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air direct ion liquid fuel cell
Journal of Power Sources 286 (2015) 232e238
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Journal of Power Sources journal homepage: www.elsevier.com/locate/jpowsour
Dithionite/air direct ion liquid fuel cell Jens Noack*, Jens Tübke, Karsten Pinkwart Applied Electrochemistry, Fraunhofer Institute for Chemical Technology, Joseph-von-Fraunhofer-Str. 7, 76327 Pfinztal, Germany
h i g h l i g h t s A dithionite/air fuel cell has been developed and tested. The maximum power density was 2 mW cm2, the open circuit voltage 0.9 V. Low kinetics of dithionite oxidation limited the performance of the cell. ORR was influenced by dithionite solution crossover.
a r t i c l e i n f o
a b s t r a c t
Article history: Received 18 September 2014 Received in revised form 20 February 2015 Accepted 26 March 2015 Available online 31 March 2015
The feasibility of an alkaline S2O2 4 /air-fuel cell was evaluated at room temperature, using a cell with an anion exchange membrane and a platinum oxygen reduction reaction catalyst. The tests performed were open circuit voltage analysis, linear sweep voltammetry, discharge analysis and electrochemical impedance spectroscopy (EIS) with registration of anode half-cell potential. With 0.85 M Na2S2O4 in 2 M KOH, the cell achieved a maximum power density of 2 mW cm2, and the open circuit cell voltage was about 0.9 V. In a potentiostatic discharging at 0.2 V cell voltage, an energy efficiency of 12.3% was achieved at an energy density of 8.6 Wh L1. The low power density was mainly due to the low reaction kinetics of dithionite oxidation at graphite electrodes. The low energy efficiency was mainly caused by a low cathode potential, which probably resulted from mixed potential formation and the low anode kinetics. © 2015 Elsevier B.V. All rights reserved.
Keywords: Dithionite Alkaline Fuel Cell Development Characterization
1. Introduction With the increase of renewable energy sources in the energy grid, fluctuating feed and the related difference between energy supply and demand [1,2] has led to an increased need for energy storage. In addition to mechanical, electrical and thermal storage, electrochemical storage offers the potential for a secure, stable and cost effective decentralized energy supply [3]. Pb-acid batteries are a long-established storage technology up to the MWh range, but have a relatively short cycle life. Especially in recent years increased effort has been made to upscale and demonstrate various technologies, including existing technologies such as nickel batteries and sodium sulfur batteries, and alternatives such as Li-ion, redox flow batteries and the hydrogen energy chain [4]. In order to achieve energy storage with low kWh prices, high cycle rates, long
* Corresponding author. E-mail address: [email protected] (J. Noack). http://dx.doi.org/10.1016/j.jpowsour.2015.03.159 0378-7753/© 2015 Elsevier B.V. All rights reserved.
service life, low material cost and good recyclability are required. The greater the ratio of energy/power, the more important the cost of the redox couples. Most current research therefore focuses on the development of energy storage with inexpensive materials such as H, Li, Na, S, Br, Zn and Fe. In this study we searched for a combination of two low-cost redox couples which also possesses an acceptable standard potential difference, to achieve viable cell voltages. The combination 2 of S2O2 4 /SO3 couple with an oxygen reduction reaction (ORR) potentially offers very low storage cost at a potential difference of 1.53 V if this kind of fuel cell can be combined with an efficient separate electrolysis process or directly into one cell as a unitized system. In contrast to fuel cells based on organic materials, this combination provides the possibility to reverse the negative reactions by electrolysis and thus to store energy. Beside hydrogenbased fuel cells, vanadium-air fuel cells are a similar system in which the anode reaction can be directly reversed; however here the anodic energy carrier material is ions in solution [5e7]. Despite the low kinetics and complicated S2O2 -oxidation [8e10], the 4
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problem of sulfur poisoning of platinum-based ORR catalysts [11,12] and the question of the electrochemical reduction of sulfite to dithionite [13,14], we have built an alkaline dithionite fuel cell and investigated its properties. This approach appears promising as non-precious metal catalysts for the ORR are in development [15] and because the S2O2 4 -oxidation could be accelerated at suitable electrode materials. Working with a complete cell set-up we aimed to investigate the combination of the individual reactions with a Ptbased catalyst, and to determine what performance could be achieved or what problems would arise. The S2O2 4 -oxidation is a two-electron mechanism of trivalent sulfur to tetravalent as sulfite SO2 3 with an electrode potential of 0.294 V vs RHE in an alkaline medium: 2 Anode : S2 O2 4 þ 4OH #2SO3 þ 2H2 O þ 2e E ¼ 0:294 V
(1) Cathode :
O2 þ 2H2 O þ 4e /4OH
O2 þ H2 O þ 2e /HO 2 þ OH HO 2 þ H2 O þ 2e /3OH
Cell :
E ¼ þ1:227 V
E ¼ þ0:761 V
(3)
E ¼ þ1:693 V
2 2S2 O2 4 þO2 þ 4OH /4SO3 þ 2H2 O
(2)
(4) E ¼ 1:530 V (5)
S2O2 4
A 1 M solution -based fuel cell offers the potential for a theoretical maximum capacity of 53.6 Ah L1. The theoretical energy density at 1.53 V cell voltage is 82.01 Wh L1. Assuming an efficiency of 40%, the energy output would be 32.8 Wh L1.
2. Experimental 2.1. Test setup With two graphite felts (GFA5, SGL Carbon, Germany) and a catalyst-coated membrane (CCM), two graphite plates (FU 4396, Schunk Carbon GmbH, Germany), two copper current collectors and two metallic end plates, a fuel cell with a geometric active area of 40 cm2 was constructed. The CCM was prepared at the Fraunhofer ICT and consisted of 2.4 mg Pt cm2 as a catalyst layer on the cathode side and an anion exchange membrane (FAA-3, Fumatech GmbH, Germany). Teflon® (DuPont, USA) served as a binder material. The catalyst layer was applied by a spraying method. Activation of the membrane was carried out by storing the MEA for 2 h in 1 M KOH and then for 1 h in H2O. A 0.85 M Na2S2O4 (technical grade, 85%, SigmaeAldrich) served as fuel in 2 M KOH (Merck, Germany). Just as in the anode half-cell, no conventional gas diffusion layer was used in the cathode half-cell. In the negative half-cell a graphite felt (GFA5, SGL-Carbon GmbH, Germany) served as the electrode. In the positive half-cell a graphite felt provided the electrical contact between the CCM and graphite plate. As illustrated in Fig. 1 a test rig was setup with the cell. A dual-piston pump (Nano Pump, Duratec GmbH, Germany) supplied the cell with dithionite solution at the anode, and a mass flow controller supplied the cell with 100 mL min1 synthetic air at the cathode. The air was not pre-wetted. A Hg/HgO reference electrode and a glassy carbon electrode were present in the reservoir of the dithionite solution. The dithionite reservoir was continuously purged with nitrogen to prevent oxidation by atmospheric oxygen. With this set-up, the cell voltage, the anode potential and the redox potential was measured and the cathode potential was calculated. The measurements were performed with a potentiostat and a
Fig. 1. Schematic representation of the experimental setup used for the operation and measurements of a dithionite/air fuel cell (4N e redox potential of dithionite, 4A e anode potential, 4Z e cell voltage).
