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State of charge indicators for alkaline zinc-air redox flow batteries
Journal of Power Sources 424 (2019) 76–81
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State of charge indicators for alkaline zinc-air redox flow batteries Christian Zelger a b
a,b
a
b
, Michael Süßenbacher , Andreas Laskos , Bernhard Gollas
T
a,∗
Institute for Chemistry and Technology of Materials, Graz University of Technology, Stremayrgasse 9, 8010, Graz, Austria CEST Competence Centre for Electrochemical Surface Technology GmbH, Viktor-Kaplan-Straße 2, 2700, Wiener Neustadt, Austria
H I GH L IG H T S
G R A P H I C A L A B S T R A C T
of charge indicators for alkaline • State zinc-air redox flow batteries are studied.
and conductivity correlate • Density with zincate concentration in NaOH and KOH.
rest potential varies linearly with • Zinc the logarithmic Nernst concentration term.
or a combination of these quan• One tities permit a simple state of charge monitoring.
A R T I C LE I N FO
A B S T R A C T
Keywords: State of charge Redox flow battery Zincate electrolyte Conductivity Density Rest potential
Physico-chemical properties of NaOH and KOH electrolytes are studied as possible state of charge indicators for the alkaline zinc-air redox flow battery. In order to predict the available energy of the battery and to prevent degradation as the result of unsuitable operating conditions, its state of charge has to be known. The state of charge is directly related to the composition of the negative electrolyte. A battery charging operation is simulated by electrodepositing zinc on a rotating cylinder electrode from NaOH electrolyte, while the simulation of a discharging operation is accomplished by dissolving ZnO in KOH electrolyte. The refractive index, conductivity, and density of NaOH- and KOH-electrolytes as well as the rest potential of the zinc electrode in these media are measured for different zincate and alkali concentrations at several temperatures. The electrolyte density and electrolyte conductivity are both linearly correlated with the zincate concentration. A linear correlation is found also between the rest potential and the logarithmic concentration term in the Nernst equation for the zinc halfcell reaction. These indicators permit a simple and straightforward monitoring of the state of charge of alkaline zinc-air redox flow batteries.
1. Introduction The increasing fraction of intermittent and difficult-to-predict renewable energy sources, like solar and wind power, requires high flexibility in power grid operation. Energy storage helps to achieve flexibility. It maximizes the utilization of generated power without affecting when and how consumers use it [1]. Among others, redox flow batteries (RFBs) are potential candidates in the field of grid-scale
∗
electrochemical energy storage. The all-vanadium and Zn-Br systems are the most developed RFB technologies in terms of market application [2]. In a typical RFB, like the all-vanadium RFB or Fe-Cr RFB [3], the electroactive species are not stored inside the electrodes. Therefore, the mechanical stress for the electrodes is reduced compared to other types of batteries, like e.g. conventional lithium-ion batteries. This helps to build long-lasting electrodes and systems [4]. The separation of the storage unit and power unit brings additional advantages of high design
Corresponding author. E-mail address: [email protected] (B. Gollas).
https://doi.org/10.1016/j.jpowsour.2019.03.099 Received 2 January 2019; Received in revised form 4 March 2019; Accepted 23 March 2019 0378-7753/ © 2019 Elsevier B.V. All rights reserved.
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flexibility and scalability. A disadvantage is the reduced energy efficiency due to the required electrolyte movement. In order to predict the available energy of the battery and to prevent side reactions, like electrolyte decomposition, the state of charge (SOC) has to be known. The open-circuit voltage (OCV) measurement has been widely used for determining the SOC in commercial all-vanadium RFB's [5–8]. Measurement of the half-cell electrolyte potentials [9], the electrolyte conductivity [10,11], spectrophotometric properties (UV-vis [6,10,12–14], IR [15]) and ultrasonic velocity [16] were also proposed for the all-vanadium RFB. For the Zn-Br RFB, electrolyte conductivity [17] and Raman spectroscopy [18,19] were suggested. In small commercial RFB systems a complex and expensive SOC determination is uneconomic and not practicable. Therefore, a low-cost and simple way of monitoring the SOC is preferred. The zinc-air RFB is particularly well suited for stationary energy storage, because its active materials are low cost, environment-friendly and abundant [1,20,21]. The reversible air electrode in one half-cell reaction supersedes the storage of the positive active material, which results in an increased energy density of the RFB. The half-cell reactions and the overall electrochemical reaction of the alkaline zinc-air RFB with a soluble zinc electrode are shown in Eqs. (1)–(3). During discharge, zincate ions [Zn(OH)4]2− are dissolved into the negative electrolyte (Eq. (1)) [22]. According to Eq. (4), hydrogen evolution occurs as a side reaction at the negative electrode. The standard electrode potentials of Eqs. (1), (2) and (4) are referred to pH = 14 [23]. [Zn(OH)4]2− + 2e− ⇆ Zn + 4OH− E0 = −1.285 V vs. SHE 2OH
−
⇆ ½O2 + H2O + 2e 2−
[Zn(OH)4]
2H2O + 2e
−
−
0
E = +0.401 V vs. SHE −
⇆ Zn + ½O2 + 2OH −
⇆ H2 + 2OH
ZnO + H2O + 2 OH− ⇆ [Zn(OH)4]2− −
CO2 + 2 OH
E = −0.828 V vs. SHE
(4)
RT a ([Zn(OH)4 ]2 − ) ·ln nF a (OH−)2⋅a (H2 O)⋅a (O2 )0.5
(7)
Electrolytes were prepared with potassium hydroxide (≥85%, p. a.) or sodium hydroxide (≥99%, p. a.) from Carl Roth GmbH and zinc oxide from Sigma-Aldrich (> 99%, puriss.). The organic electrolyte additive Lugalvan P (LP) is a commercial product of BASF SE. First, 7 M or 8 M solutions of NaOH or KOH were prepared, then zinc oxide and the electrolyte additive were dissolved. A zinc sheet (99.99%) from Advent Research Materials Ltd. was used as the negative electrode in experiments with additive-free KOH electrolytes. The inner electrolyte of the Hg/HgO reference electrode was 1 M or 8 M KOH. The respective concentration is indicated in the figures. The rotating cylinder electrode experiment was performed with a HT Rota-Hull cell from Eco Chemie B. V. (The Netherlands). The geometry of the cell was described in detail by Low et al. [32]. The cathode was a brass cylinder with 6 mm diameter, a length of 12 cm and an electrode area of 15 cm2. A cylindrical platinized titanium mesh with an inner diameter of 52 mm and a height of 23 mm served as anode. No current shield was used with the setup to ensure a uniform current density distribution along the cylinder cathode. The zincate concentration in the electrolyte was determined at regular time intervals by X-ray fluorescence (Model XAN-FD, Helmut Fischer AG) during zinc deposition on the rotating cylinder electrode. The refractive index was measured with an Abbe refractometer at a wavelength of 589 nm. At this wavelength the 7 M NaOH/ZnO/2.4 g L−1 LP electrolyte is optically transparent. The electrolyte conductivity was measured with a conductivity meter (Model 703) from Knick GmbH & Co. The NaOH electrolyte density was determined by weighing a defined electrolyte volume, while a Schott pycnometer (Germany) was used for the KOH electrolyte. The corrosion current density was determined from gravimetric measurements. Eight zinc sheets were stored in a glass beaker containing the temperature-controlled electrolyte solution, which was magnetically stirred and heated on an electric heating plate (IKA basic). The zinc sheets (99.99%) from Advent Research Materials Ltd. had geometric surface areas of 31 cm2 (NaOH electrolyte) or 18 cm2 (KOH electrolyte). The volume of the NaOH and KOH electrolytes were 1.8 L and 0.925 L, respectively. At selected times, a zinc sheet was removed from the electrolyte solution and the mass loss was measured after careful rinsing and drying. Each data point in these experiments was generated from one zinc sheet. The electrolyte solutions were open to ambient air during the experiment to allow also for oxygen reduction as cathodic corrosion reaction.
