The Influence of Free Acid in Vanadium Redox-Flow Battery Electrolyte on “Power Drop” Effect by XRG Inc.

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The Influence of Free Acid in Vanadium Redox-Flow Battery Electrolyte on “Power Drop” Effect and Thermally Induced Degradation Nataliya V. Roznyatovskaya,* Matthias Fühl, Vitaly A. Roznyatovsky, Jens Noack, Peter Fischer, Karsten Pinkwart, and Jens Tübke of a VRFB is over 10 years and the extended lifetime tests are not always practical, the focus of many studies has been placed to the degradation phenomena and accelerated aging of VRFB or its components. One of the reasons of capacity loss during battery operation under long-term[1] or practical conditions[2] is transport of electrolyte species: protons, water, vanadium ions through membrane. Signs of degradation such as direction of water crossover flux, permeability for vanadium species, and so on appeared to be different in case of the VRFBs with anion and cation exchanger membranes.[2] Accelerated aging aims to test the cell over commonly used limits. Usually, the VRFB is charged–discharged at states-ofcharge (SoC) of 20–80% because of side reactions, which can occur at high voltages. In case of constant current discharge of a VRFB, which starts from a higher SoC (over 80%), the “power drop” effect has been recently and for the first time reported for a cell with anion exchanger membrane and vanadium electrolyte based on sulfuric acid matrix.[3] This effect, i.e., reversible decrease in a discharge voltage, is more pronounced after a longer exposure of cell to a high SoC, at higher discharge current densities or at low temperatures.[3] The appearance of “power drop” is suggested to be caused by reversible adsorption or temporary precipitation of vanadium(V) species onto membrane, which is a pH-dependent process.[3,4] In this work, we apply a fully charging of a VRFB with anion exchange membrane as a stressor to ascertain if there is any correlation between the initial electrolyte composition and a tendency of the VRFB to degradation. The initial electrolyte is often composed of a 50%:50% mol mixture of vanadium(III) and vanadium(IV) species in sulfuric acid and is denoted further as V3.5þ electrolyte. This electrolyte needs to be precharged directly in the VRFB for further battery operation. The “power drop” effect and ex situ thermally induced degradation of catholyte are to be considered because both of these phenomena are pH-dependent.[2,3] This, in turn, allows one to precise the optimal ratio of free sulfuric acid to vanadium species for electrolyte preparation to target the electrolyte formulation and therefore stable VRFB operation. As the VRFB electrolyte is commonly produced by chemical or electrolytic dissolution of vanadium raw compounds (vanadium pentoxide, vanadyl sulfate) in sulfuric acid, the concentration of free sulfuric acid in the final (V3.5þ) electrolyte or in the battery

A series of vanadium redox-ﬂow battery (VRFB) electrolytes at 1.55 M vanadium and 4.5 M total sulfate concentration are prepared from vanadyl sulfate solution and tested under conditions of appearance of “power drop” effect (discharge at high current density from high state-of-charge). A correlation between the initial electrolyte composition, the thermal stability of catholyte, and the susceptibility of VRFB to exhibit a “power drop” effect is derived. The increase in total acidity to 3 M, expressed as concentration of sulfuric acid in precursor vanadyl sulfate solution, enables “power drop”-free operation of VRFB at least at 75 mA cm 2. Thermally-induced degradation of electrolyte is evaluated based on decrease in vanadium concentration in the electrolyte series after exposure to the temperature of 45 C and based on characterization of catholytes series using 51V, 17O, and 1H nuclear magnetic resonance spectroscopy.

1. Introduction In the past decade, the vanadium redox-flow battery (VRFB) has become a well-developed and commercialized technology for a long-term energy storage and conversion. As the target lifetime Dr. N. V. Roznyatovskaya, M. Fühl, J. Noack, Dr. P. Fischer, Prof. K. Pinkwart, Prof. J. Tübke Fraunhofer Institute for Chemical Technology Applied Electrochemisrty Joseph-von-Fraunhofer-Str. 7, Pfinztal 76327, Germany E-mail: nataliya.roznyatovskaya@ict.fraunhofer.de Dr. N. V. Roznyatovskaya, J. Noack, Dr. P. Fischer, Prof. K. Pinkwart, Prof. J. Tübke German-Australian Alliance for Electrochemical Technologies for Storage of Renewable Energy (CENELEST), Mechanical and Manufacturing Engineering University of New South Wales (UNSW), UNSW Sydney, NSW 2052, Australia Dr. V. A. Roznyatovsky Chemistry Department M. V. Lomonosov Moscow State University Leninskiye Gory 1-3 GSP-1, 119991 Moscow, Russian Federation The ORCID identification number(s) for the author(s) of this article can be found under https://doi.org/10.1002/ente.202000445. © 2020 The Authors. Published by Wiley-VCH GmbH. This is an open access article under the terms of the Creative Commons Attribution License, which permits use, distribution and reproduction in any medium, provided the original work is properly cited.