frequency response analyzer (FRA þ Pstat 20874A 2055A, Solartron Analytics, USA). All measurements were performed at room temperature. Unless otherwise stated, all potentials were calculated against a reversible hydrogen electrode (RHE). 2.2. Linear sweep voltammetry To prevent redox potential differences, linear sweep voltammetry was carried out using an arrangement of two reservoirs, in which fresh electrolyte was continually passed through the cell into a waste container. The scan rate was 20 mV s1 and the measurement was carried out from the open circuit voltage of the cell to 0 V. In addition to the cell voltage, the redox potential of the solution and the anode potential were measured. The cathode potential was calculated from the cell voltage and the anode potential. 2.3. Electrochemical impedance spectroscopy (EIS) EIS measurements were made with a potentiostatic AC amplitude of 10 mV at open circuit voltage in a frequency range of 100 kHz e 0.01 Hz or to 1 Hz as a single sinus frequency sweep. The measurements were checked regarding linearity by using the voltageetime curve to ensure that the measured current varied horizontally around the mean during the experiment, and did not drift in the positive or negative direction. Analysis of the spectra was performed with the software ZView (Scribner Inc., USA) or with Modulab (Solartron, USA) using a modified randles equivalent circuit or as circle fit. Inductive effects were not considered in the fitting, so correspondingly lower frequencies than the measured frequencies were evaluated. With the exception of the measurements during the discharge test, the EIS measurements were carried out without forced convection of the dithionite solution. The anode impedance spectrum was recorded by the usage of a potentiostat auxillary channel. 2.4. Discharge analysis The discharge was carried out as shown in Fig. 1 with one reservoir for the dithionite solution. The solution was pumped through the cell in circulation at a rate of 5 mL min1. The volume of dithionite solution was 60 mL. To be able to obtain impedance spectra as a function of the potentials, the discharge was carried out in multiple steps that were performed several times in the cycle:
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measurement of all potentials for 5 min, EIS measurement of 100 kHz e 1 Hz, measurement of all potentials for 5 min, potentiostatic discharging for 30 min at 0.2 V, measurement of all potentials for 5 min. The steps were repeated 100 times. 3. Results & discussion 3.1. Open circuit analyses Fig. 2 shows the cell voltage, the half-cell potential and the redox potential of the dithionite solution. The redox potential was about 0.106 V, and with 188 mV it was well below the standard 2 potential for S2O2 4 /SO3 . Although Na2S2O4 has been used directly after purchase, it can be presumed that a high proportion of S2O2 4 was already oxidized by contact with atmospheric oxygen, and the content was slightly lower than 0.85 M. The calculated cathode potential was 0.796 V, and with 431 mV it was well below the standard electrode potential of the four-electron reaction. The low cathodic potential is indicative of less than optimal conditions on the three-phase boundary layer. Since the air was not humidified it could have led to a low potential if the cathode were too dry. However, the anode half-cell was completely flushed with liquid, so the expectation was rather that the cathode would be flooded. A continuous passage of dithionite solution would lead to a reaction with oxygen, significantly reducing the percentage content of the solution and thus decreasing the cathode potential. Furthermore, the negative potential of dithionite would make a further contribution, establishing a mixed potential at the cathode. In addition, sulfur compounds may block the active catalyst centers. The cell voltage was 0.881 V, which was 640 mV below the difference of the standard potentials and was mainly caused by the low potential of the ORR. All potentials were constant during the first 10 min, which was close to an equilibrium state where no relevant side reactions occurred. 3.2. Linear sweep voltammetry Fig. 3 shows linear sweep voltammograms at dithionite solution flow rates of 23 mL min1 e 5 mL min1. The short-circuit current density was at 0 V cell voltage between 14.7 and 17.6 mA cm2, and the open circuit voltage between 0.589 and 0.819 V, with increasing values as the flow rate increased. The differences in the cell voltages at small currents were not as expected, but may be related to the
Fig. 2. Terminal voltages and half-cell potential of a cell with 0.85 M Na2S2O4, 2 M KOH.
Fig. 3. Current/voltage curves of a linear sweep voltammetry with different anolyte flow rates at a fuel cell with 0.85 M Na2S2O4.
specific construction of the cell. At low flow rates of the anolyte, and at a constant volume flow of air, the pressure in the cathode half-cell and hence the electrical contact between the CCM and graphite felt could have decreased. The substantial depolarization at low current levels is likely to be explained by a relatively high overpotential for the activation of the total reaction of the cell. To achieve a cell voltage of about 0.3 V more than 500 mV overpotential was necessary at a flow rate of 5 mL min1. At lower flow rates the voltammograms shifted towards the values of 23 mL min1. The activation overvoltage declined in relation to the open circuit voltage with decreasing flow rates, but increased in relation to the thermodynamic equilibrium state at 1.53 V, which is clearly shown by the parallel shift of the current/voltage curves. The shape of the curves showed no limitations of current densities by mass transport effects, although the flow rates were already very low. The lambda values (ratio of theoretical volume flow/adjusted volume) for 23 mL min1 at 14.6 mA cm2 were only 0.1, while lambda of 5 mL min1 with 17.6 mA cm2 was 22.8. Probably the volume of dithionite solution (about 10 mL at 85% felt porosity) was enough to not get into an area of mass transport control, but that would be expected with longer duration of discharging or higher current flows even at low flow rates. The relatively fast diffusion of dithionite provided the possibility to minimize mass transport limitations for all current densities. Although the trend with increasing flow rates goes clearly in the direction of higher currents and power densities, it is not expected that increased flow rates would result in a much higher performance of the cell. Considering the context of increasing open circuit voltage and decreasing activation overvoltage up to a lambda of the anolyte of 22, the reasons for these changes must be sought elsewhere. Since the experiments were performed in ascending order, starting with a low flow rate, a 2 S2O2 4 /SO3 e equilibrium change at the anode could have led to a lower potential of the subsequent measurement, in that the anode potential was not fully aligned to the redox potential, and a relevant quantity of SO2 3 was still present, especially in the area of the double layer. Between the measurements there was a delay of 60 s, which may have led, at very low flow rates, to a low transport of charge carriers and low mixing at the very porous anode. 1.8 mAh of charge were transferred during the experiment with 23 mL min1, which corresponded to a volume of 33.6 mL dithionite solution. The lambda value of 0.7 of the anolyte was in a region of undersupply. However, at 100 mL min1, the transferred charge was also only 1.84 mAh but the lambda value was around 3. At a flow rate of 5 mL min1, the lambda value of the anolyte was 22. Despite the
J. Noack et al. / Journal of Power Sources 286 (2015) 232e238