The formation of zinc dendrites and the hydrogen evolution side reaction in the zinc-air RFB can be prevented or minimized, if the SOC is monitored. Hydrogen evolution occurs to a large extent at the negative electrode during charging, if the current approaches the limiting current of zinc deposition. The latter conditions are also favourable for the growth of zinc dendrites [24,25]. Zinc dendrites can cause cell shortening. Apart from dendritic zinc, also other non-compact morphologies, like filamentous mossy and heavy spongy zinc poorly adhere to the substrate [22,26]. Zinc with these morphologies can easily detach from the electrode in the streaming electrolyte, which would result in a loss of capacity and/or obstruction of the electrolyte flow in the battery. The limiting current for zinc deposition depends on the zincate concentration and the electrolyte movement. Therefore, the operating limits of the negative electrolyte can be determined from the zincate concentration, which is correlated to the SOC. According to the Nernst Eq. (5), with the electrode potential E, the standard electrode potential E0, the universal gas constant R, the absolute temperature T, the number of transferred electrons per atom n and the Faraday constant F, the rest potential of the zinc electrode depends on the activities a of hydroxide, water, oxygen, and zincate ions. Therefore, the latter potential could be an indicator of the SOC and is reflected also in the OCV of the cell or battery. In aqueous alkaline electrolytes zinc corrodes [27,28], which results in a mixed potential (corrosion potential).
E = E0 +
+ H2O
(6)
2. Experimental
(2) (3)
⇆
CO32−
Here, we present the results of a fundamental investigation on state of charge indicators for the alkaline zinc-air redox flow battery. The electrolyte temperature in this study has been varied between room temperature and 60 °C, because many precious-metal-free air electrodes suitable for an alkaline zinc-air RFB show an enhanced voltage efficiency at elevated temperature [31].
(1)
+ H2O
0
molecules contained in a defined volume. When some of these molecules are replaced by other molecules, e.g. zincate ions, the electrolyte density can change. SOC indicators are affected by a change of the electrolyte composition. The electrolyte composition may change by ageing of the battery. In the zinc-air RFB, changes in electrolyte composition are caused by water evaporation, precipitation of zinc oxide (Eq. (6)), decomposition/loss of electrolyte additives or uptake of carbon dioxide from the air [29,30] according to Eq. (7).
(5)
Also other electrolyte properties are affected by the zincate concentration: the ionic strength and the electrical conductivity of the negative electrolyte, as well as its density and viscosity. The ability of an aqueous electrolyte to conduct electricity is based on the transport of solvated ions in the electric field between the electrodes. The mobility of the solvated ions and their concentration determine the electric conductivity. The electrolyte density is the weight of a defined electrolyte volume. That weight is the sum of the weight of the individual 77
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Fig. 3. Ionic conductivity σ at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−).
Fig. 1. Refractive index n at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−).
3. Results and discussion 3.1. NaOH electrolyte A battery charging operation has been simulated in a rotating cylinder electrode experiment. Zinc was electrodeposited on the rotating cylinder cathode, while oxygen was evolved at the platinized titanium counter electrode analogous to the processes in a zinc-air RFB (Eq. (3)). During the experiment, the refractive index (Fig. 1), density (Fig. 2), and conductivity (Fig. 3) of the electrolyte as well as the zinc rest potential (Fig. 4) were measured at regular time intervals. The experiment was performed in 0.425 L of a 7 M NaOH electrolyte containing 0.9 M dissolved ZnO and 2.4 g L−1 Lugalvan P (LP) additive at 60 °C, at a current density of 50 mA cm−2, and a rotational rate of 160 rpm and lasted about 27 h. LP is an additive in industrial alkaline baths for zinc plating and was used here to prevent the growth of dendritic zinc. The refractive index n shows only a small variation between 1.3915 and 1.3960 (Fig. 1) and there seems to be little correlation, if any, with the zincate concentration c([Zn(OH)4]2-). The coefficient of determination R2 (Pearson correlation coefficient) of the linear regression is as low as 0.36. Therefore, the refractive index is not suitable for determining the SOC in the 7 M NaOH/2.4 g L−1 LP electrolyte. As expected, the electrolyte density ρ decreases during the experiment with decreasing zincate concentration by 0.036 g cm−3, from 1.276 to 1.240 g cm−3 (Fig. 2). This behaviour was also observed by Siu and Evans [33]. The correlation between density and zincate
Fig. 4. Measured mixed potential and calculated equilibrium potential of the zinc electrode in a 7 M NaOH electrolyte with 2.4 g L−1 LP additive as a function of ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2) at 60 °C. The measured data are at the same zincate concentrations c([Zn(OH)4]2−) as those in Fig. 1, Fig. 2, and Fig. 3.
concentration in our experiment was found to be highly linear with a slope of 0.045 g cm−3 mol−1 L and a coefficient of determination R2 of 0.96. Hence, the electrolyte density appears to be useful for determining the SOC in 7 M NaOH/2.4 g L−1 LP electrolyte. Moreover, sensors for monitoring the density of alkaline solutions are commercially available. One has to keep in mind, however, that density variations can be caused also by concentration changes of species other than zincate, e.g. NaOH [33] and carbonates, Eq. (7), and by temperature variations [32]. Also aging of the battery may thus change the electrolyte composition and thereby affect the SOC-determination. The electric conductivity σ increases from 561 to 680 mS cm−1 during the charging experiment (Fig. 3). The increase in conductivity with decreasing zincate concentration can be explained by the high ionic strength, Ic, of the electrolyte, Eq. (8).
Ic =
1 ⋅ ∑ ci⋅z i2 2 i
(8)
Here, ci is the ion concentration of species i and zi is its ionic charge. At high concentrations, ions strongly interact and might even form ion pairs. This decreases both the mobility and the concentration of dissociated ionic charge carriers affecting the ionic conductivity. During charging, one zincate ion is reduced to metallic zinc and two hydroxide ions plus a water molecule are generated according to equation (3). The
Fig. 2. Density ρ at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−). 78
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Figs. 4 and 13.9 mV and 14.3 mV, respectively, are nevertheless in very good agreement. Electrolyte density and conductivity correlate linearly with the zincate concentration, as does the zinc rest potential with the Nernstian term ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2) in the 7 M NaOH/ 2.4 g L−1 LP. If the total concentrations of water, OH−, and the additive are known, the zincate concentration can be calculated from the rest potential of the zinc electrode in a calibration plot. All three correlations have a coefficient of determination R2 (Pearson correlation coefficient) equal or larger than 0.96 and appear to be suitable for determining the SOC in this electrolyte. An attempt was made to determine the zincate concentration from voltammetry with ultra-microelectrodes. However, no steady-state currents for the reduction of zincate could be observed in the 7 M NaOH/0.8 M ZnO/2.4 g L−1 LP electrolyte on platinum and carbon fibre ultra-microelectrodes. Also, the consumed charge for zincate reduction was not reproducible in these experiments. Thus, amperometry with ultra-microelectrodes is unsuitable for determining the SOC.