DOI: 10.1002/ente.202000445

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after the ﬁrst precharging remains only roughly known.[5] A reason for this is that the conversion of vanadium(V) or vanadium(IV) species is accompanied by protons consumption or liberation. Therefore, the free acid concentration in electrolyte depends, among other things, on transport properties of membrane and cannot be generally taken as measurable parameter. In contrast to it, the total concentration of vanadium species, sulfate (i.e., the total concentration of SO42 , HSO4 , and H2SO4 species), and conductivity are available parameters for electrolyte evaluation in practice. However, the free acid concentration in VRFB electrolyte is expected to inﬂuence: 1) the theoretical open-circuit voltage (OCV) of the battery,[6] 2) side reactions such as hydrogen gas evolution,[2,7] 3) thermal stability of V(V) electrolytes at high temperatures,[2] and 4) kinetics of positive half-cell electrode reaction and therefore activation loss as well as shunt currents. The available literature data on the optimal electrolyte formulation are controversial. The optimal electrolyte utilization is attained, if 2.5 M sulfuric acid is taken for preparation of 1.5 M vanadium electrolyte.[6] Electrolyte prepared from 1.6 M vanadyl sulfate in 2.8 M sulfuric acid is reported to be optimal because of the coupling effect of viscosity, conductivity, and electrochemical activity.[8] A 2 M vanadium electrolyte containing 5 M total sulfate is often selected as the preferred electrolyte for increased conductivity.[9] The preferred vanadium electrolyte composition for improving the electrolyte thermal stability is 1.5–1.6 M vanadium in 4–5 M sulfuric acid.[10] In this work the series of VRFB electrolyte is prepared at the constant vanadium (1.55 M) and constant total sulfate (4.6 M) concentration to investigate the effect of the free acid concentration on the aforementioned VRFB degradation.

2. Results and Discussion 2.1. Preparation of Electrolyte Series One of the ways to produce VRFB electrolyte is based on the dissolution of vanadyl sulfate in sulfuric acid followed by electrolysis. The maximal vanadium concentration in this case is limited by the solubility of vanadyl sulfate in sulfuric acid that is dependent on total sulfate/bisulfate concentration. Another parameter to choose the optimal molar ratio of VOSO4 to sulfuric acid is the conductivity of the resulting V(IV) solutions, which is primarily provided by H2SO4 and should be high to minimize the ohmic polarization.[11] Generally, the conductivity depends on the total concentrations of vanadium, sulfate/bisulfate, and free protons at each given SoC. As shown in Figure 1, the maximal conductivity of V(IV) electrolyte solutions is attained at H2SO4 concentrations higher than 4 M. This corresponds to the vanadium concentrations lower than 1 M. As the vanadium concentration of VRFB commercial electrolytes is commonly in the range from 1.6 to 2 M, the solution of 1.6 M VOSO4 in 3 M H2SO4 (total sulfate concentration of 4.5 M) was chosen to compose the V(IV)-n series. To compose series of V(IV) electrolytes with different free acid contents, the concentration of H2SO4 is commonly varied at constant VOSO4 concentration.[9,12] In vanadium-free aqueous solutions, the solvation structure of sulfuric acid remains unchanged until above 6 M of total concentration, whereby both

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Figure 1. The solubility of VOSO4·H2O in sulfuric acid matrix. Conductivity, maximal concentration of VOSO4, and total sulfate concentrations are shown in dependence on the concentration of sulfuric acid matrix.

of bisulfate and sulfate anions are present in solution according to the following dissociation equilibria because the second step of dissociation is incomplete[13] H2 SO4 ↔ Hþ þ HSO 4

(1)

þ 2 HSO 4 ↔ H þ SO4

(2)

Vanadyl sulfate is known to partially form ion pair [VOSO4]0 in aqueous solution[14] ½VOSO4 0 ↔ VO2þ þ SO2 4

(3)

The sulfate ion is supposed to be more prone to form ion pairs [VOSO4]0 with vanadyl cation than bisulfate, if sulfuric acid present in the solution.[15] However, the association of vanadyl cation with bisulfate cannot be excluded.[4,16] The electrolyte series with varied H2SO4 and constant VOSO4 concentration displays an increase in ionic strength with an increase in H2SO4 concentration because the total sulfate/bisulfate concentration does not remain constant within this electrolyte series and it contributes more to ionic strength than protons. It results consequently in increase in the ionic strength fraction of VOSO4 and viscosity of the electrolyte solution.[15] Moreover, change in ionic strength would inﬂuence the dissociation of vanadyl sulfate ion pair[15] and sulfuric acid.[13] On the contrary, as sulfate is expected to contribute more to the ionic strength of electrolyte solution, to maintain constant total sulfate concentration for all the samples of the series seems to be reasonable. To provide constant total sulfate concentration at a preset concentration of V(IV), magnesium sulfate was used to partially replace sulfuric acid in electrolyte formulation due to its high solubility in acidic solutions (Table 1). Furthermore, the conductivity was measured and compared for both series: 1) with constant total sulfate concentration (surface (1), Figure 2) and 2) with varied total sulfate content (surface (2), Figure 2). As shown in Figure 2 (surface (1)), the conductivity seems to be insensitive to the variations of vanadium concentration in the range from 0.75 to 1.6 M. However, the change in H2SO4 content from 3 to 1 M results in the gradual decrease in conductivity from 300 to 70 mS cm 1. According to the literature, there is uncertainty about the nature and extent of dissociation of magnesium sulfate in acidic solutions.[17] However, the increase in MgSO4 © 2020 The Authors. Published by Wiley-VCH GmbH