relatively high lambda values there could be a significant change in the region of the double layer, due to the low kinetics of the anode reaction and because of the foregoing measurement with a lower flow rate, at which the equilibrium between the redox potential and the anode potential had not yet been achieved. For a further analysis during the linear sweep voltammetry, potential and voltage characteristics of the cell and the half-cells were plotted in Fig. 4. In the cell with 5 mL min1 the anode potential started at 0.127 V and reached 0.019 V at 23 mL min1. The difference of 146 mV is significant and was not due to obviously different redox potentials of the dithionite solution. For 5 mL min1, the redox potential was 0.218 V, and at 23 mL min1 it was 0.225 V, so the difference was only 7 mV. The first measured current density was 0.21 mA cm2 at 5 mL min1, and 0.28 mA cm2 at 23 mL min1, so this does not explain the difference either, especially since the difference during the test remained approximately constant. The differences between redox potential and the first measured value of the anode potential were 91 mV at 5 mL min1 and 244 mV at 23 mL min1. The cell at 23 mL min1 had about 2.5 times the activation overpotential for the anodic reaction compared to the cell at 5 mL min1. For both flow rates, the cathode potential showed negligible differences at the beginning of the measurements. Thus the difference in cell voltage at the beginning of the measurements was only caused by the different activation overpotentials of the anode reactions. As well as in the analysis of the OCV measurements no further explanation could be found here, as the equilibria were different at the anode from the solution and had not adjusted during the 60 s between measurements. By the use of higher flow rates it could be possible to reach the equilibrium more quickly. The cathode potential calculated at the beginning of the measurements had a value of 0.693 V, which was well below the standard potential of the 4-electron mechanism with 1.227 V, and slightly below the peroxide formation at 0.761 V. It can be assumed that the ORR proceeded with different limitations because the activation overpotential compared to the 4-electron mechanism was 534 mV and possibly peroxide was formed. The difference between the anode potential at the start of measurement and the final value was 609 mV at 5 mL/min, and 486 mV at 23 mL min1. The corresponding differences for the cathode potential were only about 200 mV. From these values it can be concluded that the rate of the anodic reaction was about 2.5e3 times lower than that of the cathode reaction. The potential of the
235
ORR, however, had a substantially higher deviation from the equilibrium potential than the anode reaction. The cathode potential changed after about 15 s measurement time to significantly lower values. The cathode potential of the cell at 23 mL min1 depolarized more intensely than at 5 mL min1. Probably the reaction mechanism of the anode reaction had changed during the measurement. As the anode potential of the cell at 23 mL min1 was generally lower than that of the cell with 5 mL min1, the cathode depolarized slightly more because a different reaction at higher anode potential was more pronounced than in the case of the cell with 5 mL min1. The kinetics of this by-reaction was greater than the S2O2 4 -oxidation, leading to a greater depolarization of the potential of the ORR. Since the curve bending of the cathode potential starts from an anode potential of about 0.1 V and a certain overpotential for the reaction is assumed, in particular the formation of 2 SO2 4 /SO3 - must be considered as a by-reaction during the linear sweep voltammetry under the applied potential. Fig. 5 shows the IR-corrected cell voltages and the half-cell potentials as Tafel-plots. The equilibrium voltages of the cell (Fig. 5a) differed significantly from each other for the reasons discussed above. At lower flow rates two different Tafel-slopes appear; at higher flow rates there is only one. The pattern of the voltages seems to indicate the start of transport limitations. In Fig. 5b, the cathode potential shows no significant change with the change of the flow rate of dithionite solution, so it can be concluded that the flow rate has no effect on the cathodic reaction mechanism. At high current densities the Tafel-slope was 215 mV/dec (a*n ¼ 0.27), which indicates an irreversible electrochemical reaction in this system. However, the linear range was not fully formed, so no conclusions could be made concerning eventual transport limitations or an exact value for a*n. When the curves of the cathode potential are compared with those of the anode potential and the cell voltages, it can be seen that the cathode reaction was responsible for the bending of the cell voltage at higher current densities. At 5 mL/min the anode Tafelslope was 318 mV/dec (a*n ¼ 0.19). It suggests that this reaction was even slower than the ORR and which limited the cell performance significantly. A significant mass transfer limitation could not be detected. The course of the anode potentials at low flow rates and current densities showed complications in the anode reaction, which are believed to have been caused by the above-mentioned problems i.e. low fluid supply or pressure at low dithionite lambda values. 3.3. Discharge analysis
Fig. 4. Cell and half-cell potentials of a cell with 0.85 M Na2S2O4 during linear sweep voltammetry with anolyte flow rates of 23 mL min1 and 5 mL min1.
Fig. 6 shows the potential profiles and the capacity during the discharge of a dithionite fuel cell. Some spikes are artifacts of the measurements and were not removed by smoothing methods. As regards the cell voltage, a depolarization of 642 mV can be observed from the beginning of the measurement at OCV 0.842 to the discharge voltage of 0.200 V. The cathode potential at OCV was 0.774 V and quickly depolarized to 0.693 V. The value of 0.693 V corresponded exactly to the value of the cathode potential at the beginning of the investigations using linear sweep voltammetry. Both values were significantly different to the theoretical value of standard potential of the 4-electron mechanism of the ORR. The activation overvoltage of the ORR was 81 mV. The anode potential changed from 0.07 V to 0.491 V by an amount of 561 mV. The kinetics of the anode reaction was approximately seven times slower than that of the ORR. As the experiment progressed the redox potential of the solution changed from 0.107 V to a value of 0.749 V at the end of the measurement. After a test period of 30 h a potential jump of the
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Fig. 5. Tafel plots of IR corrected cell voltage (a) and half-cell potentials (b) of a fuel cell with 0.85 M Na2S2O4 with variation of anolyte flow rate.
half-cells can be recognized, which is normally indicative of a change in the reaction mechanism. Up to this time 99% of the charge carriers were transferred with 2.57 Ah capacity. The following change of potentials hasn't resulted in a relevant charge transfer process. The test can therefore be divided into two parts. At first dithionite was oxidized until a time of about 30 h, and afterward significant changes in the cell occurred. In the first part, the OCV was on average about 0.7 V, in the second part at 0.3 V. In the first part the voltage progression was unusual. After a period of about 20 h, the terminal voltage rose again slightly and then dropped
sharply in the second part. Looking at the half-cell potentials one can see the increase of cathode OCP with time, which moved together with the trend of simultaneously increasing anode potential. In the second part the cathode OCP finally reached a plateau, in which further depolarizations were hardly discernible. Probably the three-phase interface of the cathode and in particular the surface of the Pt catalyst changed. It can be assumed that even in a short time the dithionite solution can pass through the membrane to the catalyst and can therefore alter the kinetics of the ORR. With the change in the composition of the dithionite solution the activity of the catalyst could have been changed at the same time, as evidenced here in an increasing cathode potential. However, the anode and cathode reactions are strongly pH-dependent (Equations (1) and (2)). For the complete reaction of 0.85 M S2O2 4 according to Equation (5), the redox potential should increase by 100 mV due to the pH decrease. The change in the cathode potential was about 150 mV, which was in the range of the possible change because large potential differences can be expected even at moderate pH values. The increase in the cathode potential was thus mainly a cause of pH value change, but shows at the same time that there must be a clear balance between the composition of the cathodic three-phase boundary layer and the anolyte solution, because a pH value increase can only be achieved by dithionite oxidation, which normally takes place in the anode half-cell. This relationship shows a continuous passage of liquid through the membrane towards the cathode, leading to the assumption that this was rather too wet to function more efficiently Table 1. 60 mL of 0.85 M Na2S2O4 solution were used during the experiment which correspond to a capacity of 2.73 Ah in a twoelectron transition. The measured value was 2.57 Ah and hence a charge carrier efficiency of 94.1% could be achieved. The loss of 5.9% charge could have had three main reasons. The first possibility is that the dithionite solution was already oxidized, since the preparation of the solution was not carried out under inert gas atmosphere. Oxidation of the solution could have occurred during the measurement due to diffused oxygen and the diffusion of dithionite
Table 1 Theoretical and measured voltages, capacities and energy densities of a cell with 0.85 M Na2S2O4.
