a, 07743 Jena, Germany 4

Center for Energy and Environmental Chemistry Jena (CEEC Jena), Philosophenweg 7a, 07743 Jena, Germany

*Corresponding author: [email protected]

Abstract Copper oxide (CuO) can be used as electrode material for lithium-ion and sodium-ion batteries, however, the way toward application for rechargeable systems is still long. This is mainly related to the complex nature of the electrode reaction and to ageing mechanisms that are not well understood, especially in case of sodium. Main subject of this paper is to compare the electrode reaction of CuO in lithium (CuO/Li) and sodium (CuO/Na) half cells and to study side reactions in the CuO/Na system by means of differential electrochemical mass spectrometry (DEMS) and in-situ pressure monitoring during galvanostatic cycling (PMGC). Electrode processes have been studied at different current densities and temperatures. In CuO/Li cells, CuO and Cu2O form during charging, their respective fraction depending on the current density. In case of CuO/Na, Cu2O is always the charging product although oxidation to CuO can be temporarily achieved by increasing the temperature to 50 °C. The difference between both, CuO/Na and CuO/Li is related to the electrode volume expansion/shrinkage during cycling. Differences in the temporal evolution of electrode surface films are followed by electrochemical impedance spectroscopy (EIS). For the CuO/Na system, PMGC and DEMS studies reveal a periodic release of gaseous side products as a result of a slow but likely continuous electrolyte degradation. This degradation is due to the repeated formation of a surface film (discharge) and its partial dissolution (charge) accompanied by the release of H2 and CO2. The degradation processes fade during cycling but remain dynamic. This means that long cycle life of CuO electrodes in sodium-ion batteries can likely be only achieved by employing some excess electrolyte.

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Introduction The development of safe and low cost electrochemical energy storage devices with high energy densities is a key challenge for current battery research. Furthermore, alternatives to conventional lithium-based systems should be investigated due to limited resources of lithium and/or other cell components that might become an issue in the long term. Therefore, charge storage based on conversion reactions (MaXb + (bc)Li ↔ aM + bLicX with M being a transition metal and X a non-metal, e.g. O, S, F, P, N…) is a promising concept, potentially providing higher specific capacity and energy density compared to classical LIBs. In the last 10 to 15 years, a large number of studies have been devoted to conversion reactions with lithium 1,2 and more recently, by sodium-based systems are also studied an attractive approach towards batteries relying on abundant elements 3,4. From the scientific view, it is interesting to investigate the effect of the alkali ion size on the electrode reaction. Several studies on conversion reactions for sodium-ion batteries have been published so far 4–16, the focus being mostly on performance improvement (e.g. capacity, cycle life) by nanostructuring, for example 17–33. However, many scientific questions still remain under debate for sodium and but also for lithium-based systems 1. Every conversion electrode has its own story, but most of them share questions that are related to i) the first lithiation/sodiation step which can be understood as an electrode activation, ii) the complexity of the reaction mechanisms during cycling, iii) high overpotentials and voltage hysteresis, iv) severe volume changes and v) degradation processes such as surface film formation and electrolyte degradation combined with gas evolution. (i) The initial lithiation/sodiation (activation) of the electrode is linked to the formation of a nanoscopic structure with metal nanoparticles (M) embedded in an amorphous matrix of LicX or NacX, respectively. This nanostructure is the key for the reversibility of conversion reactions because the often poorly conductive phases LicX or NacX are in close contact with the metal nanoparticles. The initial lithiation/sodiation is only partially reversible resulting in a poor initial coulombic efficiency (ICE) usually between 50 – 75 % 1,8,9. (ii): Conversion reactions undergo complex reaction processes with often only partial conversion during cycling 12,34–36. Incomplete recharging leads to lower oxidation states (i.e. Cr2O3 CrO 37, Co3O4 CoO 38,39 and CuO Cu2O 5,13,40), for example. Moreover, the role of the SEI during cycling remains often unclear. Nevertheless, values for the coulombic efficiency quickly increase after the first cycle often approaching 100 % 1. (iii): An important challenge of conversion reactions are high overpotentials and large voltage hysteresis. The direct and complete conversion is kinetically not favoured as new (and often non-conductive) phases form that require significant structural rearrangement. Moreover, intermediate phases, including insertion compounds, form that complicate the situation. It is therefore no surprise that the measured redox potentials often significantly deviate from the expected ones. Unfortunately, higher temperatures and nanostructuring of electrodes so far do not decrease these overpotentials sufficiently enough 41–43. The resulting low round trip energy efficiency is a serious drawback toward application in practical batteries. (iv): Volume expansions are a common challenge for all conversion reactions. Large volume changes during cycling can lead to mechanical degradation of the electrode and poor cycle life. This volume expansion is more severe for sodium and roughly double compared to anaACS Paragon Plus Environment

