1. Introduction
The rapid growth of modern society’s need for reliable and efficient energy storage systems has driven the development and widespread use of lithium-ion battery (LIB) technology. However, the cost and limited availability of lithium sources restricts the use of these devices in a number of applications. Sodium-ion batteries (SIB) are a promising solution to this problem, showing strong potential for use in electric transport and stationary systems of buffer electricity accumulation [1,2,3,4]. This scope of SIB application necessitates their stable operation across a wide range of external temperature conditions. However, very little is known about the low-temperature properties of materials and electrolytes used in SIBs.
The temperature range from +15 °C to +35 °C is considered optimal for LIBs [5]. The deterioration of their electrochemical and mechanical properties during operation—especially charging—at low temperatures has been shown in many works [6,7,8,9,10,11,12]. In a recent study by Laforgue et al. [13], post-mortem analysis through X-ray computed tomography, scanning electron microscopy, and other techniques revealed several failure modes observed in commercial LIBs when charged at reduced temperatures. These failure modes include lithium plating, graphite exfoliation, jelly roll deformation, active materials crumbling, aluminum corrosion, and abnormal growth of the solid electrolyte interphase (SEI) on the anode side. From the point of view of the electrochemical properties of the cells, the following effects are observed with decreasing temperature: phase separation as a result of solvent freezing or salt precipitation of the electrolytes, decrease in Li+ diffusion coefficients in active materials, increase in resistance of the formed solid electrolyte interface at the electrode and electrolyte, and deceleration of the charge transfer kinetics at the electrode and electrolyte interface [14].
It is well-known that the low-temperature properties of LIBs significantly depend on the electrolyte system composition [15]. It affects the electrolyte ionic conductivity, as well as the SEI properties and charge transfer rate at the liquid–solid interface [7,16]. However, lowering the temperature can affect the composition of the electrolyte itself due to the partial crystallization of certain constituents (both solvent and salt), which causes drastic changes in the composition and physicochemical properties of the system that are challenging to describe. For example, Che et al. [17] studied the low-temperature properties of several NaPF6 solutions and observed that the conductivity of some of them dropped to almost zero after the crystallization of the solid phase.
It is worth noting that the temperature range in which the electrolyte solution remains functional can be wide despite the phase separation. For example, the solution of 1.0 M LiPF6 in EC:EMC = 3:7 is shown to start releasing the crystalline phase at −20.4 °C (the temperature at which the first solid phase precipitates in cooling from a homogeneous liquid electrolyte) [18]. Yet, it remains operational at −40 °C; graphite||NCA full cells with this electrolyte solution demonstrate a capacity of about 30 mAh/g at a C/5 rate [18].
Among nonaqueous electrolyte systems for SIBs, the most common solvents are cyclic and linear carbonates, esters, and ethers: ethylene carbonate (EC), propylene carbonate (PC), dimethyl carbonate (DMC), diethyl carbonate (DEC), 1,2-dimethoxyethane (DME), diethylene glycol dimethyl ether (DEGDME), etc. [19,20]. EC is a polar solvent with the highest dielectric constant among those mentioned. As a consequence, it has the best ability to dissolve various salts. Shakourian-Fard et al. [21] demonstrated that the solvation of the sodium ion in EC results in the greatest energy gain among single solvents (the Gibbs free energy of solvation ΔG(sol) = −71.63 kcal/mol (−299.70 kJ/mol): Na+(EC) > Na+(PC) > … > Na+(DEC) > Na+(DMC). However, EC has a high melting point and a high viscosity (Table 1). Therefore, due to their low dielectric constants, linear carbonates with low viscosity and relatively low melting points are used as co-solvents since they cannot perform as single solvents (examples can be seen in Table 1). Solvation energy for Na+ in these mixtures is close to that for EC: −70.72 kcal/mol (−295.89 kJ/mol) for Na+(EC:DMC) and −60.11 kcal/mol (−251.50 kJ/mol) for Na+(EC:DEC) [21]. The conductivity values are also high and comparable with those of a Li-ion system: ≈7.2 mSm/cm for EC:DEC = 2:3 and ≈11.0 mSm/cm for EC:DMC = 2:3 (20 °C, 0.8 M NaPF6 solutions) [17]. Moreover, EC as a co-solvent for SIB electrolytes promotes a more stable SEI, which is likely related to the formation of ether functionalities upon reduction [22,23].