Fig. 6. Potentials of a cell with 0.85 M Na2S2O4 at potentiostatic discharge at 0.2 V.
Theoretical Measured Efficiency [%]
Cell voltage [V]
Capacity [Ah L1]
Energy density [Wh L1]
1.53 0.2 13.1
45.5 42.8 94.1
69.6 8.6 12.4
J. Noack et al. / Journal of Power Sources 286 (2015) 232e238
through the membrane. The redox potential of the solution at the beginning of the measurement was a bit higher than the theoretical amount of 0.294 V. This suggests an already partial oxidation of dithionite prior to measurement. No evidence was found of the diffusion of atmospheric oxygen into the solution through the cell or in the vessel or tubing, but this was not ruled out. On the other hand, the low cathode potential provides evidence of the passage of dithionite solution. Considering the entire loss of capacity of 160 mAh as a liquid passage with unchanged concentration through the membrane, 3 mL of anolyte entered the cathode halfcell during about 45 h continuous measurement. Probably all three reasons, to an unknown extent, led to the charge carrier loss. The operating voltage of the cell at 0.2 V represented a voltage efficiency of 13.1%. The low value was caused firstly by the low reaction kinetics of the oxidation of S2O2 4 and secondly by the low equilibrium potential of the ORR. The total amount of converted energy was 0.51 Wh, corresponding to an energy density of 8.6 Wh L1. The energy efficiency of the cell was 12.4%. 3.4. Electrochemical impedance spectroscopy In Fig. 7 the Nyquist and Bode plots of a cell with fresh 1 M Na2S2O4 are shown with a cell voltage of 0.871 V in a state without flowing anolyte and air. The EIS-spectra were fitted with a simplified equivalent circuit presented in Fig. 8. The ohmic resistance R1 of the cell was about 0.1 U, and the charge transfer resistance R2 was 14.7 U. Both values were relatively high, although R2 clearly dominated. The semi-circle of the time constant in the Nyquist plot had a significantly depressed semicircle with a capacitor exponent of 0.4 and a capacity of 0.77 F. The Nyquist plot of the anode impedance had an almost similar behavior and values as the cell impedance. The low kinetics of the anode reaction were therefore dominant throughout the cell. The deviation from the ideal capacitor behavior can be assigned by an increased response on adsorbates on the electrode surface. The Bode and Nyquist plot had a constant phase shift in the frequency range of 1e10 kHz, which had the character of a transmission line and is a characteristic of a highly developed porous surface. In the Bode plot the curves of the phase shift from the cell and anode showed negative phase shifts at high frequencies, while the impedance increased. For inductive effects, however, a positive phase shift is expected. These phenomena were not further evaluated. Probably overlays of disorders and inductive effects were the cause of these results. The Nyquist or Bode plots illustrated in Fig. 9 show the impedance spectra of a cell at different OCV after discharge. The measurements are taken from the discharge experiment illustrated
237
Fig. 8. Simplified equivalent circuit diagram for the modeling of EIS measurements of a cell with 0.85 M Na2S2O4 (R1 e ohmic resistance, R2 e charge transfer resistance of the cell reaction, CPE1 e double layer capacitance).
in Fig. 6, in which the voltage values correspond to the 1st, 50th, 70th or 100th discharge step. Due to the relatively complex results and in particular the lack of low frequency range (<1 Hz) no fitting or modeling has been performed. Nevertheless, some features and trends were qualitatively evaluated. In the low frequency range of 1e10 Hz, a time constant was dominant, which could be attributed mainly to the reaction of dithionite. With increasing discharge time, the charge transfer resistance of the reaction increased significantly. In the high-frequency part at about 10 kHz inductive effects in particular were detected during the first measurement. However, with increasing discharge time in the region of 1e100 kHz, a further time constant appeared, which has been superimposed with inductive effects, and thus did not form a typical semi-circle in the Nyquist plot. The high-frequency time constant was observed only in the impedance measurements of the cell, but not in those of the anode. Because of the high resonance frequencies, this time constant can be attributed to the electrolyte conductivity which decreased significantly during discharging. The decrease occurred in parallel with decreasing pH, which resulted from the nearly complete consumption of OH e ions and thus decreased the electrolyte conductivity during discharge. 4. Conclusions The results obtained show that the construction and operation of a dithionite/air fuel cell is feasible in principle. The power densities achieved were 1e3 decades below practical values, so it can be expected that such a non-optimized converter would lead to very high costs. The values are comparable with other systems in development and could certainly be further increased by optimizing the reaction kinetics, cell and plant management. By means of appropriate and cost-effective anode catalysts the dithionite oxidation could be accelerated and increased, while further studies on the behavior of the cathodes could enable their kinetics and equilibrium potential to be increased, improving power density and energy efficiency. The values obtained here were obtained at room
Fig. 7. Nyquist and Bode plots of a cell with 0.85 M Na2S2O4 at 0.871 V open circuit voltage.
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Fig. 9. Nyquist Bode plots of a cell with 1 M Na2S2O4 in 2 M KOH at different terminal voltages by discharging.
temperature. By increasing the temperature, the individual reaction rates would be likely to increase considerably; however, this would be associated with an increase in other technical or chemical complications such as cathode management and anolyte stability. The anolyte stability is of crucial importance for the cost, in order to achieve high numbers of cycles with high charge carrier efficiencies, and to keep the cost of reconditioning and anolytebalancing as low as possible.
[6]
[7]
[8]
Acknowledgments [9]
The authors would like to thank Dr. Carsten Cremers, Dr. Peter Fischer and Thomas Berger for the useful and interesting discussions on the subject, Florina Jung for the preparation of the CCM and Carolyn Fisher for translation and spell checking.