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logue lithium systems 4. The larger volume expansion in case of sodium might also lead to more severe cracking which increases the electrolyte/electrode contact area. This larger contact area might ease the conversion reaction but on the other hand could also enhance electrolyte degradation. (v): At potentials below < 1 V vs. A/A+ common electrolytes decompose and form a solid electrolyte interphase (SEI) during discharge 40. This surface film has been found to (partially) disappear during charging44. The mechanisms and consequences of this process are, however, poorly understood. A variety of solid, dissolved and gaseous electrolyte decomposition products has been identified by XPS, NMR, FTIR and mass spectrometry (MS) analyses 45–49, for example . Gachot et al. observed CO, CO2 and C2H4 as gaseous electrolyte decomposition products by gas chromatography (GC)/MS when cycling CoO against Li at 55 °C in a carbonate electrolyte, for example. In situ monitoring of side reactions is certainly most appealing and can be obtained by differential electrochemical mass spectrometry (DEMS), for example, which provides qualitative (and eventually quantitative) information on gaseous side products that evolve during cell cycling. A simpler but yet effective method is to monitor the cell pressure during cycling using a pressure gauge. Pressure monitoring during galvanostatic cycling (PMGC) was recently used in the field of metal-air batteries 50. To the best of our knowledge, this work combines for the first time PMGC and DEMS analysis to study gas evolution in Na-based conversion systems. It is important to note that, despite the valuable information both methods provide, gaseous compounds can be generated at both electrodes or due to a crosstalk between them. Thus, to decide whether gas release occurs at the electrode of interest or also at the counter electrode (sodium or lithium) is challenging. Anyway, He et al. showed for measurements on Li4Ti5O12 and Jozwiuk et al. in Li/S-batteries that the contribution of the metal electrode to the DEMS signal is comparably small 51,52. For this study we chose copper oxide (CuO) as electrode material, as it is considered as an attractive electrode material for LIBs and SIBs due to its low cost, environmental friendliness and high theoretical capacity (674 mAh g−1) 5,9,13,40,53–74 and as it has already been applied in primary cells 75–81. Table 1 summarizes some general characteristics of the CuO electrode reaction with Na and Li. However, due to the issues discussed above, the electrode reaction is much more complex and requires detailed characterization. Recently, we studied the reaction mechanism of CuO with lithium (CuO/Li) and sodium (CuO/Na) as well as the concomitant surface film formation using sputtered CuO thin films as model electrode5. In contrast to CuO/Li for which it is known that (at least to a larger extent) CuO forms during cycling, we found that the conversion in the CuO/Na system is only partially reversible and Cu2O forms during cycling. Interestingly, oxygen also appeared to be redox active in the case of CuO/Na, indicating also anion redox activity. Morphological and chemical composition analyses of the surface films revealed important differences between both systems. For sodium, the SEI formed during discharge is significantly thicker (6 µm vs. 1 µm), more homogeneous and mainly composed by inorganic species (Na2CO3, NaF). For CuO/Li, the SEI was more heterogeneous and organic. Overall, our previous study on CuO thin films demonstrated that the reaction mechanism and the surface film formation are strongly influenced by the size of the alkali ion 5. ACS Paragon Plus Environment


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Different from our earlier publication on thin film electrodes, this study makes use of conventional CuO bulk electrodes (thickness approx. 70−100 µm) as they are closer to practical relevance and because the larger amount of material used in the cell enables the application of pressure monitoring and DEMS as analytical techniques. Moreover, we use galvanostatic cycling at different rates and temperatures, impedance spectroscopy and microscopy as classical characterization methods to compare between the CuO/Li and CuO/Na systems. Table 1. Thermodynamic data from HSC Chemistry 8.0 (p = 105 Pa, T = 25 °C, bulk thermodynamics, pure phases)a Lithium

∆rG0 / kJ mol−1

E0 vs. Li/Li+ /V

qth / mAh g−1

volume change /%

wth / Wh kg−1

(1)

 → Cu + Li2O CuO + 2Li ← 

− 431.551

2.24

674

+ 74

1285

(2)

 → Cu2O + Li2O 2CuO + 2Li ← 

− 449.800

2.33

337

+ 54

722

(3)

 → 2Cu + Li2O Cu2O + 2Li ← 

− 413.302

2.14

375

+ 22

731

0

Sodium

∆rG0 / kJ mol−1

E vs. Na/Na+ /V

qth / mAh g−1

volume change /%

wth / Wh kg−1

(4)

 → Cu + Na2O CuO + 2Na ← 

− 246.705

1.28

674

+ 173

546

(5)

 → Cu2O + Na2O 2CuO + 2Na ←  − 264.954

1.37

337

+ 103

359

 → 2Cu + Na2O Cu2O + 2Na ← 

1.18

375

+ 74

336

(6)

− 228.456

a

Theoretical capacities (qth) relate to the active electrode mass, i.e. the copper compounds only. Values for the theoretical energy density (wth) are calculated with respect to the total mass of all elements including Li/Na, Cu and O.