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www.advancedsciencenews.com Table 1. Composition of V(IV) electrolyte samples prepared by dissolution of VOSO4 in sulfuric acid with addition of MgSO4. Sample

Molar content [%]

Total concentrations [M]

Conductivity [mS cm 1]

V(IV)

Vanadium

Sulfatea)

H2SO4

V(IV)-3

100

1.53

4.53

3.0

357

V(IV)-2.5

100

1.56

4.56

2.5

0.5

245

V(IV)-2

100

1.53

4.53

2.0

1.0

173

V(IV)-1.5

100

1.56

4.56

1.5

103

MgSO4

Value calculated as a total sulfate amount from all constituents.

MgSO4 þ z M H2SO4) was chosen to prepare further model series of V3.5þ electrolytes, whereby the ratio of total sulfate to the total vanadium concentrations ([SO4]T:[V]T) remained constant and the ratio of free acid to vanadium concentration ([H2SO4]free:[V]T) was varied within the series (Table 2). The samples V(III)-n and V(V)-n, which were prepared from V(IV)-n by exhaustive electrolysis in a cell with anion exchange membrane, were analyzed by titration (Table 2). It can be seen that both of V(V)-n and V(III)-n concentrations are for 0.02– 0.03 M lower than the concentration of precursor V(IV)-n solutions. One of a reason for it may be the loss of vanadium species as vanadium pentoxide precipitates on the anion exchange membrane, which is supposed to happen during the electrolysis to high SoC.[3,11] During the electrolysis of V(IV) solutions, sulfate anions are expected to move through membrane from anolyte to catholyte according to the commonly suggested mass balance equations (Equation (4)–(5)).[19] 2VOSO4 þ 2H2 SO4 þ 2e

ðanolyteÞ

(4)

! V2 ðSO4 Þ3 þ 2H2 O þ SO2 4 ðcatholyteÞ

2VOSO4 þ 2H2 O þ SO2 4 2e

(5)

! ðVO2 Þ2 SO4 þ 2H2 SO4

Figure 2. Conductivity of VOSO4 solutions for (surface 1) series with constant total sulfate concentration x M VOSO4 þ y M MgSO4 þ z M H2SO4, where x þ y þ z ¼ 4.6 M, (surface 2) mixtures x M VOSO4 þ z M HSO4, where 0.75 < x < 1.5 and 1 < z < 3.

SO2 4 should cross through the anion exchanger. Due to this transport phenomenon, the sulfate species are to be depleted in V(III) electrolyte and to be enriched in V(V) electrolyte solution. Therefore, if such prepared V(III) solution is mixed with precursor V(IV) solution in 1:1 v/v ratio to obtain V3.5þ electrolyte, the total sulfate concentration in V3.5þ electrolyte is expected to be lower than in original V(IV) electrolyte (the total sulfate concentration is expected to be of 4.1 M in V3.5þ according to Equation (4) and (5)). However, this suggestion Table 2. Composition of V(III), V(V), and V35þ electrolyte samples. Sample

concentration in sulfuric acid mixtures results in an decrease in proton concentration.[18] Under these conditions the conductivity of electrolyte is then likely to be related to the total acid concentration and is an experimentally accessible parameter. In contrast, the conductivity of the binary mixture (VOSO4 þ H2SO4) displays a more complex dependence on vanadium and sulfuric acid concentrations (surface (2), Figure 2). In particular, the conductivity of electrolyte samples (surface (2), Figure 2) remains almost unchanged in the range of H2SO4 concentrations from 1.5 to 3 M at the V(IV) concentration of 1.6 M. The variation of V(IV) concentration from 0.75 to 1.6 M leads to considerable changes in conductivity at each value of sulfuric acid concentration. The attainable conductivity values for surface (1) in Figure 2 are generally lower than the values for surface (2). It can be explained by ion association of vanadium species, which can occur in the presence of magnesium sulfate and by dissociation of sulfuric acid, which is suppressed with increasing the total sulfate concentration (Equation (1)–(2)). However, the series of V(IV) electrolyte based on ternary mixtures (x M VOSO4 þ y M Energy Technol. 2020, 8, 2000445

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Preparation procedure

Molar content [%]

Conductivity Total concentrations [M] [mS cm 1]

V(III) V(IV) V(V) Vanadium Sulfatea) V3.5-3 V3.5-2.5 V3.5-2

1:1 v/v mixture of V(III)-n and V(IV)-n where n ¼ 3; 2.5; 2; 1.5

V3.5-1.5 V(III)-3 V(III)2.5 V(III)-2

Electroreduction of V(IV)-n where n ¼ 3; 2.5; 2; 1.5

V(V)-3

V(V)-2

1.55

4.50

283

51.3 48.7

1.55

4.38

198

49.0 51.0

1.55

4.47

130

50.6 49.4

1.55

4.38

100

1.53

250

100

1.54

158

99.3

V(III)1.5

V(V)-2.5

50.9 49.1

0.7

100 Electrooxidation of V(IV)-n where n ¼ 3; 2.5; 2; 1.5

V(V)-1.5

1.53

–

1.4

98.6

1.54

445

–

1.6

98.4

1.52

342

–

1.1

98.9

1.53

256

–

1.4

98.6

1.52

188

Determined gravimetrically.