[10]
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redox fuel cell, J. Appl. Electrochem. 41 (2011) 1223e1232, http:// dx.doi.org/10.1007/s10800-011-0342-8. S.S. Hosseiny, M. Saakes, M. Wessling, A polyelectrolyte membrane-based vanadium/air redox flow battery, Electrochem. Commun. 13 (8) (2011) 751e754, http://dx.doi.org/10.1016/j.elecom.2010.11.025. Noack Jens, Cremers Carsten, Bayer Domnik, Tübke Jens, Pinkwart Karsten, Development and characterization of a 280 cm2 vanadium/oxygen fuel cell, J. Power Sources 253 (0) (2014) 397e403, http://dx.doi.org/10.1016/ j.jpowsour.2013.12.070. E. Gasana, P. Westbroek, K. De Wael, E. Temmerman, K. De Clerck, P. Kiekens, Kinetics and mechanism of the oxidation of sodium dithionite at a platinum electrode in alkaline solution, J. Electroanal. Chem. 553 (0) (2003) 35e42, http://dx.doi.org/10.1016/S0022-0728(03)00255-9. P. Westbroek, J. de Strycker, K. van Uytfanghe, E. Temmerman, Electrochemical behaviour of sodium dithionite and sulfite at a gold electrode in alkaline solution, J. Electroanal. Chem. 516 (1e2) (2001) 83e88, http:// dx.doi.org/10.1016/S0022-0728(01)00656-8. E. Gasana, P. Westbroek, E. Temmerman, H.P. Thun, F. Twagiramungu, Influence of changes of platinum electrode surface condition on the kinetics of the oxidation of sodium dithionite and sulfite in alkaline solution, Electrochem. Commun. 2 (10) (2000) 727e732, http://dx.doi.org/10.1016/S1388-2481(00) 00109-0. Encyclopedia of Electrochemical Power Sources, vol. 2, 2009, p. 941. Handbook of Fuel Cells, vol. 2, 2003, p. 555. L. Scott Leonard, Youzhen Ding, M. Stalder Sloane, A. Kohl Paul, Jack Winnick, A. Bottomley Lawrence, Electrosynthesis of sodium hydrosulfite: I. Development of an online process control monitor, J. Electrochem. Soc. 145 (12) (1998) 4052e4056, http://dx.doi.org/10.1149/1.1838913. Youzhen Ding, L. Scott Lenonard, M. Stalder Sloane, A. Kohl Paul, Jack Winnick, A. Bottomley Lawrence, Electrosynthesis of sodium hydrosulfite: II. The effect of Cathode Material, J. Electrochem. Soc. 145 (12) (1998) 4057e4061, http:// dx.doi.org/10.1149/1.1838914. Chen Zhongwei, Higgins Drew, Yu Aiping, Zhang Lei, Jiujun Zhang, A review on non-precious metal electrocatalysts for PEM fuel cells, Energy Environ. Sci. 4 (9) (2011) 3167e3192, http://dx.doi.org/10.1039/C0EE00558D.
Contents lists available at ScienceDirect
Journal of Power Sources journal homepage: www.elsevier.com/locate/jpowsour
Dithionite/air direct ion liquid fuel cell Jens Noack*, Jens Tübke, Karsten Pinkwart Applied Electrochemistry, Fraunhofer Institute for Chemical Technology, Joseph-von-Fraunhofer-Str. 7, 76327 Pfinztal, Germany
h i g h l i g h t s A dithionite/air fuel cell has been developed and tested. The maximum power density was 2 mW cm2, the open circuit voltage 0.9 V. Low kinetics of dithionite oxidation limited the performance of the cell. ORR was influenced by dithionite solution crossover.
a r t i c l e i n f o
a b s t r a c t
Article history: Received 18 September 2014 Received in revised form 20 February 2015 Accepted 26 March 2015 Available online 31 March 2015
The feasibility of an alkaline S2O2 4 /air-fuel cell was evaluated at room temperature, using a cell with an anion exchange membrane and a platinum oxygen reduction reaction catalyst. The tests performed were open circuit voltage analysis, linear sweep voltammetry, discharge analysis and electrochemical impedance spectroscopy (EIS) with registration of anode half-cell potential. With 0.85 M Na2S2O4 in 2 M KOH, the cell achieved a maximum power density of 2 mW cm2, and the open circuit cell voltage was about 0.9 V. In a potentiostatic discharging at 0.2 V cell voltage, an energy efficiency of 12.3% was achieved at an energy density of 8.6 Wh L1. The low power density was mainly due to the low reaction kinetics of dithionite oxidation at graphite electrodes. The low energy efficiency was mainly caused by a low cathode potential, which probably resulted from mixed potential formation and the low anode kinetics. © 2015 Elsevier B.V. All rights reserved.
Keywords: Dithionite Alkaline Fuel Cell Development Characterization
1. Introduction With the increase of renewable energy sources in the energy grid, fluctuating feed and the related difference between energy supply and demand [1,2] has led to an increased need for energy storage. In addition to mechanical, electrical and thermal storage, electrochemical storage offers the potential for a secure, stable and cost effective decentralized energy supply [3]. Pb-acid batteries are a long-established storage technology up to the MWh range, but have a relatively short cycle life. Especially in recent years increased effort has been made to upscale and demonstrate various technologies, including existing technologies such as nickel batteries and sodium sulfur batteries, and alternatives such as Li-ion, redox flow batteries and the hydrogen energy chain [4]. In order to achieve energy storage with low kWh prices, high cycle rates, long
* Corresponding author. E-mail address: [email protected] (J. Noack). http://dx.doi.org/10.1016/j.jpowsour.2015.03.159 0378-7753/© 2015 Elsevier B.V. All rights reserved.
service life, low material cost and good recyclability are required. The greater the ratio of energy/power, the more important the cost of the redox couples. Most current research therefore focuses on the development of energy storage with inexpensive materials such as H, Li, Na, S, Br, Zn and Fe. In this study we searched for a combination of two low-cost redox couples which also possesses an acceptable standard potential difference, to achieve viable cell voltages. The combination 2 of S2O2 4 /SO3 couple with an oxygen reduction reaction (ORR) potentially offers very low storage cost at a potential difference of 1.53 V if this kind of fuel cell can be combined with an efficient separate electrolysis process or directly into one cell as a unitized system. In contrast to fuel cells based on organic materials, this combination provides the possibility to reverse the negative reactions by electrolysis and thus to store energy. Beside hydrogenbased fuel cells, vanadium-air fuel cells are a similar system in which the anode reaction can be directly reversed; however here the anodic energy carrier material is ions in solution [5e7]. Despite the low kinetics and complicated S2O2 -oxidation [8e10], the 4
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problem of sulfur poisoning of platinum-based ORR catalysts [11,12] and the question of the electrochemical reduction of sulfite to dithionite [13,14], we have built an alkaline dithionite fuel cell and investigated its properties. This approach appears promising as non-precious metal catalysts for the ORR are in development [15] and because the S2O2 4 -oxidation could be accelerated at suitable electrode materials. Working with a complete cell set-up we aimed to investigate the combination of the individual reactions with a Ptbased catalyst, and to determine what performance could be achieved or what problems would arise. The S2O2 4 -oxidation is a two-electron mechanism of trivalent sulfur to tetravalent as sulfite SO2 3 with an electrode potential of 0.294 V vs RHE in an alkaline medium: 2 Anode : S2 O2 4 þ 4OH #2SO3 þ 2H2 O þ 2e E ¼ 0:294 V
(1) Cathode :
O2 þ 2H2 O þ 4e /4OH
O2 þ H2 O þ 2e /HO 2 þ OH HO 2 þ H2 O þ 2e /3OH
Cell :
E ¼ þ1:227 V
E ¼ þ0:761 V
(3)
E ¼ þ1:693 V
2 2S2 O2 4 þO2 þ 4OH /4SO3 þ 2H2 O
(2)
(4) E ¼ 1:530 V (5)
S2O2 4
A 1 M solution -based fuel cell offers the potential for a theoretical maximum capacity of 53.6 Ah L1. The theoretical energy density at 1.53 V cell voltage is 82.01 Wh L1. Assuming an efficiency of 40%, the energy output would be 32.8 Wh L1.