Experimental section Preparation of CuO electrodes, cell assembly and electrochemical measurements: Electrodes were prepared by doctor blading using an aqueous slurry. The electrode composition was 65 wt. % active material (CuO from Sigma Aldrich), 25 wt. % conductive additives (15 wt. % SFG-6 and 10 wt. % Super PLi both from Imerys) and 10 wt. % binder (NaCMC, Sigma Aldrich). Aluminium was used as current collector. After drying at room temperature, electrodes (d = 12 mm) were punched out and further dried at 65 °C (3h under vacuum) followed by calendaring (DPM solutions). The electrode loading was approx. 1.8 mgCuO cm−2. Swageloktype T-cells (Giessen cell 82, 3-electrode configuration) were assembled in an Ar-filled glovebox (GS Glovebox Systemtechnik) with oxygen and water levels below 1 ppm. Sodium or lithium metals were used as counter and reference electrodes. Separators from Whatman (GF/A) and Celgard (polypropylene) were soaked with 80 µL of electrolyte, i.e. 0.5 M NaFSI (sodium-bis(fluorosulfonyl)imide from Solvionic) in EC/DMC (3:7, v/v) (ethylene carbonate, dimethyl carbonate, Sigma Aldrich) and 1 M LiPF6 in EC/DMC (LP30 from Merck, < 10 ppm of H2O). Water impurities were minimized by drying the solvents over molecular sieves for several weeks and the sodium conducting salt for 24 h under vacuum at 120 °C. NaFSI was chosen because it provided better stability and reproducibility of the experiments compared to other salts such as NaPF6 5. Galvanostatic discharge/charge tests were performed in ACS Paragon Plus Environment

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the voltage range of 0.01 – 3.0 V vs. Na/Na+ or Li/Li+ in a Maccor battery cycler (Series 4000) at different currents, namely 67.4 mA g−1, 134.8 mA g−1, 674 mA g−1 and 1348 mA g−1 (i.e. 0.1 C, 0.2 C, 1 C and 2 C). Impedance measurements (EIS) were performed with a VMP3 equipment (Bio-Logic) after every discharge (0.01 V vs. A/A+, A = Na, Li) and charge (3.0 V vs. A/A+) over 50 cycles in the frequency range 10−1 to 106 Hz under AC stimulus with 10 mV amplitude. The CuO/A half-cells (3-electrode configuration) were galvanostatically cycled with 75 mA g−1. CuO thin film model electrodes were used to avoid side reactions related to binders and/or conducting agents. 5. All electrochemical experiments were carried out at 25 °C or 50 °C, respectively. Morphological characterization: Cells were disassembled under Ar atmosphere, and electrodes were rinsed with an EC/DMC (3:7, v/v) solvent mixture. Subsequently, samples were dried in vacuum for 3 h at 65 °C. Electrodes were cut and transferred via an air tight transfer system to the scanning electron microscope (SEM) (MERLIN, Zeiss) for cross-section analysis. Pressure sensor experiments: CuO/Na half-cells (2-electrode configuration) were assembled as mentioned above and galvanostatically cycled at 75 mA g−1 between 0.01 V – 3.0 V at 25 °C on a Maccor battery cycler. Simultaneously a pressure sensor (a PAA-33X absolute pressure sensor, Omega Engineering) was used to record the pressure change of the cell. To guarantee a stable pressure background, the cell was kept under open circuit conditions for 24 h. Differential electrochemical mass spectrometry (DEMS)-experiments: A setup developed at BELLA laboratory (KIT) was used 83,84. A CuO/Na half-cell (2-electrode configuration) was assembled containing 100 µL electrolyte. Quick connects (Swagelok) were used to enclose the system. Gas evolution was analysed by a mass spectrometer (GSD 320, OmniStar Gasanalysesystem, Pfeiffer Vacuum) using a continuous He flow (2 cm³ min−1). To avoid the drying-out of the cell, the He gas was saturated with dimethyl carbonate (DMC). To ensure a stable background, the cell was enclosed into the DEMS system and purged with He for 6 h. Then, the cell was galvanostatically cycled at 75 mA g−1 between 0.01 V – 3.0 V at 25 °C using a VMP3-system.

Results and discussion Electrochemical measurements This section addresses the electrochemical behaviour of CuO electrodes with respect to specific capacity, rate capability, cycle life, coulombic efficiency and average voltage. Measurements were conducted at 25 °C and 50 °C. Results for CuO/Li and CuO/Na are compared and discussed with respect to volume expansions of the conversion reaction systems.

Specific capacity and average voltage at different C-rates and temperatures The theoretical capacity of CuO is 674 mAh g−1, while for Cu2O it is only 337 mAh g−1 assuming a conversion to Cu and A2O.



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Figure 1. Rate-dependent discharge capacity of CuO electrodes in sodium (a) and lithium (b) cells at different C-rates at constant temperature of 25 °C.