ionic strength decreases, because it depends on the square of the ionic charges and the zincate ion is doubly charged. Moreover, due to the Grotthuss mechanism [34], hydroxide ions contribute more strongly to the electric conductivity than zincate ions do. Therefore, the conductivity increases if zincate ions are replaced by the two-fold number of hydroxide ions. A linear correlation with a coefficient of determination R2 of 0.99 was found between the electrolyte conductivity and the zincate concentration (Fig. 3). The ionic conductivity can thus be used for determining the SOC in the 7 M NaOH/2.4 g L−1 LP electrolyte. Also in this case, however, one should bear in mind that conductivity variations can be caused also by concentration changes of ions other than zincate and by temperature changes. The rest potential of the zinc electrode shows a variation of 29 mV for a change of the zincate concentration from 0.18 M to 0.83 M (Fig. 4). The measured potentials are shown together with the equilibrium potentials calculated from the Nernst Eq. (5) using concentrations instead of activities. The activity of oxygen dissolved in the electrolyte is assumed to be constant in the open cell and is therefore neglected in the calculations. The first equilibrium potential was measured after the brass electrode was completely coated with zinc at c([Zn (OH)4]2−) = 0.83 M and c(OH−) = 5.3 M. The initial OH− concentration was calculated from Eq. (6), while the varying OH− concentration during the charging experiment was calculated from Eq. (3). The water concentrations were calculated from the measured densities of the electrolyte subtracting the weight of NaOH, ZnO, and LP and taking into account the consumption of one water per ZnO according to Eq. (6). In the calculation of the equilibrium potential, the standard potential given for the half-cell reaction in Eq. (1) is used and converted to the potential scale of the Hg/HgO reference electrode (+0.1154 V vs. SHE @ 1 M KOH). The universal gas constant R is 8.314 J mol−1 K−1, T = 333 K (60 °C), n = 2 and F = 96485 A s mol−1. The difference between the measured and the calculated potentials in Fig. 4 can be explained with two effects. Firstly, zinc corrodes in this electrolyte and hence, the measured potential is the free corrosion potential of zinc for the reactions in Eqs. (3) and (9) [35–37]. A corrosion current density of 36 μA cm−2 has been determined for the 0.2 M zincate/7 M NaOH/2.4 g L−1 LP electrolyte at 60 °C from gravimetric measurements shown in Fig. S1 of the supplementary information. From the mass loss Δm over time, a linear regression was generated and used to calculate the corrosion current density jcorr from Eq. (10). Here, k is the slope of the linear regression in Fig. S1 and the molar mass M of zinc is 65.39 g mol−1. The corrosion current density in the 0.2 M zincate/8 M KOH electrolyte at 60 °C, which is discussed in section 3.2, was found to be 6.4 μA cm−2. Zn + 2H2O + 2 OH− → H2 + [Zn(OH)4]2−
jcorr =
k ⋅F ⋅n 3600 M
3.2. KOH electrolyte Electrolyte conductivity, density and the zinc rest potential have been studied in additive-free 7.5 M, 8.0 M, and 8.5 M KOH electrolytes for zincate concentrations between 0.2 and 0.6 M at temperatures of 40, 60, and 70 °C (Fig. 5, Fig. 6, Fig. 7). In the rotating cylinder electrode experiment with NaOH electrolyte discussed above, the zincate concentration was varied electrochemically, according to Eq. (3). In the experiment with KOH electrolyte, the zincate concentration was varied by dissolution of ZnO, according to Eq. (6). The comparison of Eqs. (3) and (6) shows that dissolution of ZnO leads to the same variation of the electrolyte composition as does the electrochemical dissolution of zinc. Therefore, battery discharging can be simulated by simple dissolution of ZnO in these electrolytes. The electric conductivity σ decreases from 982 to 861 mS cm−1, for an increase in zincate concentration from 0.2 M to 0.6 M in 8 M KOH electrolyte at 60 °C (Fig. 5). The decrease in conductivity with increasing zincate concentration is explained by the high ionic strength of the electrolyte and this behaviour was already discussed for the NaOH electrolyte in section 3.1. The ionic conductivity of KOH solutions under varying conditions was reviewed previously by Gilliam et al. [38]. They showed that the conductivity decreases, if the KOH concentration increases from 7 M to 8 M at 60 °C (Fig. 3 in Ref. [38]). This behaviour is in agreement with the results of this work. Fig. 5 shows that the ionic conductivity is affected by the electrolyte temperature
(9) (10)
Since the measured potential of the zinc electrode is a mixed potential, it deviates from the equilibrium potential calculated from Eq. (5). The calculation of the corrosion potential is complex. Based on the corrosion current density the corrosion potential could be calculated from the kinetic parameters of anodic and cathodic reactions (Tafel slopes and exchange current densities). However, the kinetic parameters of these reactions under the present experimental conditions are not known. Secondly, no activity data is available for the present electrolyte. Instead of ionic activities in Eq. (5) that could be calculated from approximated activity coefficients [36], concentrations have been used. The difference between the measured and the calculated curve in Fig. 4 is thus explained by i) the neglected impact of the zinc corrosion and ii) the deviation of the activity coefficients from unity. Both corrosion potential and activity coefficients depend on the electrolyte composition and thus vary with the concentrations of zincate, water, and OH−. The presence of the plating additive might affect them as well. The slopes of the measured and of the calculated potentials in
Fig. 5. Electric conductivity σ of KOH electrolytes as a function of zincate concentration c([Zn(OH)4]2−) for different KOH concentrations and temperatures. 79
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Fig. 8. Measured zinc rest potential (square points) from Fig. 7 (60 °C, 8 M KOH) and calculated equilibrium potential as function of ln(c([Zn(OH)4]2−)·c (H2O)−1·c(OH−)−2). The numbers on the data points indicate the zincate concentration.
Fig. 6. Electrolyte density ρ at 60 °C of a 8 M KOH electrolyte at different zincate concentrations c([Zn(OH)4]2−).
zincate concentration from 0.2 M to 0.6 M, the calculated KOH density increases by 0.025 g cm−3 (Fig. 6). The linear regression of the calculated density shows a coefficient of determination R2 of 0.999, which makes the electrolyte density a good indicator for measuring the SOC in 8 M KOH electrolyte. ρ = a + b w(KOH) + c w(KOH)2 + d w(ZnO) + e T
w (ZnO) = 8.143
c (ZnO) ρ
(11)
(12)
The rest potential of the zinc electrode varies by 15–20 mV for a change of zincate concentration from 0.2 M to 0.6 M (Fig. 7). The shift of the potentials with KOH concentration is in qualitative agreement with the Nernst equation, if unity activity coefficients are assumed. This is not the case for the slope in the regression of the data at 40 °C, which should be smaller than for the data at 60 °C. The slight deviations are exemplified in Fig. 8, where the measured rest potential is shown together with the calculated equilibrium potential calculated from Eq. (5), analogously to the rotating cylinder electrode experiment in section 3.1. Zinc corrodes also in this electrolyte, the corrosion current density being 6.4 μA cm−2 at 60 °C (section 3.1). Therefore, the measured rest potential is again a mixed potential (corrosion potential), as in the case of the NaOH electrolyte shown in the previous section. The kinetic parameters of the concentration and temperature dependent anodic and cathodic corrosion reactions are not known, and therefore, the corrosion potential could not be determined. Instead, the equilibrium potential was calculated for a comparison with the measured mixed potential in Fig. 8. The OH− and water concentrations in Fig. 8 were calculated from Eq. (6) (dissolution of ZnO). The equilibrium potential was calculated from the standard potential given for Eq. (1) and converted to the potential scale of the Hg/HgO reference electrode (+0.0624 V vs. SHE @ 8 M KOH). The potential of the reference electrode is shifted negatively by 53 mV, when the inner electrolyte is changed from 1 M to 8 M KOH. This potential shift is calculated from the Nernst equation and was also confirmed experimentally. Again, ionic concentrations have been used rather than ionic activities in Eq. (5) and the constant activity of dissolved oxygen omitted for the same reasons as explained in section 3.1. The calculated density from Eq. (11). was used for the calculation of the water concentration. The difference between the measured and the calculated data in Fig. 8 arises from the neglected impact of the zinc corrosion and the deviation of the ionic activity coefficients from unity. The slopes of the data sets in Fig. 8 are again similar, 11.9 mV for the linear regression of the measured potential and 16.7 mV for the calculated potential.