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cannot be conﬁrmed by experimental data on total sulfate concentration in V3.5-n series (Table 2), which is determined to be 4.4–4.5 M and is very close to 4.5 M of total sulfate estimated for V(IV)-n solutions (Table 1). This indicates that either the chemical state of vanadium species and consequently mass balance differs from that in Equation (4) and (5), or the diffusion of other species than sulfate occurs through membrane during electrolysis. Figure 3 shows the conductivities of V(III), V(V), and V3.5þ electrolytes prepared from the V(IV)-n series. As the total vanadium and total sulfate concentrations within the V3.5-n series remain unchanged for all the values of n (Table 2), the difference in V3.5-n conductivity values for each n is to be linked primarily to the difference in protons concentration and conﬁrms the increasing amount of free sulfuric acid in V3.5-n samples for n-values from 1.5 to 3. Because the electrode reaction of cathodic V(V)/V(IV) couple is pH-dependent (Equation (5)), the cell

Figure 3. Conductivity of V(III), V(IV), V(V), and V3.5þ electrolyte solutions in dependence on the amount of sulfuric acid in the precursor V(IV)-n solution series.

voltage (i.e., the difference between cathodic and anodic half-cell potentials) and in particular the OCV of VRFB operated with V3.5-n electrolytes is expected to vary within the electrolyte series. To estimate the OCV in case of V3.5-n electrolytes, the half-wave potentials were considered as formal potentials of cathodic and anodic redox couples. The half-wave potentials (E1/2) and peak-topeak separations (ΔE) for V(V)/V(IV) and V(III)/V(II) couples were determined by cyclic voltammetry at glassy carbon electrode using V(IV)-n and V(III)-n probes (Figure 4 and Table 3). As shown in Table 3, the electrode reactions of both V(V)/ V(IV) and V(III)/V(II) couples are quasireversible, and the maximal difference in OCV values within the series is 0.03 V for the samples V3.5-3 and V3.5-1.5. With the decrease in n-value in V(III)-n and V(IV)-n samples, the currents for both redox couples at the voltammogramms decrease (Figure 4). As the currents are proportional to the diffusion coefﬁcient of redox active species, the diffusion of vanadium species is hindered with the decrease in n-value within V(III)-n and V(IV)-n series. The concentration of magnesium sulfate increases within the series and by analogy to the binary mixtures VOSO4 and H2SO4[15] the viscosity of electrolyte solutions within the series can increase for lower n-values. It can, in turn, inﬂuence the diffusion of vanadium species and explain the difference in diffusion currents at the cyclic voltammogramms (Figure 4). Although the expected difference in OCV for the cell with V3.5-n electrolytes is relatively small, i.e., of 30 mV, commonly used galvanostatic charge–discharge protocol with preset cutoff voltages would not be suitable for the cells with V3.5-n electrolytes. Moreover, the differences in conductivity of V3.5-n probes (Table 2) would result in variations in attainable SoC values. Therefore, the following protocol was further used to the cells operated with V3.5-n electrolytes and to investigate the degradation phenomena: 1) Evaluation of initial cell resistance from electrochemical impedance spectrum (EIS). 2) Galvanostatic charge followed by potentiostatic charge step to possible highest SoC. 3) Determination of the cell resistance from EIS to compare with initial cell resistance value. 4) Galvanostatic discharge at relatively high current density to enable the appearance of “power drop-off

Figure 4. Cyclic voltammogramms at glassy carbon electrode in electrolyte samples a) series V(III)-n, b) V(IV)-n, where n is 3, 2.5, 2, and 1.5. Curves are recorded at 0.02 V s 1 and corrected for ohmic drop using post-iR correction of data.

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www.advancedsciencenews.com Table 3. Electrochemical characteristics of redox couples of series samples V(III)-n and V(IV)-n evaluated from cyclic voltammetric measurements. Sample n

ΔE [V]

E1/2 [V] versus Hg/Hg2SO4

OCV [V]

V(III)/V(II)

V(V)/V(IV)

V(III)/V(II)

V(V)/V(IV)

0.988

0.452

0.19

0.14

1.440

2.5

0.99

0.447

0.2

0.14

1.437

0.988

0.438

0.22

0.14

1.426

1.5

0.985

0.428

0.22

0.14

1.410

effect.”[3] 5) Investigation of catholytes, i.e., V(V)-n at high SoC at room and elevated temperature. 2.2. Accelerated Aging: Charge to High SoC and Discharge (“Power Drop” Effect) The discharge curves for the cell operated with V3.5-n electrolytes are shown in Figure 5. These discharge steps were preceded by the charging of the cell to high SoC, so that the cells displayed the OCV of 1.64 V before discharge. The “power drop” effect implies the appearance of a pit at the voltage–time curve at the beginning of the discharge,[3] provided that the half-cell potentials do not exhibit any pit. It is seen that the pit is deeper for those electrolytes, which are less conductive, i.e., which are prepared from the V(IV) precursor containing less sulfuric acid. However, no pit can be seen at the curve (1). This curve corresponds to the electrolyte with a highest concentration of

Figure 5. Discharge curves for cells operated with V3.5-n electrolyte series at a current density of 75 mA cm 2. The cells displayed the cell voltage of 1.64 V before the discharge.