Rate capability measurements of CuO/Na and CuO/Li-cells (1 C corresponds to a current density of 674 mA g−1(active mass)) at 25 °C are shown in Figure 1. As expected, the capacity decreases with increasing C-rate for both systems, due to the well-known rate-capacity effect 85. Moreover, the increasing current density results in an increasing IR drop and therefore lower redox potentials (see the differential capacity plot in Figure S1). The most obvious difference between the CuO/Na and CuO/Li system is that the rate-capacity effect is much more pronounced for the latter, see Figure 1a and 1b and differential capacity plot in Figure S2. For CuO/Li capacities up to around 800 mAh g−1 are reached at low currents (0.1 C) suggesting that largely reaction 1 takes place, i.e. CuO is the major charge product. Obviously, the theoretical capacity is exceeded but this has been reported for many conversions reaction. The extra capacity is due to charge storage in the conductive additives as well as related to SEI formation and, in some cases, interface charge storage 86,87. From our experience, the contribution of the conductive additives is typically around 20 mAh g−1 to the measured electrode capacity 4,5. At higher C-rates the capacity values rapidly drop, for example to roughly 300 mAh g−1 at 2 C. Consequently, Cu2O is the main charge product and largely reaction 3 takes place (CuO is a by-product at 2 C) 40,63. In contrast, the CuO/Na system is kinetically limited already at low C-rates reaching up to 400 mAh g−1 at 0.1 C. That means that elemental copper is largely reoxidized to Cu2O only during charging, i.e. reaction 6 instead of reaction 4 takes place. The loss upon increasing the C-rate is, however, smaller as compared to the CuO/Li system. We note that graphite is inactive towards Na-ion storage in the used electrolyte so does not contribute any extra capacity 88. Overall, the results show that the capacity of the CuO/Li system exceeds the one of CuO/Na only at low C-rates. At high C-rates (> 1 C), similar capacity values are found for both systems indicating that Cu2O rather than CuO is the charging product, i.e. mostly eq. 3 and 6 take place. The CuO/Na and CuO/Li systems also differ with respect to the coulombic efficiency ηC (CE). ηC is an important parameter describing the reversibility of the reaction and is defined as the ratio of the total charge capacity and the total discharge capacity. A well-known limitation of conversion electrodes is the low efficiency of the first discharge/charge cycle (also named initial coulombic efficiency, ICE), which is mostly lower than 75 %. The first cycle can be considered as an activation cycle in which side reactions with the electrolyte (SEI for-


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mation) take place and the nanostructure forms. Upon subsequent cycling, conversion systems become stable and values for ηC usually approach 100 % 1.

Figure 2. The coulombic efficiency of CuO/A, A = Na (orange), Li (grey) a) after the 1st cycle (ICE) and b) average over 50 cycles. (Note: Data points are connected for better clarity only)

Figure 2 shows the initial coulombic efficiency (ICE) and the average efficiency (1st to 50th cycle) of the CuO/Na and CuO/Li systems in comparison. At low rate, ICE values are comparable for both systems. With increasing C-rate, ICE values decrease for both conversion systems but the effect is much stronger for CuO/Li. Values for CuO/Na drop from 64 % (0.1 C) to 51 % (2 C) whereas values for CuO/Li drop from 62 % to only 41 %. The rapid drop in CuO/Li can be understood from the aforementioned kinetic limitations to fully reoxidize Cu to CuO at increasing C-rate. Instead, Cu2O is formed (reaction 3) during charging. For CuO/Na, the average efficiency stabilizes at 98.2 % at low rates and at 98.7 % for high Crates. The increase in average efficiency with increasing rate could be due to the fact that a higher current also decreases the time for side reactions. No clear trend is observed for CuO/Li but average CE values are slightly lower compared to CuO/Na. This finding will be further discussed in section 3.1.2. The position of the observed redox potentials is particularly important for the discussion of the electrode reaction. Although conversion reactions theoretically should show well defined constant redox-potentials (see Table 1), voltage profiles appear very complex and are often difficult to interpret in reality. Redox activity is found in a broad voltage window rather than at defined potentials (see Figure S3), which means that the standard methods describing electrode kinetics/thermodynamics fail. To compare the charging and discharging voltages under different conditions, we therefore plot the average voltage, i.e. the number average over all values obtained during galvanostatic charging or discharging. This is a useful simplification because equilibrium potentials for conversion reactions are experimentally hardly accessible. We restrict the analysis to the CuO/Na system as it can be described by reaction 6 for which the equilibrium voltage would be 1.18 V. Figure 3 shows the average discharge and charge voltage at different rates. As can be seen, the combined overvoltages amount to about 1 V at the beginning and reduce to around 0.75 V upon cycling. These values are in the same range as other oxide conversion reactions with lithium 89 Not surprisingly, the average discharge voltage decreases with increasing C-rate. More unexpectedly, the average charging voltage is



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independent of the C-rate and stable values of around 1.5 V are found for at least 50 cycles. This leads us to the conclusion that the charging process is generally faster as compared to the discharge process. For completeness, the analogue results for CuO/Li are shown in the supporting information (see Figure S5). The interpretation is, however, less straightforward due to (a) the mixed behavior of reaction 1 and 3 and due to (b) the graphite additive that is only redox active in case of lithium (see Figure S2) and lowers the average voltage. The combined overpotentials are, however, slightly larger as compared to the CuO/Na system.

Figure 3. The average voltages of CuO electrodes in sodium cells at different C-rates at constant temperature of 25 °C.

From the measurements above it is clear that CuO can reform in case of lithium cells (at low rates) but not in case of sodium cells. The reason for this is not quite clear but might be related to the volume expansion of the electrodes during cycling. Conversion electrodes are generally plagued by large volume changes due to the formation of different phases with different density. As the size of the sodium ion is larger compared to the lithium ion (r(Na+) = 1.02 Å vs. r(Li+) = 0.76 Å) 90, volume expansions are larger for sodium-based conversion reactions. For oxides, volume expansions during lithiation or sodiation are typically in the range of 100 % (Li) and 200 % (Na), respectively 4. During the first discharge, the volume expansion of the ideal conversion between CuO and Na to Cu and Na2O reaches 174 % (reaction 4), which is significantly higher than in the case of Li (74 %). However, the effective difference between both systems during cycling is much smaller as in the case of CuO/Na largely reaction 6 instead of reaction 4 (see Table 1) takes place. For reaction 6, the volume change is only 74 %, i.e. the same value as for reaction 1 in the case of lithium. A graphical illustration of the volume changes is shown in Figure S4. Both systems therefore undergo comparable volume changes during cycling which might well explain the finding that the CuO/Na is only partially reversible even at low rates. This behaviour can be linked to the empirical chemical rules from Ostwald and Ostwald-Volmer 91,92. These qualitative rules state that a system will not reach the stable ground state directly but instead passes through less stable states and that on the way to dense compounds, less dense compounds are formed as intermediates. In our case this means that Cu2O forms as reaction ACS Paragon Plus Environment