Fig. 7. Rest potential E of the zinc electrode as function of ln(c([Zn(OH)4]2−)·c (H2O)−1·c(OH−)−2) at different KOH concentrations and temperatures. The lines show the linear regressions. The measurements were carried out at the same zincate concentrations c([Zn(OH)4]2−) as those shown in Fig. 5.
and KOH concentration. A temperature decrease from 60 to 40 °C lowers the conductivity by 250 mS cm−1 at a zincate concentration of 0.2 M. The linear correlations between electrolyte conductivity and zincate concentration in Fig. 5 show a coefficient of determination R2 of 0.99. If the conductivity is used for determining the SOC, one has to bear in mind though that variations can be caused also by concentration changes of ions other than zincate (e.g. potassium, hydroxide, carbonates) and by temperature changes. The electrolyte densities for different zincate concentrations have been calculated with Eqs. (11) and (12), as suggested by Siu and Evans [33] and the results are shown in Fig. 6. The KOH weight concentration w(KOH) of 34.0% (8 M) was calculated according to Gilliam et al. (Equation (1) and Table 4 in Ref. [38] for a temperature of 60 °C). In Eq. (11), w(ZnO) is the weight concentration of zinc oxide, the temperature T is 60 °C, and the coefficients a to e are 1.01137, 0.008466, 0.00003071, 0.01039, −0.0005106 (Table I in Ref. [33], column “CRC data included, initial KOH concentration”). For comparison, also two measured data points are shown in Fig. 6. The measured data is shifted to smaller values by approximately 0.030 g cm−3, which might be due to a small error in either w(KOH) or w(ZnO). However, the slope of 0.077 g cm−3 mol−1 L in the measured data is close to that of 0.062 g cm−3 mol−1 L in the calculated data. The KOH electrolyte density drops with decreasing zincate concentration, which was found also for the NaOH electrolyte (section 3.1). For an increase of the 80
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acknowledge funding (no. 848933) by the Climate and Energy Fund of the Austrian Federal Government through the 1st Call “Energieforschung”.
Density and ionic conductivity correlate linearly with the zincate concentration in the KOH electrolyte, as does the rest potential with the term ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2). The coefficient of determination R2 (Pearson correlation coefficient) is in all cases equal or larger than 0.99. Density and electric conductivity of the electrolyte as well as the rest potential can thus be used for determining the SOC in 8 M KOH electrolyte.
Appendix A. Supplementary data Supplementary data to this article can be found online at https:// doi.org/10.1016/j.jpowsour.2019.03.099.
4. Conclusions References Several possible state of charge indicators for the alkaline zinc-air redox flow battery have been studied during a battery charging operation simulated in a rotating cylinder electrode experiment with 7 M NaOH/ZnO/2.4 g L−1 LP electrolyte. The measurements show no significant correlation between refractive index and zincate concentration. However, density and conductivity of the NaOH electrolyte correlate linearly with the zincate concentration, while the rest potential correlates linearly with the Nernstian term ln(c([Zn(OH)4]2−)·c(H2O)−1·c (OH−)−2). Also in the case of the 8 M KOH electrolyte, density and conductivity correlate linearly with the zincate concentration, and the zinc rest potential correlates linearly with the Nernstian term ln(c([Zn (OH)4]2−)·c(H2O)−1·c(OH−)−2). In all linear regressions of density, conductivity, and rest potential for both NaOH and KOH electrolytes, the coefficient of determination R2 is either equal to 0.96 or in most cases even around 0.99. Such strong linear correlations between measured parameters and zincate concentration allow for a simple determination of the SOC. The density, conductivity, and rest potential of NaOH and KOH zincate electrolytes appear to be reliable indicators of the SOC of the zinc-air redox flow battery, if the temperature is kept constant or measured simultaneously. The formation of carbonate in the alkaline electrolytes upon absorption of CO2 might affect the three indicators. Furthermore, degradation of battery components such as cathode materials (catalysts, conductive support, binders, and current collector), materials for cell housing, and possibly membranes could result in accumulation of impurities that might also affect the electrolyte properties. These processes, however, occur slowly on a long timescale and would be an issue of battery maintenance. They depend on the specific battery installation and an assessment of their possible impact on the respective indicators thus requires thorough long-term studies. The variation of the rest potential is 29 mV in the NaOH electrolyte and 20 mV in KOH, and thus rather small in the investigated SOC range. Such small potential differences are difficult to measure. The conductivity varies by approximately 125 mS cm−1 in the NaOH and KOH electrolytes. Variations of such a magnitude are much simpler to measure than the density or rest potential. The reliability of the SOC monitoring could nevertheless be improved even further by combining two of the indicators.
[1] C.J. Barnhart, S.M. Benson, Energy Environ. Sci. 6 (2013) 1083–1092. [2] D. Rastler, Electric Power Research Institute Technical Report, (2010), p. 1020676. [3] P. Leung, X. Li, d.L. Ponce, L. Berlouis, C.T.J. Low, F.C. Walsh, RSC Adv. 2 (2012) 10125–10156. [4] T. Nguyen, R.F. Savinell, Electrochem. Soc. Interface. 19 (Fall) (2010) 54–56. [5] K.W. Knehr, E.C. Kumbur, Electrochem. Commun. 13 (2011) 342–345. [6] Z. Tang, D.S. Aaron, A.B. Papandrew, T.A. Zawodzinski, ECS Trans. 