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sulfuric acid within the V3.5-n series. The occurrence of “power drop” effect is likely to depend on the molar ratio of sulfuric acid to vanadium in the V3.5þ electrolyte. As shown in Figure 5, the cell can be discharged from high SoC without power drop, if the ratio [H2SO4]free:[V]T in V(IV) electrolyte precursor solution is 2:1 (at least for the series at [SO4]T:[V]T of 2.9). The “power drop” is attributed previously to an increase in membrane resistance, which is caused by adsorption of V(V) species or temporary precipitation of vanadium pentoxide within the pores of anion exchange membrane in contact with catholyte.[3] The accumulation of insoluble V(V) species from catholyte was previously conﬁrmed for cell with anion exchange separator operated at SoC from 10% to 80%.[11] However, anion exchange membranes are often used for VRFB cell construction due to lower permeability of V(II), V(III), and V(IV) components of electrolyte than cation exchanger.[11] To verify, if the membrane in the cell becomes blocked by V(V) species, the cell resistance was measured by electrochemical impedance before the charging step and directly before the discharge, i.e., at the high SoC of the cell (Figure 6). Indeed, apart from the membrane resistance other components such as electrolyte conductivity, resistance of electrodes, and contact resistance contribute to the common cell resistance. We assume that all these parameters for the same cell except for membrane resistance in case of deposition of vanadium species undergo inconsiderable changes or remain unchanged. As shown in Figure 6, impedance spectra for the cells with V3.5-n electrolytes move to the direction of higher resistances within the series V3.5-n with the decrease in n value. Moreover, there is a difference in the cell resistance between the cell with initial V3.5þ electrolyte and the same cell being fully charged. This difference is more pronounced for the low n-values within V3.5-n series (Table 4), i.e., for the cells, which exhibit “power drop” effect (Figure 5). Therefore, the decrease in current at the beginning of discharge step, i.e., “power drop,” seems to be linked to the increased cell resistance before the discharge step, which, in turn, is to be attributed to the increase in membrane resistance presumably because of adsorption of V(V) species. As the

Figure 6. Nyquist plots for the freshly assembled cells (1a–4a), which are fed with electrolytes V3.5-n, and at high SoC before discharge (1b–4b).

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www.advancedsciencenews.com Table 4. Changes of the cell resistance derived from Nyquist plots for the cells, which are fed with electrolytes V3.5-n before the charge step (initial state) and directly before the discharge step. Sample n

Rcell, initial [Ohm]

Rcell,

before discharge

ΔR ¼ Rbefore

discharge – Rcell, initial

[Ohm]

0.074

0.075

0.001

2.5

0.076

0.082

0.006

0.089

0.101

0.012

1.5

0.119

0.147

0.030

adsorption of V(V) species at the membrane is assumed to be pH-dependent, the data in Figure 5 can be interpreted in terms of optimal molar ratio [Hþ]:[V]T, which disables power drop during discharge at given cell operation conditions and cell configuration. The sufficient amount of free acid in electrolyte and efficient transport or generation of protons is required for the stable operation of VRFB at high discharge currents or at high SoC. On the contrary, the curve (1) in Figure 5 corresponds to the magnesium-free electrolyte solutions in contrast to the curves (2)–(4), which show the pit. In this case, it can be generally assumed that the “power drop” effect can be caused as well by magnesium ions in electrolyte solutions. However, this assumption would be inconsistent with the fact that magnesium sulfate is often used as a reference vanadium-free electrolyte to study the ions transport or crossover through membranes for redox-flow batteries.[20] To summarize, the initial V(IV)-n electrolyte solutions, which were used to prepare the V3.5-n series, display a decrease in conductivity for lower n-value. This decrease in conductivity within the series can be related to the increase in viscosity due to increasing amount of MgSO4 and suppression of the second dissociation step of sulfuric acid (Equation (2)). The cells operated with V3.5-n electrolytes derived from V(IV)-n precursors display “power drop” effect for n-values from 2.5 to 1.5.

Figure 7. Vanadium concentration of V(V)-n probes a) before and b) after being kept a day at þ45 C. The gray bars in Figure 7a show to the conductivity values of V(V)-n electrolytes.

had the highest stability and remained without any sign of precipitation after a day at þ45 C. All the other samples of V(V)-n series displayed lower concentration of V(V) because of thermally induced precipitation of V(V) species. The V(V) species are known to be involved in several equilibria, and vanadium(V) speciation in catholyte is complex.[4,21] Therefore, to ascertain further if the variation of sulfuric acid concentration in V(IV)-n can inﬂuence the chemical structure of the V(V) species within the V(V)-n series, the V(V)-n samples were investigated by dynamic light scattering (DLS) and nuclear magnetic resonance (NMR) techniques. The V(V) species in electrolyte are prone to form dimeric and oligomeric structures,[4] and the particle size distributions in V(V)-n samples at room temperature are found to be bimodal except for the sample V(V)-1.5, where three fractions of particles are present (Figure 8).