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intermediate, i.e. during charging Cu2O forms prior to CuO. After Cu2O has formed, the large volume change in case of CuO/Na then hinders further oxidation to form CuO which result in lower capacities compared to CuO/Li systems. Moreover, it has been observed that many sodium-based conversion systems exhibit lower capacities than their analogous lithium systems which might be a result of larger volume changes. Nevertheless, the volume changes in CuO/Na systems might also positively affect the cycling stability. In the first cycle (formation of Cu from CuO and reformation of Cu2O), the expansion is much larger than the shrinkage, eventually creating voids that buffer the volume changes in the forthcoming cycles. Finally, as kinetic limitations might be overcome by increasing temperature, we studied the CuO/Na cells at 50 °C. Note: The shift in equilibrium redox potentials due to the temperature increase is negligible (see Table S1). In case of CuO/Na, the first charge capacities increase from 309 mAh g−1 (25 °C) to 483 mAh g−1 at 50 °C, respectively and therefore initial coulombic efficiencies increase from 59 % to 72.4 %. At 50 °C the capacity exceeds the one of reaction 6 (337 mAh g−1) and therefore reaction 4 partially takes place as well. Indeed, the oxidation of Cu2O to CuO at 50 °C can be seen from the differential capacity plots shown in Figure 4. Unfortunately, this situation is not permanent and the signal due to Cu2O oxidation vanishes in the subsequent cycles. Capacity values drop to around 300 mAh g−1 within 30 cycles (see Figure 5a) reaching values similar to the ones obtained at 25 °C. It is quite likely that this rapid capacity loss is also due to enhanced electrolyte decomposition at 50 °C. In any case, the temperature increase reduces the overpotentials, see Figure 5b. This is especially obvious for the discharging process which has been already identified to be slower as compared to the charging process (see Figure 3). By increasing the temperature, the combined overpotentials reduce to around 500 mV to 600 mV.

Figure 4. Differential capacity of CuO electrodes in the sodium system for different cycles at 0.1 C and 50 °C.



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Figure 5. Results for CuO/Na at 25 °C (blue) and 50 °C (red) a) discharge (closed circles) and charge capacity values (open circles) as well as coulombic efficiency and b) average voltages.

Interphase formation monitored by EIS measurements Formation and stability of the SEI have been studied by electrochemical impedance spectroscopy (EIS). It is worth noting that the quantitative analysis of impedance spectra of thick and porous electrodes is very challenging due to a multitude of different effects related to the nature of real electrodes (electrode porosity, additives, surface films, etc.) and the cell geometry. However, EIS is a very useful tool to identify differences between the different cells (Na vs. Li) as well as to study ageing effects. We used dense CuO thin film electrodes free of any binder and conductive additive and focus on qualitative differences between the CuO/A systems. As will be seen, this is useful to describe interphase formation as a function of time. The Nyquist plots with fits and SEM images of the discharged CuO electrode for the Na- and Li-system are shown in Figure 6. A more detailed analysis of the impedance data and the fitting model are shown in Figure S6.

Figure 6. a) Nyquist-plot and fits of CuO/A (A = Na (orange), Li (grey)) half-cells after the 1st discharge to 0.01 V vs. A/A+ and b) cross-sectional SEM images of CuO electrodes after discharge and charge in the first cycle in CuO/A systems.


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The resistances are much larger for CuO/Na systems as compared to CuO/Li (see Figure 6a and Figure S6), suggesting a thinner SEI for the latter. This is well supported by the cross-section images shown in Figure 6b, where the surface film of the CuO/Na electrode is much thicker (6 µm) than that of the CuO/Li electrode (1 µm) 5 . The comparison also reveals noticeable differences in the semicircles of the Nyquist plots. The semicircle asymmetry observed in the CuO/Li system could be due to the multi-layered structure of the SEI, which is consistent with Aurbach’s impedance model 93. In contrast, the non-layered structure of the SEI in the CuO/Na system, with a more homogeneous (and inorganic) composition results in a more symmetric semicircle. Both cases are also in line with our previous XPS study on the interphase formation 5.

Figure 7. Nyquist-plot of a) CuO/Na and b) CuO/Li half-cells after 1st, 2nd, 10th and 50th discharge. For a clearer overview only the fit of the 1st cycle is shown.

Figure 7 shows impedance spectra recorded after different number of cycles. For CuO/Na the impedance decreases with time and stabilizes after several cycles (see Figure S6). This means that the SEI formation is very dynamic and it takes several formation cycles before stabilization. For CuO/Li, the SEI appears to change not much. Both is in good agreement with the SEM images in Figure 6b. The lack of variations in the semicircle shapes for CuO/Na indicates that the observed changes are mostly caused by the dynamic formation/dissolution of the SEI rather than by variations in the chemical composition 5. Thus, the homogeneous inorganic SEI of CuO/Na systems and the multi-layered heterogeneous SEI of CuO/Li systems are both maintained.

