Molybdenum-assisted reduction of VO2+ for lowcost electrolytes of vanadium redox �ow batteriesDeokhyun Hwang 

Soulbrain Co., Ltd.Jong-Wook Ha  ( jwha@soulbrain.co.kr )

Soulbrain Co., Ltd. https://orcid.org/0000-0001-9491-0176Young Soo Park 

Soulbrain Co., Ltd.

Article

Keywords: energy storage systems, all-vanadium redox �ow batteries (VRFB), electrolyte preparation

Posted Date: April 16th, 2021

DOI: https://doi.org/10.21203/rs.3.rs-402509/v1

License: This work is licensed under a Creative Commons Attribution 4.0 International License.  Read Full License

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Molybdenum-assisted reduction of VO2+ for low cost

electrolytes of vanadium redox flow batteries

Deokhyun Hwang, Jong-Wook Ha* & Young Soo Park

Central R&D Center, Soulbrain Co., Ltd.

Seongnam 13486, Gyeonggi-do, Republic of Korea

*email: jwha@soulbrain.co.kr

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Abstract

Efficient and affordable energy storage systems are indispensable to accomplish a successful

energy transition from fossil fuels to renewable sources. Although all-vanadium redox flow

batteries (VRFB) possess many distinctive advantages, much improvement in the process for

electrolyte preparation is needed to overcome low productivity and complexity of the current

electrolysis process. Herein, we demonstrate a simple one-pot process for the preparation of

V3.5+ electrolytes from V2O5 by utilizing hydrazine monohydrate as a residue-free reducing

agent and molybdenum as a homogeneous catalyst accelerating the reduction of VO2+. It is

confirmed that the performance of the electrolytes prepared by the newly developed process is

identical to that by electrolysis in terms of charge-discharge efficiency and capacity up to

current density of 200 mA cm-2. This study can contribute to the wide spread of VRFB by

providing a scalable process suitable for the mass production of V3.5+ electrolyte.

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Introduction

Energy transition from fossil fuels to renewable resources is an important step to alleviate

threats accompanied by climate change because most carbon dioxide, a representative

greenhouse gas, is emitted in the course of electricity production1. Accordingly, solar and wind

power plants have occupied increasing portions of distributed power generation in recent

years2-4. It is clear this trend will be accelerated in the future, but nevertheless, many problems

must be addressed in order for renewable energies to become the ultimate solution for

electricity grid decarbonization5. Above all, the intrinsic intermittence of renewable energies

makes it difficult to balance energy supply and demand, and may even destabilize the power

grid6,7. Therefore, it is necessary to develop efficient and cost-effective energy storage systems

suitable for various purposes8. Despite the active distribution of lithium ion batteries for

stationary applications9, there remains an unyielding demand for development of better

electrochemical energy storage systems.

The redox flow battery is one of the promising technologies for large scale and long duration

stationary energy storage systems with improved safety10. The design and operation of the

redox flow battery are flexible because the power and energy capacities can be decoupled. It

has received a lot of attention because of its excellent scalability and relatively inexpensive

operation and maintenance costs. As a result, various kinds of aqueous and non-aqueous flow

batteries utilizing redox couples based on metal ions and organics have been reported11,12.

Among these, the all-vanadium redox flow battery (VRFB) is of particular interest because of

its reduced occurrence of cross-contamination of electrolytes across the membrane through

application of the same active species to positive and negative electrolytes13,14. Another

attractive point of VRFBs is their excellent sustainability15. Vanadium is an abundant element

more commonly found in earth’s crust than zinc or copper and the vanadium electrolyte is not

consumed during operation, and thus, can be recycled16. Although significant progress toward

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commercialization of VRFBs has been made in the last three decades or so, further

improvement is required, particularly from an economic point of view, to widen their

acceptance17.

Compared with other key components constituting a VRFB stack such as ion exchange

membranes and electrode materials, electrolytes have received relatively little attention18,19. It

is known that the portion of the electrolytes in the installation cost of VRFB is significant, and

its importance increases as the power to energy ratio increases20. In the early stages of VRFB

development, VOSO4 and V2O3 were used as raw materials for electrolytes, but it is now

common to use V2O5 as a starting material, which is much cheaper and commercially available

in large quantities. In addition, the recent trend is to prepare a so-called V3.5+ electrolyte

containing equimolar concentrations of VO2+ and V3+ ions and to use this as both the positive

and negative electrolytes of the system at the same time. However, since V2O5 has insufficient

solubility in aqueous sulfuric acid solution, the preparation of electrolyte must be preceded by

reduction of VO2+ to VO2+ in order to reach a vanadium concentration of 1.5 to 1.8 M. Various

chemical reducing agents can be used in this step. On the other hand, there is no known

chemical reducing agent effective in the reduction of VO2+ to V3+ under mild reaction

conditions. Thus, the manufacture of V3.5+ electrolyte almost entirely relies on electrolysis21.