2.3. Accelerated Aging: Exposure of Catholyte at High SoC to Higher Temperatures The V(V) electrolytes, i.e., fully charged catholytes, are known to be unstable at elevated temperatures because of irreversible formation and precipitation of vanadium pentoxide.[4] The stability of electrolyte at high temperatures, expressed in terms of lag period before the precipitation starts, is to be pH-dependent.[4] Therefore, the behavior of V(V) at elevated temperatures was used as accelerated degradation test to verify the role of free acid in electrolyte in the degradation process. The probes of V(V)-n prepared by electrolysis of V(IV)-n electrolyte series were exposed to temperature of þ45 C for a day. The concentration of V(V) species was determined in V(V)-n probes before and after the thermally induced precipitation (Figure 7). As it was expected, the decrease in V(V) concentration in the aqueous phase (Figure 7b) correlates with the conductivity of V(V)-n solution or with the concentration of free sulfuric acid in the precursor V(IV)-n electrolyte. Therefore, the sample V(V)-3

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Figure 8. Size of vanadium species in the V(V)-n samples derived from DLS measurements (n is 3, 2.5, 2, 1.5 and corresponds to the concentration of sulfuric acid in the parent V(IV)-n series).

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The DLS measurements were repeated for reproducibility and the data in Figure 8 are shown for two trials (serie 1 and serie 2) as an average value of five measurements for each sample. All the samples feature the fraction of particles centered at about 4000 nm independently of sulfuric acid concentration in the parent V(IV)-n series. The size of the particles with the diameter below 10 nm seems to slightly depend on the concentration of sulfuric acid or conductivity of the samples, so that the larger particles are present in the sample with the lowest amount of acid. This particle size distribution differs from that one registered for 1.6 M vanadium electrolyte at 4 M total sulfate, stabilized by phosphoric acid.[21] In this case, there was only one fraction of particles centered at 1 nm at room temperature. This fraction was attributed to vanadium(V) dimers. All the samples of V(V)-n series do not have any stabilizing additive and exhibit the presence of larger particles (oligomers) already at room temperature. Apart from the oligomers, V(V)-n solutions are supposed to contain a mixture of VO2þ monomer, dimer, deprotonated forms or species associated with sulfate, or bisulfate.[21] The NMR spectra give usually an unresolved signal for this mixture, which, however, can be sensitive to the distribution of vanadium species, i.e., to the share of dimers. The 51V, 17O, and 1H NMR spectra were recorded for V(V)-n samples with n to be 3, 2.5, 2 because the V(V)-1.5 demonstrated the signs of aging, i.e., precipitate formation, during storage at already room temperature over several months. The 51V NMR spectra of V(V)-n series (Figure 9a) are characterized by the presence of a line at 558 ppm for all three samples. Compared with the literature data,[22] the 51V NMR spectra of V(V)-n series can be interpreted as a signal corresponding to the mixture VO2þ and [V2O3]4þ, whereby the share of dimers is low and monomer dominates. The 17O NMR spectra (Figure 9b) display two peaks for V(V)-n series: O(1) and O(2). The line O(1) is typically attributed to the —OH, —O— groups of sulfuric acid or water, and to the oxygen groups of vanadium core, which are in a rapid exchange with water ligand by the protonation/deprotonation reaction. The chemical shift of the peak O(1) and its line width depend on the acid concentration. The peak O(1) shifts upfield and line O(1) becomes broader with the decrease in acid concentration in the parent V(IV)-n series. In contrast, the peak O(2) is centered at 163 ppm (0.6 kHz) for all the V(V)-n samples and is attributed to the oxygen moiety of sulfuric acid. The 1H NMR spectra of the V(V)-n series (Figure 9c) are presumably a superposition of 1H resonances for H3Oþ, —OH groups of sulfuric acid, and vanadium species. The chemical shift varies from 7.46 to 6.8 ppm for the samples V(V)-n with the decrease in acid content. However, the chemical shift of 1H NMR is generally small, so that it is difficult to assign its variation to any process. Compared with the NMR spectra of 1.6 M vanadium electrolyte stabilized by phosphoric acid,[21] the 51V and 17O lines in NMR spectra of V(V)-n series are much broader with line width of 2.2–2.7 kHz and of 0.9–1.4 kHz, respectively. It may be consistent with the fact that V(V)-n electrolytes contain a fraction of larger species (Figure 9) than the stabilized V(V) electrolyte. The temperature-dependent changes in 51V, 17O, and 1H NMR spectra (Figure S1, Supporting Information) are shown in Table 5. As the thermally induced precipitation of vanadium species is assumed to proceed through the protonation of oxo groups at V═O moiety, broad 51V and 17O lines at low

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Figure 9. a) 51V, b) 17O, and c) 1H NMR spectra of V(V)-n (n is 3, 2.5, 2) solutions at room temperature ( 23 C).