41 (2012) 1–9. [7] M.R. Mohamed, H. Ahmad, M.N. Abu Seman, Elektron Elektrotech 19 (2013) 37–42. [8] K. Ngamsai, A. Arpornwichanop, J. Power Sources 282 (2015) 534–543. [9] S. Corcuera, M. Skyllas-Kazacos, Eur. Chem. Bull. 1 (2012) 511–519. [10] M. Skyllas-Kazacos, M. Kazacos, J. Power Sources 196 (2011) 8822–8827. [11] K. Ngamsai, A. Arpornwichanop, J. Power Sources 298 (2015) 150–157. [12] R.P. Brooker, C.J. Bell, L.J. Bonville, H.R. Kunz, J.M. Fenton, J. Electrochem. Soc. 162 (2015) A608–A613. [13] W. Zhang, L. Liu, L. Liu, RSC Adv. 5 (2015) 100235–100243. [14] C. Petchsingh, N. Quill, J.T. Joyce, D.N. Eidhin, D. Oboroceanu, C. Lenihan, X. Gao, R.P. Lynch, D.N. Buckley, J. Electrochem. Soc. 163 (2016) A5068–A5083. [15] S. Rudolph, U. Schröder, I.M. Bayanov, K. Blenke, D. Hage, J. Electroanal. Chem. 694 (2013) 17–22. [16] Y. Chou, N. Hsu, K. Jeng, K. Chen, S. Yen, Appl. Energy 182 (2016) 253–259. [17] J.J. Bolstad, R.C. Miles, Proc. - Intersoc. Energy Convers. Eng. Conf. 24th (1989) 1311–1318. [18] W. Pell, Zinc/bromine Battery Electrolytes: Electrochemical, Physicochemical and Spectroscopic Studies, Ph.D thesis University of Ottawa, 1994. [19] H.J. Lee, D. Kim, J.H. Yang, J. Electrochem. Soc. 164 (2017) A754–A759. [20] C. Ponce de León, A. Frías-Ferrer, J. González-García, D.A. Szánto, F.C. Walsh, J. Power Sources 160 (2006) 716–732. [21] G.X. Zhang, J. Power Sources 285 (2015) 580–587. [22] C. Zelger, J. Laumen, A. Laskos, B. Gollas, Electrochim. Acta 213 (2016) 208–216. [23] A.F. Holleman, N. Wiberg, Lehrbuch der Anorganischen Chemie, 101st ed., Walter de Gruyter & Co., 1995. [24] R.Y. Wang, D.W. Kirk, G.X. Zhang, J. Electrochem. Soc. 153 (2006) C357–C364. [25] K. Wang, P. Pei, Z. Ma, H. Chen, H. Xu, D. Chen, X. Wang, J. Mater. Chem. 3 (2015) 22648–22655. [26] G.X. Zhang, Secondary batteries - zinc systems, zinc electrodes, in: OverviewIn, J. Garche (Eds.), Encyclopedia of Electrochemical Power Sources, first ed., Elsevier, 2009, pp. 384–392. [27] A.R. Suresh Kannan, S. Muralidharan, K.B. Sarangapani, V. Balaramachandran, V. Kapali, J. Power Sources 57 (1995) 93–98. [28] D.P. Gregory, P.C. Jones, D.P. Redfearn, J. Electrochem. Soc. 119 (1972) 1288–1292. [29] Y. Sato, H. Niki, T. Takamura, J. Electrochem. Soc. 118 (1971) 1269–1272. [30] G.F. McLean, T. Niet, S. Prince-Richard, N. Djilali, Int. J. Hydrogen Energy 27 (2002) 507–526. [31] D. Pletcher, X. Li, S.W.T. Price, A.E. Russell, T. Sönmez, S.J. Thompson, Electrochim. Acta 188 (2016) 286–293. [32] C.T.J. Low, E.P.L. Roberts, F.C. Walsh, Electrochim. Acta 52 (2007) 3831–3840. [33] S. Siu, J.W. Evans, J. Electrochem. Soc. 144 (1997) 1278–1280. [34] S. Cukierman, Biochim. Biophys. Acta Bioenerg. 1757 (2006) 876–885. [35] G.X. Zhang, Corrosion and Electrochemistry of Zinc, first ed., Springer, New York, 1996, p. 474. [36] J.O. Bockris, Z. Nagy, A. Damjanovic, J. Electrochem. Soc. 119 (1972) 285–295. [37] W.L. Fielder, NASA Tech. Memo. 79235 (1979). [38] R.J. Gilliam, J.W. Graydon, D.W. Kirk, S.J. Thorpe, Int. J. Hydrogen Energy 32 (2007) 359–364.
Acknowledgements The support by the European Commission (Theme 2010.7.3.1: Energy storage systems for power distribution networks, no. 256759) is gratefully acknowledged for parts of this work. C.Z., M.S., and B.G.
81
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Journal of Power Sources journal homepage: www.elsevier.com/locate/jpowsour
State of charge indicators for alkaline zinc-air redox flow batteries Christian Zelger a b
a,b
a
b
, Michael Süßenbacher , Andreas Laskos , Bernhard Gollas
T
a,∗
Institute for Chemistry and Technology of Materials, Graz University of Technology, Stremayrgasse 9, 8010, Graz, Austria CEST Competence Centre for Electrochemical Surface Technology GmbH, Viktor-Kaplan-Straße 2, 2700, Wiener Neustadt, Austria
H I GH L IG H T S
G R A P H I C A L A B S T R A C T
of charge indicators for alkaline • State zinc-air redox flow batteries are studied.
and conductivity correlate • Density with zincate concentration in NaOH and KOH.
rest potential varies linearly with • Zinc the logarithmic Nernst concentration term.
or a combination of these quan• One tities permit a simple state of charge monitoring.
A R T I C LE I N FO
A B S T R A C T
Keywords: State of charge Redox flow battery Zincate electrolyte Conductivity Density Rest potential
Physico-chemical properties of NaOH and KOH electrolytes are studied as possible state of charge indicators for the alkaline zinc-air redox flow battery. In order to predict the available energy of the battery and to prevent degradation as the result of unsuitable operating conditions, its state of charge has to be known. The state of charge is directly related to the composition of the negative electrolyte. A battery charging operation is simulated by electrodepositing zinc on a rotating cylinder electrode from NaOH electrolyte, while the simulation of a discharging operation is accomplished by dissolving ZnO in KOH electrolyte. The refractive index, conductivity, and density of NaOH- and KOH-electrolytes as well as the rest potential of the zinc electrode in these media are measured for different zincate and alkali concentrations at several temperatures. The electrolyte density and electrolyte conductivity are both linearly correlated with the zincate concentration. A linear correlation is found also between the rest potential and the logarithmic concentration term in the Nernst equation for the zinc halfcell reaction. These indicators permit a simple and straightforward monitoring of the state of charge of alkaline zinc-air redox flow batteries.
1. Introduction The increasing fraction of intermittent and difficult-to-predict renewable energy sources, like solar and wind power, requires high flexibility in power grid operation. Energy storage helps to achieve flexibility. It maximizes the utilization of generated power without affecting when and how consumers use it [1]. Among others, redox flow batteries (RFBs) are potential candidates in the field of grid-scale
∗
electrochemical energy storage. The all-vanadium and Zn-Br systems are the most developed RFB technologies in terms of market application [2]. In a typical RFB, like the all-vanadium RFB or Fe-Cr RFB [3], the electroactive species are not stored inside the electrodes. Therefore, the mechanical stress for the electrodes is reduced compared to other types of batteries, like e.g. conventional lithium-ion batteries. This helps to build long-lasting electrodes and systems [4]. The separation of the storage unit and power unit brings additional advantages of high design
Corresponding author. E-mail address: [email protected] (B. Gollas).
https://doi.org/10.1016/j.jpowsour.2019.03.099 Received 2 January 2019; Received in revised form 4 March 2019; Accepted 23 March 2019 0378-7753/ © 2019 Elsevier B.V. All rights reserved.