The serious drawback of this process is that undesirable oxidation of VO2+ occurs

simultaneously at the anode during the reduction of VO2+ at the cathode22,23. An additional

process is required to chemically reduce the surplus product back into VO2+. According to

recent techno-economic studies for VRFB, the manufacturing process cost accounts for 37 -

50% of the electrolyte cost24,25. Therefore, it is obvious that improvement of the manufacturing

process will have a great impact on the cost-effectiveness of the overall VRFB systems. For

this purpose, achieving simplicity of the process is more important than anything else because

costs for labor, initial investment and maintenance for equipment, as well as quality control of

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the manufactured electrolytes are expected to be cut down.

In this study, we develop a one-pot process using a chemical reducing agent to prepare a

V3.5+ electrolyte from V2O5. For a chemical reducing agent to be desirable, it is required not

only to reduce VO2+ and VO2+ to a lower oxidation state, but also to leave no residue in order

to ensure the performance of the electrolyte. Hydrazine monohydrate is known to be oxidized

in an acidic solution as follows26:

N2H5+ → N2 + H5

+ + 4e- (1)

It can be seen that N2 is the only by-product, which is inert and can be easily removed, and

hydrazine monohydrate has an excellent atomic economy by emitting 4 electrons per molecule.

Surprisingly, while some patents have disclosed that hydrazine compounds could be used in

the preparation of VRFB electrolytes27,28, detailed research has not yet been reported. This is

unexpected considering that the hydrazine compound is often used as a reducing agent in the

preparation of vanadium oxide nanoparticles29,30. Moreover, the serendipitous discovery of the

role of molybdenum in the reaction herein makes it possible to prepare V3.5+ electrolytes at an

acceptable reaction rate. Since the vanadium oxidation state (VOS) of the electrolyte can be

controlled by the amount of reducing agent and the reaction rate by the molybdenum

concentration, the process is concise and highly reproducible. The process developed here is

suitable for mass-production of electrolytes with a consistent quality.

Results

One-pot preparation of V3.5+ electrolytes.

It has been reported that the reduction of VO2+ to VO2+ proceeds easily in acidic medium with

various reducing agents such as oxalic acid22,23,31, methanol32 and glycerol33 due to the strong

oxidizing power of VO2+. However, according to our preliminary experiments, large amounts

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of reducing agents in excess of the stoichiometric amount are usually required in order to

complete the reductive dissolution of V2O5. The performance of electrolytes can be degraded

due to the presence of the organic residue in this case. Even though oxalic acid fully converted

V2O5 to VO2+ at a stoichiometric amount, unfortunately, it was not effective in the reduction

of VO2+ to V3+, likely due to the high activation energy of the VO2+/V3+ redox couple34. For

this reason, electrolysis has been regarded as the only way to produce V3.5+ electrolytes, as

mentioned earlier. Recently, it has been reported that the sluggish VO2+ reduction with formic

acid as an organic reducing agent can be promoted using a Pt catalyst35. However, in order to

achieve a continuous catalytic reduction process, two different reducing agents must be used

for the reduction of VO2+ (oxalic acid) and VO2+ (formic acid), respectively. Therefore, a two-

step process is still necessary to produce V3.5+ electrolytes.

According to many previous studies18,19,36-38, VRFB electrolytes with optimum performance

consist of 1.5 ~ 1.8 M vanadium and 2 ~ 3 M free H2SO4 if one considers the balance between

the energy density and electrolyte stability within the operating temperature window. In this

context, we demonstrate the effectiveness of the developed process by considering the

preparation of electrolytes with 1.6 M vanadium and 4.0 M total sulfate. In a typical one-pot

preparation of 1L V3.5+ electrolyte, 0.624 mole of hydrazine monohydrate was poured into the

reactor containing an aqueous dispersion of 0.8 moles of V2O5. Hydrazine monohydrate was

expected to be applicable to the reduction of VO2+ because it has a standard redox potential of