2. Experimental 2.1. Test setup With two graphite felts (GFA5, SGL Carbon, Germany) and a catalyst-coated membrane (CCM), two graphite plates (FU 4396, Schunk Carbon GmbH, Germany), two copper current collectors and two metallic end plates, a fuel cell with a geometric active area of 40 cm2 was constructed. The CCM was prepared at the Fraunhofer ICT and consisted of 2.4 mg Pt cm2 as a catalyst layer on the cathode side and an anion exchange membrane (FAA-3, Fumatech GmbH, Germany). Teflon® (DuPont, USA) served as a binder material. The catalyst layer was applied by a spraying method. Activation of the membrane was carried out by storing the MEA for 2 h in 1 M KOH and then for 1 h in H2O. A 0.85 M Na2S2O4 (technical grade, 85%, SigmaeAldrich) served as fuel in 2 M KOH (Merck, Germany). Just as in the anode half-cell, no conventional gas diffusion layer was used in the cathode half-cell. In the negative half-cell a graphite felt (GFA5, SGL-Carbon GmbH, Germany) served as the electrode. In the positive half-cell a graphite felt provided the electrical contact between the CCM and graphite plate. As illustrated in Fig. 1 a test rig was setup with the cell. A dual-piston pump (Nano Pump, Duratec GmbH, Germany) supplied the cell with dithionite solution at the anode, and a mass flow controller supplied the cell with 100 mL min1 synthetic air at the cathode. The air was not pre-wetted. A Hg/HgO reference electrode and a glassy carbon electrode were present in the reservoir of the dithionite solution. The dithionite reservoir was continuously purged with nitrogen to prevent oxidation by atmospheric oxygen. With this set-up, the cell voltage, the anode potential and the redox potential was measured and the cathode potential was calculated. The measurements were performed with a potentiostat and a
Fig. 1. Schematic representation of the experimental setup used for the operation and measurements of a dithionite/air fuel cell (4N e redox potential of dithionite, 4A e anode potential, 4Z e cell voltage).
frequency response analyzer (FRA þ Pstat 20874A 2055A, Solartron Analytics, USA). All measurements were performed at room temperature. Unless otherwise stated, all potentials were calculated against a reversible hydrogen electrode (RHE). 2.2. Linear sweep voltammetry To prevent redox potential differences, linear sweep voltammetry was carried out using an arrangement of two reservoirs, in which fresh electrolyte was continually passed through the cell into a waste container. The scan rate was 20 mV s1 and the measurement was carried out from the open circuit voltage of the cell to 0 V. In addition to the cell voltage, the redox potential of the solution and the anode potential were measured. The cathode potential was calculated from the cell voltage and the anode potential. 2.3. Electrochemical impedance spectroscopy (EIS) EIS measurements were made with a potentiostatic AC amplitude of 10 mV at open circuit voltage in a frequency range of 100 kHz e 0.01 Hz or to 1 Hz as a single sinus frequency sweep. The measurements were checked regarding linearity by using the voltageetime curve to ensure that the measured current varied horizontally around the mean during the experiment, and did not drift in the positive or negative direction. Analysis of the spectra was performed with the software ZView (Scribner Inc., USA) or with Modulab (Solartron, USA) using a modified randles equivalent circuit or as circle fit. Inductive effects were not considered in the fitting, so correspondingly lower frequencies than the measured frequencies were evaluated. With the exception of the measurements during the discharge test, the EIS measurements were carried out without forced convection of the dithionite solution. The anode impedance spectrum was recorded by the usage of a potentiostat auxillary channel. 2.4. Discharge analysis The discharge was carried out as shown in Fig. 1 with one reservoir for the dithionite solution. The solution was pumped through the cell in circulation at a rate of 5 mL min1. The volume of dithionite solution was 60 mL. To be able to obtain impedance spectra as a function of the potentials, the discharge was carried out in multiple steps that were performed several times in the cycle:
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measurement of all potentials for 5 min, EIS measurement of 100 kHz e 1 Hz, measurement of all potentials for 5 min, potentiostatic discharging for 30 min at 0.2 V, measurement of all potentials for 5 min. The steps were repeated 100 times. 3. Results & discussion 3.1. Open circuit analyses Fig. 2 shows the cell voltage, the half-cell potential and the redox potential of the dithionite solution. The redox potential was about 0.106 V, and with 188 mV it was well below the standard 2 potential for S2O2 4 /SO3 . Although Na2S2O4 has been used directly after purchase, it can be presumed that a high proportion of S2O2 4 was already oxidized by contact with atmospheric oxygen, and the content was slightly lower than 0.85 M. The calculated cathode potential was 0.796 V, and with 431 mV it was well below the standard electrode potential of the four-electron reaction. The low cathodic potential is indicative of less than optimal conditions on the three-phase boundary layer. Since the air was not humidified it could have led to a low potential if the cathode were too dry. However, the anode half-cell was completely flushed with liquid, so the expectation was rather that the cathode would be flooded. A continuous passage of dithionite solution would lead to a reaction with oxygen, significantly reducing the percentage content of the solution and thus decreasing the cathode potential. Furthermore, the negative potential of dithionite would make a further contribution, establishing a mixed potential at the cathode. In addition, sulfur compounds may block the active catalyst centers. The cell voltage was 0.881 V, which was 640 mV below the difference of the standard potentials and was mainly caused by the low potential of the ORR. All potentials were constant during the first 10 min, which was close to an equilibrium state where no relevant side reactions occurred. 3.2. Linear sweep voltammetry Fig. 3 shows linear sweep voltammograms at dithionite solution flow rates of 23 mL min1 e 5 mL min1. The short-circuit current density was at 0 V cell voltage between 14.7 and 17.6 mA cm2, and the open circuit voltage between 0.589 and 0.819 V, with increasing values as the flow rate increased. The differences in the cell voltages at small currents were not as expected, but may be related to the
Fig. 2. Terminal voltages and half-cell potential of a cell with 0.85 M Na2S2O4, 2 M KOH.
Fig. 3. Current/voltage curves of a linear sweep voltammetry with different anolyte flow rates at a fuel cell with 0.85 M Na2S2O4.