Pressure monitoring during galvanostatic cycling (PMGC) The formation of the solid electrolyte interphase comes along with the release of gaseous compounds that cause a change in cell pressure. Studies on the SEI formation on graphite or other electrode materials showed that the major gaseous compounds released are H2, CO2 and C2H4, for example 46,48,51,52,83,84,94–104. Some gas evolution during initial lithiation/sodiation is not an issue but continuous gas release during cycling is a clear sign for an unstable electrode/electrolyte interface and hence cell degradation. In this section, we summarize our results on pressure monitoring of CuO/Na cells that are cycled galvanostatically. DEMS measurements will be discussed in the following section.



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Figure 8. Voltage (black/grey) and pressure change (blue) during galvanostatic discharge/charge for CuO/Na half-cells during a) 1st cycle and b) 2nd cycle.

In order to achieve a stable starting signal, the assembled CuO/Na cells were equilibrated for 24 h prior to cycling. CuO/Na half-cells were then galvanostatically cycled at 75 mA g−1CuO between 3 V and 0.01 V at 25 °C. During cycling the change in cell pressure was monitored in-situ. The voltage profiles and the corresponding cell pressure of the first and second cycle are shown in Figure 8. The cell pressure changes non-linear and can be divided into three phases: i) a slight increase between 1.3 V and 0.8 V, ii) a steep rise between 0.8 V to 0.1 V and iii) again a moderate increase between 0.1 V and the cut-off potential of 0.01 V. In total, the cell pressure increases by 10.6 mbar indicating that major degradation processes take place during the first sodiation step. Assuming that only gaseous degradation products are generated this corresponds to approximately 2 µmol gas using the ideal gas law. In the subsequent charge a pressure decrease of ≈ 2 mbar occurs. This pressure decrease is significantly lower than the pressure increase during previous discharge. During the second discharge, the pressure decreased continuously at first by about 1 mbar until a potential of ≈ 0.7 V is reached. Then, the pressure starts increasing again. This increase correlates well with the onset of phase ii) in the first cycle, which was the phase with the strongest pressure increase. Afterwards, from 0.7 V to 0.1 V the cell pressure slightly rises again (≈ 0.9 mbar) and subsequently decreases until the end of the 2nd charge. The pressure change over the cell amounts then to 2.9 mbar. The decrease in cell pressure is surprising at first but might be due to several effects that occur at the same time: i) gas molecules formed during discharge (or during open circuit conditions) are consumed during charge forming secondary solid products such as Na2CO3, ii) the gas molecules dissolve partially in the electrolyte 102,103 and/or iii) the total volume of the electrodes or the electrolyte decrease as a result of the electrolyte decomposition and partial reversibility of the reaction. The formation of sodium carbonate has been previously demonstrated 5 and the dissolution of some potentially formed gases has been shown for Li-ion batteries (LIBs) 102,103. The last effect (iii), however, is negligible as the volume changes due to the conversion reactions result in pressure changes smaller than 0.25 mbar only. At the end of the 2nd cycle, the cell was rested for 10 hours before the experiment was continued. A notable increase in cell pressure of about 3 mbar occurred in the 3rd cycle, see Figure 9. Obviously, resting the cell leads to a more intense pressure change during the following cycle.


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Figure 9. Voltage (black/grey) and pressure change (blue) during galvanostatic discharge/charge for CuO/Na half-cells during continuous cycling.

The deposition of surface films on the electrodes when extending the rest time has been previously reported for other battery systems 105, and suggests that also a fraction of the increase observed during the first cycle could be ascribed to this process. In the following cycles, the pressure changes become smaller which is consistent with the results presented in the impedance section where the stabilization of the SEI after several cycles has been demonstrated. The recorded pressure signals do not allow a more detailed interpretations of the underlying side reactions. The main advantage of this method is, however, its simplicity and that it can be applied in closed cells. Any pressure change during cycling is indicative for side reactions that are otherwise easily overlooked. Although pressure sensor experiments are a suitable tool to monitor the evolution of gaseous products during cycling, no direct evidence of the chemical nature of the formed gases can be attained. Thus, in order to complete the gas analysis DEMS measurements were carried out.

DEMS measurements during cycling A powerful tool for the in-situ analysis of gas evolution reactions is differential electrochemical mass spectrometry (DEMS). Evolving gases can be detected and a relationship between their formation and the electrode potential voltage can be established. However, cells eventually can dry out easily because solvent vapour is removed from the cell as well, so often only few cycles are reported. We therefore saturated the gas atmosphere with DMC solvent, enabling DEMS measurements for 10 cycles. CuO/Na cells were assembled and rested for several hours under open circuit voltage (OCV) conditions before cycling them galvanostatically at 75 mA g−1CuO. At the same time, DEMS measurements were running to detect the gaseous compounds. The voltage profiles of CuO/Na cells and the DEMS signals of the first cycle are shown in Figure 10. The OCV is coloured in yellow, the discharge in green and the charge in red.