It should be noted that the researchers are not limited to mixtures of EC and linear carbonates in the development of low-temperature electrolytes. For example, 1,2-dimethoxyethane (DME) is a fairly popular co-solvent. Ponrouch et al. [24] used the low-temperature differential scanning calorimetry (DSC) method to study several electrolyte systems, such as 1 M NaClO4 in EC:DEC, EC:DMC, and EC:DME. On the heating curve, the first endothermic peak indicating the first crystallization was observed at −25, −50, and −75 °C for EC:DMC, EC:DEC, and EC:DME mixtures, respectively. This observation supports utilizing DME as one of the components in the solvent system for low-temperature SIBs.
In this work, the low-temperature properties of liquid electrolyte systems with NaPF6 salt dissolved in ethylene carbonate: diethyl carbonate (EC:DEC), ethylene carbonate: dimethyl carbonate (EC:DMC), or ethylene carbonate: dimethoxyethane (EC:DME) mixtures were studied using low-temperature DSC and electrochemical methods.
2. Materials and Methods
2.1. Solution Preparation
Battery grade 99% anhydrous ethylene carbonate (EC, Merck, formerly Sigma Aldrich, Rahway, NJ, USA), 99.5% anhydrous inhibitor-free 1,2-dimethoxyethane (DME, Sigma Aldrich), 99+% extra dry dimethyl carbonate (DMC, Acros Organic, Geel, Belgium), and 99% anhydrous diethyl carbonate (DEC, Acros Organic) were used as solvents. Solvent mixtures of EC:DEC, EC:DME, and EC:DMC (1:1 by volume; 64.8, 61.1, and 56.1 mol.% of ethylene carbonate, respectively) were dried with molecular sieves 4Å granules (Sigma-Aldrich). Sodium hexafluorophosphate (NaPF6, 98%, Sigma Aldrich) was dried under vacuum at 60 °C and dissolved into solvent mixtures in various molar ratios, creating the following compositions of the electrolyte systems: 0 M, 0.5 M, and 1 M NaPF6 in each solvent mixture (EC:DEC, EC:DME, or EC:DMC).
2.2. Characterization
Differential scanning calorimeter DSC 204F1 Phoenix (NETZSCH, Selb, Germany) with dynamic nitrogen atmosphere (40 cm3/min) was used for the DSC measurements. Sensitivity and temperature calibration was performed with the following standards: Hg, C6H12, C10H16, In, Sn, Bi, H2O, Zn, and CsCl. An amount of 3–5 μL of the sample was added to an aluminum pan and hermetically sealed in a glove box. The pans were cooled (0.3 K/min) to –70.0 °C and then heated (0.3 K/min) to +25 °C. Peak temperature values were recorded from the cooling and heating thermograms.
Measurements of physical characteristics and electrochemical studies at low temperatures were carried out using the low-temperature thermostat “LT-912” (LOIP, Saint-Petersburg, Russia). The temperature of freezing was measured in a test tube using the ADA Thermotester 330 (ADA Instruments, Hong Kong, China). Electrochemical measurements were carried out using the P-45x potentiostat-galvanostat (Electrochemical Instruments, Chernogolovka, Russia) in two-electrode cells with steel–steel electrodes (ionic conductivity evaluation) or Na3V2(PO4)3-Na3V2(PO4)3 electrodes (NVP-NVP; galvanostatic experiments and impedance spectroscopy). The Na3V2(PO4)3/C (NVP) electrodes were composed of 93% of active material, synthesized according to the route described in [25], 1% of carbon, and 6% of polyvinylidenefluoride. The dry mixture was blended with N-metyl-2-pyrrolidone, spread on carbon-coated aluminum foil, dried at 90 °C in a heating chamber, calendared, cut into ø15 mm electrodes, then vacuum dried at 110 °C for at least 1 h. The loading of electrodes was 2.8–3.4 mAh/cm2 for the chronopotentiometry experiments and 0.5–0.9 mAh/cm2 for the electrochemical impedance spectroscopy (EIS) experiments. The cells were assembled inside an Ar-filled glove box. The borosilicate separator GF 50 (Schleicher & Schuell MicroScience GmbH, Dassel, Germany) when assembling an electrochemical cell was used. The rate was set to C/100 (1.2 mA/g) for the NVP-NVP cell measurements in the galvanostatic mode of chronopotentiometry studies. EC-Lab BioLogic software was used to fit the impedance spectra.