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flexibility and scalability. A disadvantage is the reduced energy efficiency due to the required electrolyte movement. In order to predict the available energy of the battery and to prevent side reactions, like electrolyte decomposition, the state of charge (SOC) has to be known. The open-circuit voltage (OCV) measurement has been widely used for determining the SOC in commercial all-vanadium RFB's [5–8]. Measurement of the half-cell electrolyte potentials [9], the electrolyte conductivity [10,11], spectrophotometric properties (UV-vis [6,10,12–14], IR [15]) and ultrasonic velocity [16] were also proposed for the all-vanadium RFB. For the Zn-Br RFB, electrolyte conductivity [17] and Raman spectroscopy [18,19] were suggested. In small commercial RFB systems a complex and expensive SOC determination is uneconomic and not practicable. Therefore, a low-cost and simple way of monitoring the SOC is preferred. The zinc-air RFB is particularly well suited for stationary energy storage, because its active materials are low cost, environment-friendly and abundant [1,20,21]. The reversible air electrode in one half-cell reaction supersedes the storage of the positive active material, which results in an increased energy density of the RFB. The half-cell reactions and the overall electrochemical reaction of the alkaline zinc-air RFB with a soluble zinc electrode are shown in Eqs. (1)–(3). During discharge, zincate ions [Zn(OH)4]2− are dissolved into the negative electrolyte (Eq. (1)) [22]. According to Eq. (4), hydrogen evolution occurs as a side reaction at the negative electrode. The standard electrode potentials of Eqs. (1), (2) and (4) are referred to pH = 14 [23]. [Zn(OH)4]2− + 2e− ⇆ Zn + 4OH− E0 = −1.285 V vs. SHE 2OH
−
⇆ ½O2 + H2O + 2e 2−
[Zn(OH)4]
2H2O + 2e
−
−
0
E = +0.401 V vs. SHE −
⇆ Zn + ½O2 + 2OH −
⇆ H2 + 2OH
ZnO + H2O + 2 OH− ⇆ [Zn(OH)4]2− −
CO2 + 2 OH
E = −0.828 V vs. SHE
(4)
RT a ([Zn(OH)4 ]2 − ) ·ln nF a (OH−)2⋅a (H2 O)⋅a (O2 )0.5
(7)
Electrolytes were prepared with potassium hydroxide (≥85%, p. a.) or sodium hydroxide (≥99%, p. a.) from Carl Roth GmbH and zinc oxide from Sigma-Aldrich (> 99%, puriss.). The organic electrolyte additive Lugalvan P (LP) is a commercial product of BASF SE. First, 7 M or 8 M solutions of NaOH or KOH were prepared, then zinc oxide and the electrolyte additive were dissolved. A zinc sheet (99.99%) from Advent Research Materials Ltd. was used as the negative electrode in experiments with additive-free KOH electrolytes. The inner electrolyte of the Hg/HgO reference electrode was 1 M or 8 M KOH. The respective concentration is indicated in the figures. The rotating cylinder electrode experiment was performed with a HT Rota-Hull cell from Eco Chemie B. V. (The Netherlands). The geometry of the cell was described in detail by Low et al. [32]. The cathode was a brass cylinder with 6 mm diameter, a length of 12 cm and an electrode area of 15 cm2. A cylindrical platinized titanium mesh with an inner diameter of 52 mm and a height of 23 mm served as anode. No current shield was used with the setup to ensure a uniform current density distribution along the cylinder cathode. The zincate concentration in the electrolyte was determined at regular time intervals by X-ray fluorescence (Model XAN-FD, Helmut Fischer AG) during zinc deposition on the rotating cylinder electrode. The refractive index was measured with an Abbe refractometer at a wavelength of 589 nm. At this wavelength the 7 M NaOH/ZnO/2.4 g L−1 LP electrolyte is optically transparent. The electrolyte conductivity was measured with a conductivity meter (Model 703) from Knick GmbH & Co. The NaOH electrolyte density was determined by weighing a defined electrolyte volume, while a Schott pycnometer (Germany) was used for the KOH electrolyte. The corrosion current density was determined from gravimetric measurements. Eight zinc sheets were stored in a glass beaker containing the temperature-controlled electrolyte solution, which was magnetically stirred and heated on an electric heating plate (IKA basic). The zinc sheets (99.99%) from Advent Research Materials Ltd. had geometric surface areas of 31 cm2 (NaOH electrolyte) or 18 cm2 (KOH electrolyte). The volume of the NaOH and KOH electrolytes were 1.8 L and 0.925 L, respectively. At selected times, a zinc sheet was removed from the electrolyte solution and the mass loss was measured after careful rinsing and drying. Each data point in these experiments was generated from one zinc sheet. The electrolyte solutions were open to ambient air during the experiment to allow also for oxygen reduction as cathodic corrosion reaction.
The formation of zinc dendrites and the hydrogen evolution side reaction in the zinc-air RFB can be prevented or minimized, if the SOC is monitored. Hydrogen evolution occurs to a large extent at the negative electrode during charging, if the current approaches the limiting current of zinc deposition. The latter conditions are also favourable for the growth of zinc dendrites [24,25]. Zinc dendrites can cause cell shortening. Apart from dendritic zinc, also other non-compact morphologies, like filamentous mossy and heavy spongy zinc poorly adhere to the substrate [22,26]. Zinc with these morphologies can easily detach from the electrode in the streaming electrolyte, which would result in a loss of capacity and/or obstruction of the electrolyte flow in the battery. The limiting current for zinc deposition depends on the zincate concentration and the electrolyte movement. Therefore, the operating limits of the negative electrolyte can be determined from the zincate concentration, which is correlated to the SOC. According to the Nernst Eq. (5), with the electrode potential E, the standard electrode potential E0, the universal gas constant R, the absolute temperature T, the number of transferred electrons per atom n and the Faraday constant F, the rest potential of the zinc electrode depends on the activities a of hydroxide, water, oxygen, and zincate ions. Therefore, the latter potential could be an indicator of the SOC and is reflected also in the OCV of the cell or battery. In aqueous alkaline electrolytes zinc corrodes [27,28], which results in a mixed potential (corrosion potential).
E = E0 +
+ H2O
(6)
2. Experimental
(2) (3)
⇆
CO32−
Here, we present the results of a fundamental investigation on state of charge indicators for the alkaline zinc-air redox flow battery. The electrolyte temperature in this study has been varied between room temperature and 60 °C, because many precious-metal-free air electrodes suitable for an alkaline zinc-air RFB show an enhanced voltage efficiency at elevated temperature [31].
(1)
+ H2O
0
molecules contained in a defined volume. When some of these molecules are replaced by other molecules, e.g. zincate ions, the electrolyte density can change. SOC indicators are affected by a change of the electrolyte composition. The electrolyte composition may change by ageing of the battery. In the zinc-air RFB, changes in electrolyte composition are caused by water evaporation, precipitation of zinc oxide (Eq. (6)), decomposition/loss of electrolyte additives or uptake of carbon dioxide from the air [29,30] according to Eq. (7).
(5)
Also other electrolyte properties are affected by the zincate concentration: the ionic strength and the electrical conductivity of the negative electrolyte, as well as its density and viscosity. The ability of an aqueous electrolyte to conduct electricity is based on the transport of solvated ions in the electric field between the electrodes. The mobility of the solvated ions and their concentration determine the electric conductivity. The electrolyte density is the weight of a defined electrolyte volume. That weight is the sum of the weight of the individual 77
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Fig. 3. Ionic conductivity σ at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−).
Fig. 1. Refractive index n at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−).
3. Results and discussion 3.1. NaOH electrolyte A battery charging operation has been simulated in a rotating cylinder electrode experiment. Zinc was electrodeposited on the rotating cylinder cathode, while oxygen was evolved at the platinized titanium counter electrode analogous to the processes in a zinc-air RFB (Eq. (3)). During the experiment, the refractive index (Fig. 1), density (Fig. 2), and conductivity (Fig. 3) of the electrolyte as well as the zinc rest potential (Fig. 4) were measured at regular time intervals. The experiment was performed in 0.425 L of a 7 M NaOH electrolyte containing 0.9 M dissolved ZnO and 2.4 g L−1 Lugalvan P (LP) additive at 60 °C, at a current density of 50 mA cm−2, and a rotational rate of 160 rpm and lasted about 27 h. LP is an additive in industrial alkaline baths for zinc plating and was used here to prevent the growth of dendritic zinc. The refractive index n shows only a small variation between 1.3915 and 1.3960 (Fig. 1) and there seems to be little correlation, if any, with the zincate concentration c([Zn(OH)4]2-). The coefficient of determination R2 (Pearson correlation coefficient) of the linear regression is as low as 0.36. Therefore, the refractive index is not suitable for determining the SOC in the 7 M NaOH/2.4 g L−1 LP electrolyte. As expected, the electrolyte density ρ decreases during the experiment with decreasing zincate concentration by 0.036 g cm−3, from 1.276 to 1.240 g cm−3 (Fig. 2). This behaviour was also observed by Siu and Evans [33]. The correlation between density and zincate
Fig. 4. Measured mixed potential and calculated equilibrium potential of the zinc electrode in a 7 M NaOH electrolyte with 2.4 g L−1 LP additive as a function of ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2) at 60 °C. The measured data are at the same zincate concentrations c([Zn(OH)4]2−) as those in Fig. 1, Fig. 2, and Fig. 3.
concentration in our experiment was found to be highly linear with a slope of 0.045 g cm−3 mol−1 L and a coefficient of determination R2 of 0.96. Hence, the electrolyte density appears to be useful for determining the SOC in 7 M NaOH/2.4 g L−1 LP electrolyte. Moreover, sensors for monitoring the density of alkaline solutions are commercially available. One has to keep in mind, however, that density variations can be caused also by concentration changes of species other than zincate, e.g. NaOH [33] and carbonates, Eq. (7), and by temperature variations [32]. Also aging of the battery may thus change the electrolyte composition and thereby affect the SOC-determination. The electric conductivity σ increases from 561 to 680 mS cm−1 during the charging experiment (Fig. 3). The increase in conductivity with decreasing zincate concentration can be explained by the high ionic strength, Ic, of the electrolyte, Eq. (8).