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Figure 10. i) Voltage profiles and DEMS signals for ii) m/z = 2 and iii) m/z = 44 of a CuO/Na half-cell at 75 mA g−1 and 25 °C of the 1st cycle. The open circuit voltage (OCV), discharge and charge regions are coloured in yellow, green and red, respectively.

The voltage profiles in panel (i) show the expected complex behaviour with redox activity in a broad voltage range. In our previous study on thin film electrodes we found that reduction of CuO to Cu2O occurs at approximately at 1 V vs. Na/Na+ and to metallic Cu at about 0.48 V during discharge. Then, the oxidation to Cu2O takes place during charge at about 2.0 V. Moreover, the electrolyte decomposes below 0.4 V vs. Na/Na+ 5. Panel (ii) and (iii) show the intensity signals detected by the mass spectrometer. During the open circuit period, the ion currents for both m/z = 2 and 44 increase. These signals correspond most likely to hydrogen (H2, 2) and carbon dioxide (CO2, 44), which show different patterns. As no current flows under OCV conditions, the evolution of H2 and CO2 are caused by chemical reactions due to side reactions with the metallic sodium electrode. Considering the high reactivity of the Na anode and the presence of traces of residual water in the electrolyte and/or electrodes, a fast reaction occurs during the cell assembly resulting in the H2 formation 106. During the following discharge, both H2 and CO2 evolution is observed. The H2 evolution rate increases between OCV and 1.3 V vs. Na/Na+, then decreases between 1.3 V – 0.8 V and further increases between 0.8 V – 0.48 V and again decreases in the range of 0.48 V – 0.01 V. Thus, the degree of H2 formation is linked to the electrode potential. Moreover, a correlation between H2 and CO2 can be observed as CO2 evolution proceeds contrary to the H2 one, i.e. if the formation of H2 increases, the one of CO2 decreases. Quantification of the signals is shown in Figure S7 and S8. Despite the uncertainty at low ppm-levels, it can be stated that the H2 concentrations are significantly higher than those of CO2. During the subsequent charging, the formation rate of H2 decreases between 0.01 V – 1.0 V vs. Na/Na+. Then it increases from 1.0 V – 2.25 V and remains almost constant up to 3.0 V. During charging, the CO2 signal continuously increases. Formation mechanisms of H2 and CO2 have been suggested in literature and can be summarized as follows 46,48,51,52,83,84,94–102: H2 can be produced by the electrochemical reduction of trace water in the electrolyte on the sodium electrode (reaction 7). Then, a nucleophilic attack ACS Paragon Plus Environment

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by the hydroxide (OH−) anions could open the ring structure of the ethylene carbonate (EC) solvent (reaction 8) and generate CO2 gas. (7)

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Furthermore, the alkyl di-alcoholate (−O-CH2-CH2-OH) can repeatedly decompose EC molecules by nucleophilic attack (reaction 9), leading to CO2 formation without consuming significant fractions of the trace water and/or without releasing H2. (9)

The generated CO2 can be consumed according to reaction 10. On the one hand, OH−-ions could react with CO2 to form carbonates (e.g. Na2CO3 and HCO3−) on the other hand oligomerization reactions of di-alcoholates could proceed. (10)

CO2 might be also released due to decomposition of Na2CO3 during charging. Nevertheless, during the 1st discharge firstly H2 is generated and then CO2, thus reaction 7 and 8 are favoured above 0.8 V vs. Na/Na+. The evolution of H2 mainly proceed because of the reduction of residual water traces in the electrolyte and/or the electrodes. Below 0.8 V, H2 is formed and the reaction rate of CO2 decreases due to the formation of carbonates (reaction 10). Moreover, carbonates are the main component of the SEI in CuO/Na systems 5. Then below 0.48 V reaction 8 is again favoured. During the following charge, both CO2 and H2 are formed. During charge, the SEI or more specific the carbonates dissolve and thus could result in the generation of CO2. Due to electrolyte decomposition protons, H+ and/or R-H+ (organic protons with R = organic group) could be formed and reduced to H2 96–98.



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Figure 11. i) Voltage profiles and DEMS signals for ii) m/z = 2 and iii) m/z = 44 of a CuO/Na half-cell at 75 mA g−1 and 25 °C of several cycles. The green shading represents the discharge cycles, while the OCV is coloured in yellow.

DEMS signals for a series of consecutive cycles are shown in Figure 11. It is seen that the first cycle is clearly different from the following ones. The gas signals during the first cycle are stronger and exhibit distinct maxima at different potentials, which is clearly related to the initial SEI formation, the formation of the nanoscale structure of the electrode as well as breaking the SEI on the sodium counter electrode. Starting from the second cycle, the voltage behaviour is quite reversible and accompanied by periodic changes in the H2 and CO2 evolution. H2 formation decreases during cycling because the residual trace water is gradually removed over time. For CO2 it is seen that its formation rate increases during charging while its reaction rate decreases during discharge. The signal remains on similar intensity levels, indicating that CO2 formation is a continuous, although periodic, process. This periodicity might also explain the dynamic formation/decomposition of the SEI (see SEM images in Figure 6b). At the same time, the renewed SEI exhibits the same properties resulting in constant impedance values after a few cycles (see Figure 7a). We emphasize that we never visually observed that cycled CuO/Na cells have dried-out after running for as much as 100 cycles. The degradation is therefore a comparably slow process. We remind that these conclusions are of course only true if the periodic gas release is not due to the sodium metal counter electrode. Although the contribution of the metal electrode has been proven to be small 51, separating the contributions of both electrodes is therefore a worthwhile goal for future studies. Based on the present findings and the result from our previous study on CuO thin films a complete picture of the processes in a CuO/Na cell evolves that is illustrated in Figure 12.