3. Results and Discussion
The general scheme of the studies carried out in this work is shown in Figure 1a. It is worth noting that standard methods for studying thermodynamic properties imply heating a pre-frozen object of interest. This is motivated by the fact that liquid can transition into a metastable supercooled state during cooling from room temperature, which introduces errors in the values of the determined parameters and jeopardizes reproducibility. However, from a battery operation perspective, the cooling process is of primary interest. Therefore, the following experiment was conducted in the first part of this study. All solutions were placed in hermetic test tubes with a ThermoTester; after that, test tubes were cooled in a cryothermostat at a rate of 0.3 K/min. Table 2 shows the initial crystallization temperatures of the studied electrolyte systems. These points were easily identified both visually and by a small (several degrees) increase in the ThermoTester temperature caused by the exothermic crystallization process. It is worth noting the studied mixtures increasingly resembled melting snow as the temperature decreased. Therefore, the points of initial crystallization are marked as “wet snow”, or WS, hereinafter. As an example, the solution of 1 M NaPF6 in EC:DEC before crystallization and its “wet snow” stage is shown in Figure 1.
Most studies concerning the low-temperature properties of metal-ion batteries tend to overlook the state of matter of electrolytes, although this aspect is extremely important from our perspective. Unfortunately, it seems unfeasible to determine the precise composition of the liquid and solid phases following a phase transition, since phase diagrams for such systems are not described in the literature and their development is associated with high experimental costs. We assume that the main component of the crystal phase formed below the WS point is EC, given that its melting point is the highest among the studied solvents. In addition, since EC exhibits the highest dielectric constant and the ability to dissolve NaPF6, the precipitation of a large amount of EC must inevitably lead to the partial crystallization of the salt. It is also noteworthy that the freezing points of solvent mixtures without salt (“0 M” solutions) correlate with the data from the literature [26] on the phase diagrams of these mixtures, although lower by several degrees due to kinetic factors.
To better understand precipitation processes for these compositions and to be able to vary solvent composition in other ratios, it would be useful to perform a solid–liquid phase equilibria (SLE) calculation using a thermodynamic model. Although there are approaches to fit a thermodynamic model for a mixed solvent [27], this model, close to polynomial ones, is unfit for electrolytes. To model the thermodynamic properties of liquid in such systems using any thermodynamic model, e.g., ePC-SAFT [28], one needs to conduct further research concerning NaPF6-one solvent systems (solvent activity, SLE, mean ionic activity coefficient) and to do a new assessment for a mixed solvent [27], using the model. For a few selected compositions, it is easier to do an empirical study of SLE. Additionally, any computed phase diagram requires verification with experimental data.
The DSC method was used to study phase transitions that occur in electrolytes during cooling and heating. The cooling rate was chosen to be the same as that used for the test tubes. Thermograms are shown in Figure 2.
DSC thermogram data indicate that, even at extremely low cooling rates, the solutions crystallize at temperatures from −60 °C to −40 °C, despite the difference in phase transition temperatures. The cooling graphs do not contain peaks corresponding to the temperatures of the crystallization onset that were observed in test tubes, nor do they mirror the heating graphs for the electrolyte systems. A similar phenomenon has been observed previously [18,29,30]. This suggests that all solutions become metastable supercooled liquids during cooling in the DSC experiment.
A DSC curve represents a polythermal cross-section of a phase diagram if it is performed on heating. Thus, the first effect on the curve is usually regarded as a solidus temperature (the first point where a portion of liquid appears), but sometimes at the conditions of measurement not all sample becomes crystalline and the effect can not be registered [31]. The last effect usually indicates the liquidus—the last point where the crystals exist on heating—and the first point where a solid starts to precipitate, a solvent or a solvate [31]. In a ternary system, there should be more effects, e.g., indicating a temperature where a second solid phase starts melting.
The EC:DMC binary blend has the highest eutectic temperatures from all three solvent mixtures [32], which may explain why there are all three expected effects in the curve. The eutectic temperatures for systems with DEC and DME should be lower than their melting points (Table 1) and, thus, we can conclude that we did not manage to obtain a completely solid sample before heating.
Thus, on the one hand, several DSC peaks were obtained upon heating for each sample, which indicates the melting of the electrolyte components. Some of these peaks correspond well to the “wet snow” points that we observed in the experiment described above: −7.1 °C for 1 M NaPF6 in EC:DEC = 1:1 (Tws = −7.8 °C); −11.1 °C for 1 M NaPF6 in EC:DME = 1:1 (Tws = −13.1 °C); and −19.5 °C for 1 M NaPF6 in EC:DMC = 1:1 (Tws = −21.2 °C).