Ic =
1 ⋅ ∑ ci⋅z i2 2 i
(8)
Here, ci is the ion concentration of species i and zi is its ionic charge. At high concentrations, ions strongly interact and might even form ion pairs. This decreases both the mobility and the concentration of dissociated ionic charge carriers affecting the ionic conductivity. During charging, one zincate ion is reduced to metallic zinc and two hydroxide ions plus a water molecule are generated according to equation (3). The
Fig. 2. Density ρ at 60 °C of a 7 M NaOH electrolyte with 2.4 g L−1 LP additive at different zincate concentrations c([Zn(OH)4]2−). 78
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Figs. 4 and 13.9 mV and 14.3 mV, respectively, are nevertheless in very good agreement. Electrolyte density and conductivity correlate linearly with the zincate concentration, as does the zinc rest potential with the Nernstian term ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2) in the 7 M NaOH/ 2.4 g L−1 LP. If the total concentrations of water, OH−, and the additive are known, the zincate concentration can be calculated from the rest potential of the zinc electrode in a calibration plot. All three correlations have a coefficient of determination R2 (Pearson correlation coefficient) equal or larger than 0.96 and appear to be suitable for determining the SOC in this electrolyte. An attempt was made to determine the zincate concentration from voltammetry with ultra-microelectrodes. However, no steady-state currents for the reduction of zincate could be observed in the 7 M NaOH/0.8 M ZnO/2.4 g L−1 LP electrolyte on platinum and carbon fibre ultra-microelectrodes. Also, the consumed charge for zincate reduction was not reproducible in these experiments. Thus, amperometry with ultra-microelectrodes is unsuitable for determining the SOC.
ionic strength decreases, because it depends on the square of the ionic charges and the zincate ion is doubly charged. Moreover, due to the Grotthuss mechanism [34], hydroxide ions contribute more strongly to the electric conductivity than zincate ions do. Therefore, the conductivity increases if zincate ions are replaced by the two-fold number of hydroxide ions. A linear correlation with a coefficient of determination R2 of 0.99 was found between the electrolyte conductivity and the zincate concentration (Fig. 3). The ionic conductivity can thus be used for determining the SOC in the 7 M NaOH/2.4 g L−1 LP electrolyte. Also in this case, however, one should bear in mind that conductivity variations can be caused also by concentration changes of ions other than zincate and by temperature changes. The rest potential of the zinc electrode shows a variation of 29 mV for a change of the zincate concentration from 0.18 M to 0.83 M (Fig. 4). The measured potentials are shown together with the equilibrium potentials calculated from the Nernst Eq. (5) using concentrations instead of activities. The activity of oxygen dissolved in the electrolyte is assumed to be constant in the open cell and is therefore neglected in the calculations. The first equilibrium potential was measured after the brass electrode was completely coated with zinc at c([Zn (OH)4]2−) = 0.83 M and c(OH−) = 5.3 M. The initial OH− concentration was calculated from Eq. (6), while the varying OH− concentration during the charging experiment was calculated from Eq. (3). The water concentrations were calculated from the measured densities of the electrolyte subtracting the weight of NaOH, ZnO, and LP and taking into account the consumption of one water per ZnO according to Eq. (6). In the calculation of the equilibrium potential, the standard potential given for the half-cell reaction in Eq. (1) is used and converted to the potential scale of the Hg/HgO reference electrode (+0.1154 V vs. SHE @ 1 M KOH). The universal gas constant R is 8.314 J mol−1 K−1, T = 333 K (60 °C), n = 2 and F = 96485 A s mol−1. The difference between the measured and the calculated potentials in Fig. 4 can be explained with two effects. Firstly, zinc corrodes in this electrolyte and hence, the measured potential is the free corrosion potential of zinc for the reactions in Eqs. (3) and (9) [35–37]. A corrosion current density of 36 μA cm−2 has been determined for the 0.2 M zincate/7 M NaOH/2.4 g L−1 LP electrolyte at 60 °C from gravimetric measurements shown in Fig. S1 of the supplementary information. From the mass loss Δm over time, a linear regression was generated and used to calculate the corrosion current density jcorr from Eq. (10). Here, k is the slope of the linear regression in Fig. S1 and the molar mass M of zinc is 65.39 g mol−1. The corrosion current density in the 0.2 M zincate/8 M KOH electrolyte at 60 °C, which is discussed in section 3.2, was found to be 6.4 μA cm−2. Zn + 2H2O + 2 OH− → H2 + [Zn(OH)4]2−
jcorr =
k ⋅F ⋅n 3600 M
3.2. KOH electrolyte Electrolyte conductivity, density and the zinc rest potential have been studied in additive-free 7.5 M, 8.0 M, and 8.5 M KOH electrolytes for zincate concentrations between 0.2 and 0.6 M at temperatures of 40, 60, and 70 °C (Fig. 5, Fig. 6, Fig. 7). In the rotating cylinder electrode experiment with NaOH electrolyte discussed above, the zincate concentration was varied electrochemically, according to Eq. (3). In the experiment with KOH electrolyte, the zincate concentration was varied by dissolution of ZnO, according to Eq. (6). The comparison of Eqs. (3) and (6) shows that dissolution of ZnO leads to the same variation of the electrolyte composition as does the electrochemical dissolution of zinc. Therefore, battery discharging can be simulated by simple dissolution of ZnO in these electrolytes. The electric conductivity σ decreases from 982 to 861 mS cm−1, for an increase in zincate concentration from 0.2 M to 0.6 M in 8 M KOH electrolyte at 60 °C (Fig. 5). The decrease in conductivity with increasing zincate concentration is explained by the high ionic strength of the electrolyte and this behaviour was already discussed for the NaOH electrolyte in section 3.1. The ionic conductivity of KOH solutions under varying conditions was reviewed previously by Gilliam et al. [38]. They showed that the conductivity decreases, if the KOH concentration increases from 7 M to 8 M at 60 °C (Fig. 3 in Ref. [38]). This behaviour is in agreement with the results of this work. Fig. 5 shows that the ionic conductivity is affected by the electrolyte temperature
(9) (10)
Since the measured potential of the zinc electrode is a mixed potential, it deviates from the equilibrium potential calculated from Eq. (5). The calculation of the corrosion potential is complex. Based on the corrosion current density the corrosion potential could be calculated from the kinetic parameters of anodic and cathodic reactions (Tafel slopes and exchange current densities). However, the kinetic parameters of these reactions under the present experimental conditions are not known. Secondly, no activity data is available for the present electrolyte. Instead of ionic activities in Eq. (5) that could be calculated from approximated activity coefficients [36], concentrations have been used. The difference between the measured and the calculated curve in Fig. 4 is thus explained by i) the neglected impact of the zinc corrosion and ii) the deviation of the activity coefficients from unity. Both corrosion potential and activity coefficients depend on the electrolyte composition and thus vary with the concentrations of zincate, water, and OH−. The presence of the plating additive might affect them as well. The slopes of the measured and of the calculated potentials in
Fig. 5. Electric conductivity σ of KOH electrolytes as a function of zincate concentration c([Zn(OH)4]2−) for different KOH concentrations and temperatures. 79
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Fig. 8. Measured zinc rest potential (square points) from Fig. 7 (60 °C, 8 M KOH) and calculated equilibrium potential as function of ln(c([Zn(OH)4]2−)·c (H2O)−1·c(OH−)−2). The numbers on the data points indicate the zincate concentration.