Table 5. Variations of NMR spectra features for V(V)-n series with increase in temperature. Line

Chemical shift

Line width

Downﬁeld congruently

Parabolic-shaped/minimum at 20–30 C

Downﬁeld

Narrowing to the same, lower value of 0.7 kHz

Unchanged

Narrowing to the same, lower value of 0.4 kHz

Upﬁeld

Parabolic-like/maximum at 25 C

V O O(1) O O(2)

temperatures are signs of slow proton exchange processes, which means slow kinetics of vanadium pentoxide formation. The difference in chemical shift of 51V and 17O lines for V(V)-n samples with various acidity, which can be seen at low temperatures, becomes minimal at the temperatures over 30 C, whereby the 51V NMR exhibit the line at 555 ppm, which is more close to the value of 545 ppm for monomeric VO2þ species. VO2þ is the ﬁnal species, remaining in the electrolyte solution after thermally induced precipitation. To summarize, the difference in V(V) species within the V(V)-n series concerns the vanadium-oxygen core and no any other changes in chemical speciation of vanadium are assessed by NMR or DLS technique. With the increase in sulfuric acid concentration, the processes involving proton exchange become

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more rapid and the corresponding lines in the NMR spectra are more narrow.

3. Conclusion The series of 1.55 M V(IV) electrolyte solution was prepared with constant total sulfate concentration of 4.5 M and varied concentration of sulfuric acid. This series of electrolytes is characterized by variation in conductivity and presumably by suppression of second dissociation step of sulfuric acid. The investigation of V3.5þ electrolytes and catholyte at high SoC, obtained from the precursor V(IV) series, shows that the amount of free sulfuric acid in VRFB electrolyte or presumably the molar ratio [Hþ]:[V]T at given total sulfate concentration is likely to define whether the VRFB operated with this electrolyte is prone to exhibit such degradation processes as “power drop” or thermally induced precipitation of vanadium(V) species at 45 C. The amount of free sulfuric acid in the final V3.5þ electrolyte is expected to depend on the way of electrolyte preparation, i.e., on the type of vanadium source compound (vanadyl sulfate or vanadium oxides). The transfer of sulfate/bisulfate during the electrolysis of V(IV) solutions through anion exchange membrane to keep mass balance could not be confirmed by gravimetric analysis of resulting V(III) and V(V) electrolyte samples within the series under consideration. The estimate of the sole value of total sulfate concentration in electrolyte is not enough to predict susceptibility of VRFB to reversible or irreversible degradation. The electrolyte conductivity is to be taken as an important parameter of electrolyte evaluation.

4. Experimental Section Chemicals: Sulfuric acid (95%, supra) was Carl-Roth reagent. MgSO4·7H2O (puriss. p.a.) was purchased from Sigma-Aldrich. Vanadyl sulfate hydrate (technical grade) was purchased from Chempur and used without puriﬁcation. The amount of crystallization water in VOSO4·xH2O was determined by potentiometric titration. VOSO4·H2O was obtained by thermal dehydratation of VOSO4·xH2O and used for solubility trials. Electrolyte Preparation: The series of vanadium(IV) electrolytes (for brevity, the VRFB electrolytes are denoted in the text according to the oxidation state of vanadium species as V(V), V(IV), V(III), and V(II)) was prepared by dissolution of vanadyl sulfate (VOSO4·xH2O, where x is from 5 to 6) and MgSO4·7H2O in water after addition of preset amount of sulfuric acid (Table 1) to provide the constant total sulfate concentration throughout the series. For comparison, the series of vanadyl sulfate solutions was prepared in sulfuric acid also without addition of magnesium sulfate and used for conductivity measurements. V(III) and V(V) forms of electrolyte in sulfuric acid were produced by electrolysis of V(IV) electrolytes (Figure 10). V3.5þ electrolyte probes were obtained by mixing the electrolytes containing V(III) and V(IV) redox forms. A 40 cm2 cell of the same design as described for the cell test (see Section 4.4) was used for electrolysis. The electrolysis was conducted under galvanostatic–potentiostatic mode with maximal voltage of 1.65 V and maximal current of 3 A. Titration and Gravimetric Analysis: The potentiometric titration was conducted with 0.1 M cerium(IV) sulfate standard solution (Carl-Roth, Germany) using Titrator T70 (Mettler Toledo Int. Inc., Germany) to determine the total vanadium concentration and the molar content of vanadium in various redox forms in electrolyte samples. The total concentration of sulfate ions electrolyte samples was determined by gravimetric analysis with barium chloride.

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Figure 10. Schematic representation of the preparation pathways for series of electrolyte samples used in this work.