On the other hand, in the case of 1 M NaPF6 in EC:DME = 1:1 and 1 M NaPF6 in EC:DMC = 1:1 there are effects with peak temperatures 15.9 °C and 3.9 °C. It is highly likely that it is caused by EC melting/precipitating, but the amount of precipitate is not high enough to see in the tubes or to fill the cell to prevent the liquid part to flow. For EC:DMC, the point is close to liquidus temperatures in binary EC:DMC mixtures [32]. The second peaks, which are close to “wet snow”, are likely to be the points where EC starts to co-precipitate with another solid phase.
Thus, we believe that the behavior of the electrolytes during their cooling can strongly depend on the experimental procedure. To understand which of the observed patterns best correspond to the electrochemical properties of the solutions under study, we carried out a series of experiments in sodium-ion cells during cooling. First, the specific resistance of solutions in two-electrode cells with steel electrodes was determined by electrochemical impedance spectroscopy [33]. Figure 3 illustrates the temperature dependence of the specific electrical conductivity.
In the case of DME (red graph), two inflection points are visible regardless of concentration. The other two systems (black and blue graphs) have only one inflection point. Further, 1 M NaPF6 in EC:DME exhibits the highest conductivity during initial cooling, which may result from a higher permittivity and lower viscosity of DME in comparison to DMC and DEC (see Table 1). However, 1 M NaPF6 in EC:DMC performs better at −20 °C < t < 0 °C, while 1 M NaPF6 in EC:DEC maintains higher conductivity below this temperature region.
To understand how changes in the composition and conductivity of electrolytes during their cooling affect Na+ (de)intercalation into electrode materials, the following test was carried out in symmetrical two-electrode cells with the NASICON-type Na3V2(PO4)3 (NVP) material. It is well-known that Na3V2(PO4)3 can be both oxidized and reduced to NaV2(PO4)3 and Na4V2(PO4)3, respectively. Both processes are characterized by a flat voltage plateau. Therefore, at low current densities, symmetric NVP-NVP cells show a linear E-x dependence at a voltage of ≈1.75 V. The cells were placed in a cryothermostat and cooled at a rate of 0.3 K/min, with a constant current of C/100 applied to the electrodes. The temperature dependencies of the overvoltage are shown in Figure 4a–c. The registered ΔE was transformed into resistance as R = (E − 1.75)/I; its reciprocal value, which can be considered “effective conductivity” k, is shown in Figure 4d,f.
Based on the data obtained, the following features of the systems under study can be noted. First, all electrolytes demonstrate a change in the course of the temperature dependence at certain points, which is expressed in the breakpoints of the k exponential graphs in Figure 4d,f. For EC:DEC and EC:DME, these points are several degrees lower than the temperature of “wet snow” or DSC peaks, but in the case of EC:DMC, these values match closely. Second, the change in the temperature dependence slope of the parameter k in Arrhenius coordinates strongly depends on the co-solvent. For DEC, this change is minimal, whereas the course of both curves transforms drastically in the case of DME and DMC. As a result, the overvoltage in the case of EC:DEC stays below 0.2 V when −40 °C is reached, but for the other two systems it exceeds 1 V.
To analyze the overvoltage that occurs when cells are cooled in more detail, electrochemical impedance spectroscopy was also performed in the NVP-NVP cells, as illustrated by Figure 5a–d. Figure 5e depicts the equivalent circuit that was used to process the data. The resulting values of the ohmic (Rs) and charge transfer (Rct) resistance are presented in Table 3.
The temperature dependence of the inverse values of the solution resistance (Rs) and the charge transfer resistance (Rct) in Arrhenius coordinates are shown in Figure 6.
As observed in Table 2 and Figure 6, 1 M NaPF6 in EC:DMC exhibits the best electrochemical performance at temperatures above −20 °C. Upon further cooling, its Rct in NVP-NVP cells increases drastically (more than 30× growth from −20 °C to −40 °C), and 1 M NaPF6 in EC:DEC becomes a better choice for low-temperature batteries due to its Rct growing only by 1.7 times within the same temperature interval. Remarkably, in the case of DEC and DMC co-solvents, the slope of the Arrhenius 1/Rct plot depends less significantly on the liquid–solid phase transitions of the electrolyte than its ionic conductivity. This is most noticeable for 1 M NaPF6 in EC:DEC: the exponential plot of the temperature dependence of its 1/Rct remains almost linear over the entire range despite the liquid-solid phase transition. A similar phenomenon was observed earlier for commercial Li-ion cells [34]. This observation supports the notion that the charge transfer resistance does not consistently exhibit a direct relationship with the ionic conductivity of the electrolyte. Previously, it was reported that increasing the concentration of salts in the electrolytes for metal-ion batteries led to a decrease in the resistance at the electrode–electrolyte interface, although the maximum conductivity was generally registered for approximately 1 M solutions [35,36]. We believe that in our study, the amount of liquid phase remaining after partial freezing of 1 M NaPF6 in the EC:DEC electrolyte is higher than in the case of EC:DME and EC:DMC. Moreover, the latter must have been completely solid, according to the DMC-EC phase diagram. Different NaPF6 solvation energies in DEC and DME lead to the drastically different behavior of the charge transfer resistance at the cathode–electrolyte interface upon cooling.