specific construction of the cell. At low flow rates of the anolyte, and at a constant volume flow of air, the pressure in the cathode half-cell and hence the electrical contact between the CCM and graphite felt could have decreased. The substantial depolarization at low current levels is likely to be explained by a relatively high overpotential for the activation of the total reaction of the cell. To achieve a cell voltage of about 0.3 V more than 500 mV overpotential was necessary at a flow rate of 5 mL min1. At lower flow rates the voltammograms shifted towards the values of 23 mL min1. The activation overvoltage declined in relation to the open circuit voltage with decreasing flow rates, but increased in relation to the thermodynamic equilibrium state at 1.53 V, which is clearly shown by the parallel shift of the current/voltage curves. The shape of the curves showed no limitations of current densities by mass transport effects, although the flow rates were already very low. The lambda values (ratio of theoretical volume flow/adjusted volume) for 23 mL min1 at 14.6 mA cm2 were only 0.1, while lambda of 5 mL min1 with 17.6 mA cm2 was 22.8. Probably the volume of dithionite solution (about 10 mL at 85% felt porosity) was enough to not get into an area of mass transport control, but that would be expected with longer duration of discharging or higher current flows even at low flow rates. The relatively fast diffusion of dithionite provided the possibility to minimize mass transport limitations for all current densities. Although the trend with increasing flow rates goes clearly in the direction of higher currents and power densities, it is not expected that increased flow rates would result in a much higher performance of the cell. Considering the context of increasing open circuit voltage and decreasing activation overvoltage up to a lambda of the anolyte of 22, the reasons for these changes must be sought elsewhere. Since the experiments were performed in ascending order, starting with a low flow rate, a 2 S2O2 4 /SO3 e equilibrium change at the anode could have led to a lower potential of the subsequent measurement, in that the anode potential was not fully aligned to the redox potential, and a relevant quantity of SO2 3 was still present, especially in the area of the double layer. Between the measurements there was a delay of 60 s, which may have led, at very low flow rates, to a low transport of charge carriers and low mixing at the very porous anode. 1.8 mAh of charge were transferred during the experiment with 23 mL min1, which corresponded to a volume of 33.6 mL dithionite solution. The lambda value of 0.7 of the anolyte was in a region of undersupply. However, at 100 mL min1, the transferred charge was also only 1.84 mAh but the lambda value was around 3. At a flow rate of 5 mL min1, the lambda value of the anolyte was 22. Despite the
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relatively high lambda values there could be a significant change in the region of the double layer, due to the low kinetics of the anode reaction and because of the foregoing measurement with a lower flow rate, at which the equilibrium between the redox potential and the anode potential had not yet been achieved. For a further analysis during the linear sweep voltammetry, potential and voltage characteristics of the cell and the half-cells were plotted in Fig. 4. In the cell with 5 mL min1 the anode potential started at 0.127 V and reached 0.019 V at 23 mL min1. The difference of 146 mV is significant and was not due to obviously different redox potentials of the dithionite solution. For 5 mL min1, the redox potential was 0.218 V, and at 23 mL min1 it was 0.225 V, so the difference was only 7 mV. The first measured current density was 0.21 mA cm2 at 5 mL min1, and 0.28 mA cm2 at 23 mL min1, so this does not explain the difference either, especially since the difference during the test remained approximately constant. The differences between redox potential and the first measured value of the anode potential were 91 mV at 5 mL min1 and 244 mV at 23 mL min1. The cell at 23 mL min1 had about 2.5 times the activation overpotential for the anodic reaction compared to the cell at 5 mL min1. For both flow rates, the cathode potential showed negligible differences at the beginning of the measurements. Thus the difference in cell voltage at the beginning of the measurements was only caused by the different activation overpotentials of the anode reactions. As well as in the analysis of the OCV measurements no further explanation could be found here, as the equilibria were different at the anode from the solution and had not adjusted during the 60 s between measurements. By the use of higher flow rates it could be possible to reach the equilibrium more quickly. The cathode potential calculated at the beginning of the measurements had a value of 0.693 V, which was well below the standard potential of the 4-electron mechanism with 1.227 V, and slightly below the peroxide formation at 0.761 V. It can be assumed that the ORR proceeded with different limitations because the activation overpotential compared to the 4-electron mechanism was 534 mV and possibly peroxide was formed. The difference between the anode potential at the start of measurement and the final value was 609 mV at 5 mL/min, and 486 mV at 23 mL min1. The corresponding differences for the cathode potential were only about 200 mV. From these values it can be concluded that the rate of the anodic reaction was about 2.5e3 times lower than that of the cathode reaction. The potential of the
235
ORR, however, had a substantially higher deviation from the equilibrium potential than the anode reaction. The cathode potential changed after about 15 s measurement time to significantly lower values. The cathode potential of the cell at 23 mL min1 depolarized more intensely than at 5 mL min1. Probably the reaction mechanism of the anode reaction had changed during the measurement. As the anode potential of the cell at 23 mL min1 was generally lower than that of the cell with 5 mL min1, the cathode depolarized slightly more because a different reaction at higher anode potential was more pronounced than in the case of the cell with 5 mL min1. The kinetics of this by-reaction was greater than the S2O2 4 -oxidation, leading to a greater depolarization of the potential of the ORR. Since the curve bending of the cathode potential starts from an anode potential of about 0.1 V and a certain overpotential for the reaction is assumed, in particular the formation of 2 SO2 4 /SO3 - must be considered as a by-reaction during the linear sweep voltammetry under the applied potential. Fig. 5 shows the IR-corrected cell voltages and the half-cell potentials as Tafel-plots. The equilibrium voltages of the cell (Fig. 5a) differed significantly from each other for the reasons discussed above. At lower flow rates two different Tafel-slopes appear; at higher flow rates there is only one. The pattern of the voltages seems to indicate the start of transport limitations. In Fig. 5b, the cathode potential shows no significant change with the change of the flow rate of dithionite solution, so it can be concluded that the flow rate has no effect on the cathodic reaction mechanism. At high current densities the Tafel-slope was 215 mV/dec (a*n ¼ 0.27), which indicates an irreversible electrochemical reaction in this system. However, the linear range was not fully formed, so no conclusions could be made concerning eventual transport limitations or an exact value for a*n. When the curves of the cathode potential are compared with those of the anode potential and the cell voltages, it can be seen that the cathode reaction was responsible for the bending of the cell voltage at higher current densities. At 5 mL/min the anode Tafelslope was 318 mV/dec (a*n ¼ 0.19). It suggests that this reaction was even slower than the ORR and which limited the cell performance significantly. A significant mass transfer limitation could not be detected. The course of the anode potentials at low flow rates and current densities showed complications in the anode reaction, which are believed to have been caused by the above-mentioned problems i.e. low fluid supply or pressure at low dithionite lambda values. 3.3. Discharge analysis
Fig. 4. Cell and half-cell potentials of a cell with 0.85 M Na2S2O4 during linear sweep voltammetry with anolyte flow rates of 23 mL min1 and 5 mL min1.
Fig. 6 shows the potential profiles and the capacity during the discharge of a dithionite fuel cell. Some spikes are artifacts of the measurements and were not removed by smoothing methods. As regards the cell voltage, a depolarization of 642 mV can be observed from the beginning of the measurement at OCV 0.842 to the discharge voltage of 0.200 V. The cathode potential at OCV was 0.774 V and quickly depolarized to 0.693 V. The value of 0.693 V corresponded exactly to the value of the cathode potential at the beginning of the investigations using linear sweep voltammetry. Both values were significantly different to the theoretical value of standard potential of the 4-electron mechanism of the ORR. The activation overvoltage of the ORR was 81 mV. The anode potential changed from 0.07 V to 0.491 V by an amount of 561 mV. The kinetics of the anode reaction was approximately seven times slower than that of the ORR. As the experiment progressed the redox potential of the solution changed from 0.107 V to a value of 0.749 V at the end of the measurement. After a test period of 30 h a potential jump of the
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Fig. 5. Tafel plots of IR corrected cell voltage (a) and half-cell potentials (b) of a fuel cell with 0.85 M Na2S2O4 with variation of anolyte flow rate.