Fig. 6. Electrolyte density ρ at 60 °C of a 8 M KOH electrolyte at different zincate concentrations c([Zn(OH)4]2−).
zincate concentration from 0.2 M to 0.6 M, the calculated KOH density increases by 0.025 g cm−3 (Fig. 6). The linear regression of the calculated density shows a coefficient of determination R2 of 0.999, which makes the electrolyte density a good indicator for measuring the SOC in 8 M KOH electrolyte. ρ = a + b w(KOH) + c w(KOH)2 + d w(ZnO) + e T
w (ZnO) = 8.143
c (ZnO) ρ
(11)
(12)
The rest potential of the zinc electrode varies by 15–20 mV for a change of zincate concentration from 0.2 M to 0.6 M (Fig. 7). The shift of the potentials with KOH concentration is in qualitative agreement with the Nernst equation, if unity activity coefficients are assumed. This is not the case for the slope in the regression of the data at 40 °C, which should be smaller than for the data at 60 °C. The slight deviations are exemplified in Fig. 8, where the measured rest potential is shown together with the calculated equilibrium potential calculated from Eq. (5), analogously to the rotating cylinder electrode experiment in section 3.1. Zinc corrodes also in this electrolyte, the corrosion current density being 6.4 μA cm−2 at 60 °C (section 3.1). Therefore, the measured rest potential is again a mixed potential (corrosion potential), as in the case of the NaOH electrolyte shown in the previous section. The kinetic parameters of the concentration and temperature dependent anodic and cathodic corrosion reactions are not known, and therefore, the corrosion potential could not be determined. Instead, the equilibrium potential was calculated for a comparison with the measured mixed potential in Fig. 8. The OH− and water concentrations in Fig. 8 were calculated from Eq. (6) (dissolution of ZnO). The equilibrium potential was calculated from the standard potential given for Eq. (1) and converted to the potential scale of the Hg/HgO reference electrode (+0.0624 V vs. SHE @ 8 M KOH). The potential of the reference electrode is shifted negatively by 53 mV, when the inner electrolyte is changed from 1 M to 8 M KOH. This potential shift is calculated from the Nernst equation and was also confirmed experimentally. Again, ionic concentrations have been used rather than ionic activities in Eq. (5) and the constant activity of dissolved oxygen omitted for the same reasons as explained in section 3.1. The calculated density from Eq. (11). was used for the calculation of the water concentration. The difference between the measured and the calculated data in Fig. 8 arises from the neglected impact of the zinc corrosion and the deviation of the ionic activity coefficients from unity. The slopes of the data sets in Fig. 8 are again similar, 11.9 mV for the linear regression of the measured potential and 16.7 mV for the calculated potential.
Fig. 7. Rest potential E of the zinc electrode as function of ln(c([Zn(OH)4]2−)·c (H2O)−1·c(OH−)−2) at different KOH concentrations and temperatures. The lines show the linear regressions. The measurements were carried out at the same zincate concentrations c([Zn(OH)4]2−) as those shown in Fig. 5.
and KOH concentration. A temperature decrease from 60 to 40 °C lowers the conductivity by 250 mS cm−1 at a zincate concentration of 0.2 M. The linear correlations between electrolyte conductivity and zincate concentration in Fig. 5 show a coefficient of determination R2 of 0.99. If the conductivity is used for determining the SOC, one has to bear in mind though that variations can be caused also by concentration changes of ions other than zincate (e.g. potassium, hydroxide, carbonates) and by temperature changes. The electrolyte densities for different zincate concentrations have been calculated with Eqs. (11) and (12), as suggested by Siu and Evans [33] and the results are shown in Fig. 6. The KOH weight concentration w(KOH) of 34.0% (8 M) was calculated according to Gilliam et al. (Equation (1) and Table 4 in Ref. [38] for a temperature of 60 °C). In Eq. (11), w(ZnO) is the weight concentration of zinc oxide, the temperature T is 60 °C, and the coefficients a to e are 1.01137, 0.008466, 0.00003071, 0.01039, −0.0005106 (Table I in Ref. [33], column “CRC data included, initial KOH concentration”). For comparison, also two measured data points are shown in Fig. 6. The measured data is shifted to smaller values by approximately 0.030 g cm−3, which might be due to a small error in either w(KOH) or w(ZnO). However, the slope of 0.077 g cm−3 mol−1 L in the measured data is close to that of 0.062 g cm−3 mol−1 L in the calculated data. The KOH electrolyte density drops with decreasing zincate concentration, which was found also for the NaOH electrolyte (section 3.1). For an increase of the 80
Journal of Power Sources 424 (2019) 76–81
C. Zelger, et al.
acknowledge funding (no. 848933) by the Climate and Energy Fund of the Austrian Federal Government through the 1st Call “Energieforschung”.
Density and ionic conductivity correlate linearly with the zincate concentration in the KOH electrolyte, as does the rest potential with the term ln(c([Zn(OH)4]2−)·c(H2O)−1·c(OH−)−2). The coefficient of determination R2 (Pearson correlation coefficient) is in all cases equal or larger than 0.99. Density and electric conductivity of the electrolyte as well as the rest potential can thus be used for determining the SOC in 8 M KOH electrolyte.
Appendix A. Supplementary data Supplementary data to this article can be found online at https:// doi.org/10.1016/j.jpowsour.2019.03.099.
4. Conclusions References Several possible state of charge indicators for the alkaline zinc-air redox flow battery have been studied during a battery charging operation simulated in a rotating cylinder electrode experiment with 7 M NaOH/ZnO/2.4 g L−1 LP electrolyte. The measurements show no significant correlation between refractive index and zincate concentration. However, density and conductivity of the NaOH electrolyte correlate linearly with the zincate concentration, while the rest potential correlates linearly with the Nernstian term ln(c([Zn(OH)4]2−)·c(H2O)−1·c (OH−)−2). Also in the case of the 8 M KOH electrolyte, density and conductivity correlate linearly with the zincate concentration, and the zinc rest potential correlates linearly with the Nernstian term ln(c([Zn (OH)4]2−)·c(H2O)−1·c(OH−)−2). In all linear regressions of density, conductivity, and rest potential for both NaOH and KOH electrolytes, the coefficient of determination R2 is either equal to 0.96 or in most cases even around 0.99. Such strong linear correlations between measured parameters and zincate concentration allow for a simple determination of the SOC. The density, conductivity, and rest potential of NaOH and KOH zincate electrolytes appear to be reliable indicators of the SOC of the zinc-air redox flow battery, if the temperature is kept constant or measured simultaneously. The formation of carbonate in the alkaline electrolytes upon absorption of CO2 might affect the three indicators. Furthermore, degradation of battery components such as cathode materials (catalysts, conductive support, binders, and current collector), materials for cell housing, and possibly membranes could result in accumulation of impurities that might also affect the electrolyte properties. These processes, however, occur slowly on a long timescale and would be an issue of battery maintenance. They depend on the specific battery installation and an assessment of their possible impact on the respective indicators thus requires thorough long-term studies. The variation of the rest potential is 29 mV in the NaOH electrolyte and 20 mV in KOH, and thus rather small in the investigated SOC range. Such small potential differences are difficult to measure. The conductivity varies by approximately 125 mS cm−1 in the NaOH and KOH electrolytes. Variations of such a magnitude are much simpler to measure than the density or rest potential. The reliability of the SOC monitoring could nevertheless be improved even further by combining two of the indicators.
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Acknowledgements The support by the European Commission (Theme 2010.7.3.1: Energy storage systems for power distribution networks, no. 256759) is gratefully acknowledged for parts of this work. C.Z., M.S., and B.G.
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