In general, we can summarize as follows: in a number of works, the authors noted the fact that the process of charge transfer at the electrode–electrolyte interface, in particular, the (de)solvation of an alkali metal cation, is the main factor determining the low-temperature properties of metal-ion batteries [14,37,38,39,40]. From this point of view, the results obtained in our work for the 1 M NaPF6 in EC:DEC electrolyte look quite promising, since the increase in Rct upon cooling down to −40 °C is surprisingly moderate. Obviously, further development in this direction will require modification of the electrolyte, including the creation of ternary or quaternary mixtures of solvents, but it is already obvious that 1 M NaPF6 in EC:DEC is a good basis for low-temperature sodium-ion batteries.
4. Conclusions
We have demonstrated that the electrochemical properties of Na-ion electrolytes at low temperatures do not align precisely with their thermodynamic characteristics. Yet, there is some relationship between the DSC peaks and the dynamics of conductivity changes upon cooling. In addition, the relationship between the individual properties of solvents and the electrochemical properties of solutions does not appear to be direct either. Among the systems studied, 1 M NaPF6 in EC:DMC demonstrates the best performance from −20 °C to +20 °C, even though both solvents have melting points above 0 °C. However, further cooling induces the liquid-solid phase transition with a drastic decrease in ionic conductivity and growth in charge transfer resistance in Na3V2(PO4)3-based cells. Thus, within the range from −40 °C to −20 °C, 1 M NaPF6 in the EC:DEC electrolyte is the one demonstrating satisfactory performance, although it undergoes the liquid–solid transition at ≈−7 °C. We attribute this to the presence of some amount of liquid phase which wets the system and provides the necessary interface contact and diffusion paths for Na+. Interestingly, the situation is much worse in the case of DME, although this solvent also has a relatively low melting point and higher dielectric constants compared to DEC. Although the resistance of the EC:DME-based solutions is lower than that of the other two electrolytes at room temperature, lowering the temperature down to ≈−20 °C results in a very sharp increase in cell resistance, similar to what is observed in the case of the EC:DMC-based electrolyte.
Conceptualization, D.S.L. and O.A.D.; methodology, D.S.L., E.V.B. and O.A.D.; investigation, D.S.L., E.V.B. and M.V.Z.; writing—original draft preparation, D.S.L., E.V.B. and M.V.Z.; writing—review and editing, O.A.D. and E.V.A.; supervision, O.A.D. and E.V.A.; project administration, E.V.A.; funding acquisition, E.V.A. All authors have read and agreed to the published version of the manuscript.
All the data obtained is presented within the article.
Thermal analysis was performed under the state assignment «Chemical thermodynamics and theoretical materials science» (CITIS № 121031300039-1). Authors acknowledge support from M.V.Lomonosov Moscow State University Program of Development in part of DSC measurements and electrochemical experiments.
The authors declare no conflict of interest. The funders had no role in the design of the study; in the collection, analyses, or interpretation of data; in the writing of the manuscript; or in the decision to publish the results.
Footnotes
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Enlarge this image.Figure 1. General scheme of the experiments carried out in the work (a); an example of visual control of electrolyte freezing for 1 M NaPF6 in EC:DEC: liquid electrolyte and “wet snow” stage (b).
Enlarge this image.Figure 2. DSC curves for (a) 1 M NaPF6 in EC:DEC = 1:1; (b) 1 M NaPF6 in EC:DME = 1:1; and (c) 1 M NaPF6 in EC:DMC = 1:1. Cooling and heating rate—0.3 K/min. Snowflakes on the temperature scale mark the points of “wet snow”.
Enlarge this image.Figure 3. Specific conductivity of NaPF6 solutions in the EC:DEC, EC:DME, EC:DMC mixtures (1:1 by volume) depending on temperature (+20 °C… −40 °C): (a) 0.5 M NaPF6, (b) 1 M NaPF6. Snowflakes on the Celsius scale mark the points of “wet snow”.