half-cells can be recognized, which is normally indicative of a change in the reaction mechanism. Up to this time 99% of the charge carriers were transferred with 2.57 Ah capacity. The following change of potentials hasn't resulted in a relevant charge transfer process. The test can therefore be divided into two parts. At first dithionite was oxidized until a time of about 30 h, and afterward significant changes in the cell occurred. In the first part, the OCV was on average about 0.7 V, in the second part at 0.3 V. In the first part the voltage progression was unusual. After a period of about 20 h, the terminal voltage rose again slightly and then dropped
sharply in the second part. Looking at the half-cell potentials one can see the increase of cathode OCP with time, which moved together with the trend of simultaneously increasing anode potential. In the second part the cathode OCP finally reached a plateau, in which further depolarizations were hardly discernible. Probably the three-phase interface of the cathode and in particular the surface of the Pt catalyst changed. It can be assumed that even in a short time the dithionite solution can pass through the membrane to the catalyst and can therefore alter the kinetics of the ORR. With the change in the composition of the dithionite solution the activity of the catalyst could have been changed at the same time, as evidenced here in an increasing cathode potential. However, the anode and cathode reactions are strongly pH-dependent (Equations (1) and (2)). For the complete reaction of 0.85 M S2O2 4 according to Equation (5), the redox potential should increase by 100 mV due to the pH decrease. The change in the cathode potential was about 150 mV, which was in the range of the possible change because large potential differences can be expected even at moderate pH values. The increase in the cathode potential was thus mainly a cause of pH value change, but shows at the same time that there must be a clear balance between the composition of the cathodic three-phase boundary layer and the anolyte solution, because a pH value increase can only be achieved by dithionite oxidation, which normally takes place in the anode half-cell. This relationship shows a continuous passage of liquid through the membrane towards the cathode, leading to the assumption that this was rather too wet to function more efficiently Table 1. 60 mL of 0.85 M Na2S2O4 solution were used during the experiment which correspond to a capacity of 2.73 Ah in a twoelectron transition. The measured value was 2.57 Ah and hence a charge carrier efficiency of 94.1% could be achieved. The loss of 5.9% charge could have had three main reasons. The first possibility is that the dithionite solution was already oxidized, since the preparation of the solution was not carried out under inert gas atmosphere. Oxidation of the solution could have occurred during the measurement due to diffused oxygen and the diffusion of dithionite
Table 1 Theoretical and measured voltages, capacities and energy densities of a cell with 0.85 M Na2S2O4.
Fig. 6. Potentials of a cell with 0.85 M Na2S2O4 at potentiostatic discharge at 0.2 V.
Theoretical Measured Efficiency [%]
Cell voltage [V]
Capacity [Ah L1]
Energy density [Wh L1]
1.53 0.2 13.1
45.5 42.8 94.1
69.6 8.6 12.4
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through the membrane. The redox potential of the solution at the beginning of the measurement was a bit higher than the theoretical amount of 0.294 V. This suggests an already partial oxidation of dithionite prior to measurement. No evidence was found of the diffusion of atmospheric oxygen into the solution through the cell or in the vessel or tubing, but this was not ruled out. On the other hand, the low cathode potential provides evidence of the passage of dithionite solution. Considering the entire loss of capacity of 160 mAh as a liquid passage with unchanged concentration through the membrane, 3 mL of anolyte entered the cathode halfcell during about 45 h continuous measurement. Probably all three reasons, to an unknown extent, led to the charge carrier loss. The operating voltage of the cell at 0.2 V represented a voltage efficiency of 13.1%. The low value was caused firstly by the low reaction kinetics of the oxidation of S2O2 4 and secondly by the low equilibrium potential of the ORR. The total amount of converted energy was 0.51 Wh, corresponding to an energy density of 8.6 Wh L1. The energy efficiency of the cell was 12.4%. 3.4. Electrochemical impedance spectroscopy In Fig. 7 the Nyquist and Bode plots of a cell with fresh 1 M Na2S2O4 are shown with a cell voltage of 0.871 V in a state without flowing anolyte and air. The EIS-spectra were fitted with a simplified equivalent circuit presented in Fig. 8. The ohmic resistance R1 of the cell was about 0.1 U, and the charge transfer resistance R2 was 14.7 U. Both values were relatively high, although R2 clearly dominated. The semi-circle of the time constant in the Nyquist plot had a significantly depressed semicircle with a capacitor exponent of 0.4 and a capacity of 0.77 F. The Nyquist plot of the anode impedance had an almost similar behavior and values as the cell impedance. The low kinetics of the anode reaction were therefore dominant throughout the cell. The deviation from the ideal capacitor behavior can be assigned by an increased response on adsorbates on the electrode surface. The Bode and Nyquist plot had a constant phase shift in the frequency range of 1e10 kHz, which had the character of a transmission line and is a characteristic of a highly developed porous surface. In the Bode plot the curves of the phase shift from the cell and anode showed negative phase shifts at high frequencies, while the impedance increased. For inductive effects, however, a positive phase shift is expected. These phenomena were not further evaluated. Probably overlays of disorders and inductive effects were the cause of these results. The Nyquist or Bode plots illustrated in Fig. 9 show the impedance spectra of a cell at different OCV after discharge. The measurements are taken from the discharge experiment illustrated
237
Fig. 8. Simplified equivalent circuit diagram for the modeling of EIS measurements of a cell with 0.85 M Na2S2O4 (R1 e ohmic resistance, R2 e charge transfer resistance of the cell reaction, CPE1 e double layer capacitance).
in Fig. 6, in which the voltage values correspond to the 1st, 50th, 70th or 100th discharge step. Due to the relatively complex results and in particular the lack of low frequency range (<1 Hz) no fitting or modeling has been performed. Nevertheless, some features and trends were qualitatively evaluated. In the low frequency range of 1e10 Hz, a time constant was dominant, which could be attributed mainly to the reaction of dithionite. With increasing discharge time, the charge transfer resistance of the reaction increased significantly. In the high-frequency part at about 10 kHz inductive effects in particular were detected during the first measurement. However, with increasing discharge time in the region of 1e100 kHz, a further time constant appeared, which has been superimposed with inductive effects, and thus did not form a typical semi-circle in the Nyquist plot. The high-frequency time constant was observed only in the impedance measurements of the cell, but not in those of the anode. Because of the high resonance frequencies, this time constant can be attributed to the electrolyte conductivity which decreased significantly during discharging. The decrease occurred in parallel with decreasing pH, which resulted from the nearly complete consumption of OH e ions and thus decreased the electrolyte conductivity during discharge. 4. Conclusions The results obtained show that the construction and operation of a dithionite/air fuel cell is feasible in principle. The power densities achieved were 1e3 decades below practical values, so it can be expected that such a non-optimized converter would lead to very high costs. The values are comparable with other systems in development and could certainly be further increased by optimizing the reaction kinetics, cell and plant management. By means of appropriate and cost-effective anode catalysts the dithionite oxidation could be accelerated and increased, while further studies on the behavior of the cathodes could enable their kinetics and equilibrium potential to be increased, improving power density and energy efficiency. The values obtained here were obtained at room
Fig. 7. Nyquist and Bode plots of a cell with 0.85 M Na2S2O4 at 0.871 V open circuit voltage.
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Fig. 9. Nyquist Bode plots of a cell with 1 M Na2S2O4 in 2 M KOH at different terminal voltages by discharging.
temperature. By increasing the temperature, the individual reaction rates would be likely to increase considerably; however, this would be associated with an increase in other technical or chemical complications such as cathode management and anolyte stability. The anolyte stability is of crucial importance for the cost, in order to achieve high numbers of cycles with high charge carrier efficiencies, and to keep the cost of reconditioning and anolytebalancing as low as possible.
[6]
[7]
[8]
Acknowledgments [9]
The authors would like to thank Dr. Carsten Cremers, Dr. Peter Fischer and Thomas Berger for the useful and interesting discussions on the subject, Florina Jung for the preparation of the CCM and Carolyn Fisher for translation and spell checking.
[10]
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