Conductivity Measurement: The measurement of electrolytes conductivity was performed at room temperature in a four-electrode glass cell with a cell constant of 7.3 mS cm 1 and glassy carbon or graphite electrodes using AC impedance technique. The cell was calibrated with a standard conductivity solution (500 mS cm 1). The frequency range from 10 to 100 kHz with AC of 10 mA and 0 DC was used to measure the highfrequency resistance values. Cyclic Voltammetry: V(III) and V(IV) samples were used for cyclic voltammetric investigations in anodic and cathodic regions of potential, respectively. Cyclic voltammetric curves were recorded in a three-electrode glass cell at room temperature using a Gamry Reference 3000 potentiostat. All potentials were measured and reported versus a Hg/Hg2SO4 reference electrode (0.65 V vs NHE). A platinum plate served as a counter electrode. A glassy carbon disc working electrode (1 mm diameter, ALS, Japan) was polished using diamond pastes (6, 1, 0.25 μm) and thoroughly washed with water. After conventional polishing, the glassy carbon electrode was additionally electrochemically pretreated. This pretreatment was conducted in a separate three-electrode cell filled with 2 M H2SO4 solution by application of six cyclic potential scans in the range from 0 to 2.0 V versus an Ag/AgCl reference electrode at 20 mV s 1. The final applied potential was 0 V. Particle Size Distribution Measurement: The presence of particles in the V(V) electrolyte samples was examined by means of DLS technique by use of Zetasizer Nano ZS (Malvern Instruments Ltd., England) with a noninvasive backscatter optic (173 scattering optic). The light source was a He– Ne laser with 633 nm. The size distribution curves were registered at room temperature periodically within the preset time intervals. The used dispersant refractive index for the DLS is 1.4. The DLS measurements were repeated for reproducibility and the results are given for both trials (serie 1 and serie 2) as an average value of five measurements for each sample V(V)-n (n ¼ 3; 2.5; 2; 1.5). NMR Spectroscopy: Agilent MR-400 NMR spectrometer with 5 mm OneNMR probe was used to record 1H, 51V, and 17O NMR spectra at 400.0, 105.0, and 54.2 MHz, respectively. All measurements were done without stabilization of resonance conditions, i.e., without internal lock, whereby D2O was used as an external reference. VOCl3, H2O, and tetramethylsilane were taken as standards to measure chemical shifts for 51V, 17 O, and 1H, respectively. To acquire the 1H spectrum, single scan was used within 6.4 kHz spectral window. The parameters to record the 51 V, 17O, and NMR spectra were as follows: 512 scans within 89.0 kHz and 15 000 scans within 73.5 kHz window. Variable temperature measurements were conducted in 7 C (7 K) steps from 0 to 63 C (273 to 336 K). 1H, 51V, and 17O NMR spectra registration was performed under the same parameters as registration at room temperature. The electrolyte samples were first tempered at preset temperature and then the NMR spectra were taken after 10–20 min in the following sequence: 1H, 17O, and 51V. The integral intensities and frequencies of resonance lines in 17O, 51V, and 1H spectra were determined using the homemade software Intspect.[23] Ex Situ Thermal Stability Test: For thermal stability test, electrolyte samples of 1 mL in closed vials were kept in a test chamber conditioned

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www.advancedsciencenews.com at 45 C (Weiss WKL 64, Germany) for a predeﬁned period of time and the molar ratio of V(V) to V(IV) and total vanadium concentrations were determined afterward by titration. Charge–Discharge Experiments: Charge–discharge trials were performed at room temperature in a cell with 40 cm2 active surface using a Gamry Reference 3000 potentiostat. This cell was assembled using Fumasep FAP450 membrane (FuMa-Tech GmbH, Germany) as separator, surfacetreated GFD 4.6 graphite felt SIGRACELL (SGL Group, Germany) and FU 4369 graphite bipolar plates (Schunk Kohlenstofftechnik GmbH, Germany) as electrode and current collector materials, respectively. The setup was equipped by additional pair of reference and glassy carbon electrodes, which were positioned in electrolyte tanks to monitor the cathodic, anodic redox, and half-cell potentials (for details of cell design, see ref. [24]). The charging was conducted in galvanostatic–potentiostatic mode: the end-of-charge voltage was set to 1.65 V for galvanostatic step at 75 mA cm 2 and then the cutoff current was 2.5 mA cm 2 at applied voltage of 1.65 V. For discharge, the end-of-discharge voltage was set from 0.8 to 0.2 V depending on the electrolyte sample, i.e., appearance of “power drop” effect. The electrolyte ﬂow rate was 75–80 mL min 1. EIS Measurement: The EIS measurements were conducted at the frequencies in the range from 1 106 to 10 Hz. The direct current was 0 A and the alternative current 0.01 A. The cell resistances were evaluated at high-frequency range from an intercept at Nyquist plots.

Supporting Information Supporting Information is available from the Wiley Online Library or from the author.

Acknowledgements This work was ﬁnancially supported by the German Federal Ministry for Economic Affairs and Energy (BMWi) in the context of the project “DegraBat” (03ET6129B). The responsibility for the content of this publication lies with the authors. Open access funding enabled and organized by Projekt DEAL.

Conﬂict of Interest The authors declare no conﬂict of interest.

Keywords electrolytes, “power drop” effect, thermally induced aging, vanadium redox-ﬂow batteries

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Received: May 15, 2020 Revised: July 16, 2020 Published online: August 17, 2020

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