Enlarge this image.Figure 4. Chronopotentiometry curves of NVP-NVP cells at a C/100 current density upon cooling and their respective “effective conductivity” values in Arrhenius coordinates for 1 M NaPF6 in (a,d) EC:DEC; (b,e) EC:DME; and (c,f) EC:DMC. Snowflakes on the Celsius scale mark the points of “wet snow”.
Enlarge this image.Figure 5. EIS data for 1 M NaPF6 in EC:DEC, in EC:DME, and in EC:DMC at (a) T = +20 °C; (b) T = 0 °C; (c) T = −20 °C; (d) T = −40 °C; (e) equivalent circuit: Rs—solution resistance; Rct—charge transfer resistance; CPE—constant phase element; M—restricted linear diffusion impedance.
Enlarge this image.Figure 6. Inverse resistance for the steel electrode and NVP-NVP cells containing solutions of (a) 1 M NaPF6 in EC:DEC = 1:1; (b) 1 M NaPF6 in EC:DME = 1:1; and (c) 1 M NaPF6 in EC:DMC = 1:1. Snowflakes on the temperature scale mark the points of “wet snow”.
Physical properties of solvents [
| Solvent | Structural Formula | Tm, °C | Tb, °C | Viscosity, mPa·s |
Dielectric Constant |
Density, g/cm3 |
|---|---|---|---|---|---|---|
| EC | [Image omitted. Please see PDF.] | 36.4 | 248.0 | 1.90 |
89.780 | 1.321 |
| DEC | [Image omitted. Please see PDF.] | −74.3 | 126.0 | 0.75 | 2.800 | 0.969 |
| DMC | [Image omitted. Please see PDF.] | 4.6 | 91.0 | 0.59 |
3.107 | 1.063 |
| DME | [Image omitted. Please see PDF.] | −58.0 | 84.0 | 0.46 | 7.200 | 0.860 |
Tm—melting temperature; Tb—boiling temperature; EC—ethylene carbonate; DEC—diethyl carbonate; DMC—dimethyl carbonate; DME—1,2-dimethoxyethane.
Initial crystallization temperatures.
| Solvent Mixture | C(NaPF6), M | T, °C |
|---|---|---|
| EC:DEC |
0.0 | 3.3 ± 1 |
| 0.5 | −5.2 ± 1 | |
| 1.0 | −7.8 ± 1 | |
| EC:DME |
0.0 | 14.0 ± 1 |
| 0.5 | −11.9 ± 1 | |
| 1.0 | −13.1 ± 1 | |
| EC:DMC |
0.0 | −2.7 ± 1 |
| 0.5 | −17.1 ± 1 | |
| 1.0 | −21.2 ± 1 |
Electrolyte and charge transfer resistance determined from EIS data for 1 M NaPF6 solutions.
| T, °C | Rs, Ohm | Rct, Ohm | ||||
|---|---|---|---|---|---|---|
| In EC:DEC | In EC:DME | In EC:DMC | In EC:DEC | In EC:DME | In EC:DMC | |
| +20 | 2.8 | 1.9 | 2.1 | 101.7 | 107.0 | 66.9 |
| +10 | 3.1 | 2.3 | 2.4 | 110.6 | 119.0 | 72.5 |
| 0 | 3.1 | 7.2 | 2.9 | 131.3 | 171.8 | 81.0 |
| −10 | 3.7 | 17.5 | 3.6 | 148.3 | 229.8 | 92.5 |
| −20 | 8.7 | 47.6 | 9.0 | 197.4 | 396.0 | 141.9 |
| −30 | 16.0 | 864.4 | 67.0 | 231.7 | 9707.0 | 398.2 |
| −40 | 38.8 | 7244.0 | 1002.0 | 329.1 | 179,666.0 | 5150.0 |
Rs—solution resistance; Rct—charge transfer resistance.
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; Drozhzhin, Oleg A 1
; Antipov, Evgeny V 2
1 Department of Chemistry, Lomonosov Moscow State University, GSP-1, 1-3 Leninskiye Gory, 119991 Moscow, Russia
2 Department of Chemistry, Lomonosov Moscow State University, GSP-1, 1-3 Leninskiye Gory, 119991 Moscow, Russia; Center for Energy Science and Technology, Skolkovo Institute of Science and Technology, Bolshoy Boulevard 30, 121205 Moscow, Russia