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Figure 12. Schematic model for the reaction mechanism, the SEI formation and gas evolution of CuO/Na systems during a) open circuit conditions and after b) discharge (sodiation) and c) charge (desodiation).

Under open circuit conditions (Figure 12a), traces of water are chemically reduced on the Na counter electrode to form H2. Furthermore, a thin surface layer on Na, is generated. During the first discharge/1st sodiation (Figure 12b), CuO is reduced to metallic Cu (via Cu2O) and Na is oxidized to Na2O. A thick surface film which is mainly composed of inorganics like carbonates and NaF forms below 0.4 V vs. Na/Na+. Because of the reduction of residual water traces in the electrolyte and/or the electrode, mainly during the first discharge, hydrogen is formed. The decomposition of electrolyte solvents, namely EC, usually results in the generation of CO2 96–98. Moreover, CO2 could lead to the formation of solid carbonates, which have been often reported when analysing the electrode surfaces. During the first charge/1st desodiation (Figure 12c), Cu is oxidized to Cu2O and a small amount of Na2O2 is generated. The SEI, mainly composed by carbonates, is partially dissolved at this stage, being consistent with the formation of CO2. Additionally, (organic) protons (H+ and R-H+) and/or residual water are reduced to generate H2. From the second cycle on, a reversible conversion process between Cu2O and Cu occurs as well as the continuous formation of H2 and CO2. Although the intensity of the side reactions decreases by the time, they do not completely disappear.

Conclusion The use of CuO as electrode material in lithium and sodium cells was compared and electrolyte degradation has been evaluated by means of in-situ pressure measurements and DEMS analyses. At low currents, the capacity of the CuO/Li system largely exceeds that of the CuO/Na system. This is simply because CuO (or more precisely a CuO/Cu2O mixture) forms during charging for the former whereas only the intermediate Cu2O phase forms for the latter. ACS Paragon Plus Environment


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The degree of CuO formation is highly sensitive to the current density and the capacity rapidly drops when increasing the current. The higher the current, the larger the fraction of the Cu2O intermediate. The capacity of the CuO/Na with Cu2O as charging product depends much less on current density providing 280 mAh g−1 (2 C)–400 mAh g−1 (0.1 C). The oxidation of Cu to Cu2O during charging was found to be a comparably fast process. The lack of full oxidation to CuO in the sodium cell even at low current densities might be linked to the volume changes during cycling. Indeed these changes are the same for reaction 1 (formation of CuO as charging product in lithium cells) and for reaction 6 (formation of Cu2O as charging product in sodium cells). The sluggish kinetics of full oxidation to CuO during charging in sodium cells can be overcome by increasing temperature although so far at the expense of cycle life. Ageing processes related to electrolyte degradation and surface film formation have been observed for both systems. Impedance measurements show that surface films form and stabilize for CuO/Na and CuO/Li systems, but it takes longer in case of CuO/Na. Pressure monitoring and DEMS analysis provided further insight into the ageing mechanisms of CuO/Na cells. The use of PMGC has been shown to be a simple yet straightforward tool to proof gas release during cycling. DEMS analysis provides more information but the experimental setup is more complicated and drying out of the electrolyte has to be minimized in order to analyse several cycles. PMGC and DEMS measurements showed periodic changes in cell pressure and gas release (H2, CO2) during cycling. Although the intensity of the recorded signals slowly fades during cycling it is likely that the SEI is periodically formed and dissolved during cycling. This essentially means that the electrolyte is consumed in side reactions during cycling. We note that we never observed any drying-out of our cells, so the degradation process is not very fast. Nevertheless it is likely that very long cycle life can be only achieved by using excess electrolyte unless other solutions are found.

Supporting information The Supporting Information include differential capacity plots of CuO/A (A = Na, Li) at different C-rates (Figure S1 and S2), voltage profiles of CuO/Na (Figure S3), volume expansions of CuO/A (Figure S4), average voltages of CuO/Li (Figure S5), the resistance of the SEI in CuO/A (Figure S6), voltage profiles and gas concentrations in CuO/Na (Figure S7 and S8) and thermodynamic data of several reactions for CuO/A at 50 °C (Table S1).

Author information [email protected]

Acknowledgements Financial support is gratefully acknowledged from the German Research Foundation (DFGProject Thermodynamik und Kinetik von Konversionsreaktionen in neuen, natriumbasierten Batteriesystemen), from the state of Hesse (Landes-Offensive zur Entwicklung Wissenschaftlich-ökonomischer Exzellenz, LOEWE) within the project Store-E (Stoffspeicherung in Grenzflächen) and from the state of Thuringia (ProExzellenz program). We thank Imerys Graphite & Carbon for providing carbon materials and we thank P. Hering and A. Polity from


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the 1st Physics Institute (Justus-Liebig University Gießen, Germany) for providing CuO thin film model electrodes.

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Kinetics and Degradation Processes of CuO as Conversion Electrode

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