-0.23 V in an acidic solution, which is much lower than that of VO2+/VO2+ (1.0 V vs standard

hydrogen electrode). The color of V2O5 changed from yellow to dark gray immediately with

moderate N2 evolution. This slight exothermic reaction seems to reflect the reduction of VO2+

to VO2+, but due to limited solubility of V2O4 in water, it was not possible to obtain a

homogeneous solution even after prolonged agitation or temperature elevation. Fig. 1a shows

the temperature change inside the reactor measured while adding 4 moles of concentrated

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sulfuric acid over a period of 1 h. Addition of sulfuric acid evoked vigorous N2 evolution and

liberated considerable reaction heat in the first 25 min, while the reaction mixture became

homogeneous. Roughly, the amount of sulfuric acid added up to this point was 1.7 moles,

which is consistent with that required for formation of VOSO4. After that, a slight decrease in

temperature was observed despite the continuous addition of sulfuric acid. UV-Vis spectra

measured at 30 min after start of the addition of sulfuric acid indicated that all vanadium was

converted to VO2+ (Fig. 1b). It is worthy of note that three different V2O5 materials of industrial

grade (99.5% min) were examined in these experiments, and no significant difference was

observed until this point. In order to increase the temperature of the reaction mixture to 100 °C,

external heating of the reactor was initiated. When the addition of sulfuric acid was completed,

the VOS was checked again and the corresponding UV-vis spectra are shown in Fig. 1c. An

interesting result that can be found in this figure is that the existence of V3+ ions is evident in

one of the three electrolytes tested39-41. That is, the absorbance of VO2+ (760 nm) was decreased

and that of V3+ (401 nm) was appreciable in the case of V2O5-C. While maintaining the

temperature inside the reactor at 100 ℃, the changes of VOS were further tracked by UV-Vis

spectroscopy (Fig. 1d). Although there was a significant difference in reaction rates, the

appearance of V3+ by the reduction of VO2+ was obvious in all electrolytes. In one experiment

(V2O5-C), it was even possible to produce the targeted V3.5+ electrolyte after 10 h. To elucidate

the reason for this encouraging observation, the V2O5 raw materials were analyzed by X-ray

fluorescence spectroscopy and a clear relationship between the reaction rate of VO2+ reduction

and the molybdenum content could be found (Supplementary Table 1). As a validation

experiment, 1 mol% of MoO3 (16 mmole) with respect to total vanadium was added to V2O5-

A as a molybdenum source and the electrolyte was prepared with the same procedure described

above. It was again confirmed that the V3.5+ electrolyte could be prepared by a simple one-pot

synthesis with hydrazine monohydrate as the sole reducing agent in the presence of

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molybdenum. In this case, only 2 h was needed to reach a VOS of 3.50 (Fig. 1d and e).

Reaction mechanism of molybdenum-assisted electrolyte preparation.

Separate experiments were conducted to explain the role of molybdenum in the preparation of

V3.5+ electrolytes when hydrazine monohydrate was used as a reducing agent. First, a VO2+

electrolyte (1.6 M total vanadium and 4.0 M total sulfate) was prepared from V2O5-A and a

stoichiometric amount of hydrazine monohydrate (0.4 M). To this, MoO3 was then added, but

no noticeable reduction of VO2+ occurred. Therefore, it could be speculated that the electrons

produced from the oxidation of hydrazine monohydrate were responsible for the reduction of

VO2+, and molybdenum might play a certain role in promoting the reaction rate. Secondly, in

the absence of vanadium ion, the kind of reaction that could proceed between MoO3 and

hydrazine monohydrate in acidic solution was considered. 80 mmole of MoO3 powder was

dispersed in 400 mL of 2.0 M aqueous sulfuric acid solution. To this opaque suspension, 20

mmole of hydrazine monohydrate was injected at 90 ℃. Immediately, the color of the

suspension changed from white to turbid blue, and finally to transparent orange-red

(Supplementary Fig. 1a). Evolution of N2 resulting from the oxidation of hydrazine

monohydrate was also observed. The UV-Vis spectrum revealed a characteristic peak at 295

nm (Supplementary Fig. 1b), which can be attributed to the existence of dimers of Mo5+

(Mo2O42+)42-45. This indicates the occurrence of reductive dissolution of MoO3 by hydrazine

monohydrate. Thirdly and lastly, it was observed that the peak intensity of Mo5+ dimers at 295

nm was consistently decreased upon mixing with VO2+ solution, while that of V3+ at 401 nm

increased at the same time. Based on these observations and the fact that molybdenum has an

ionic form of MoO2+ in a strong acidic solution46, the oxidation-reduction reaction between

hydrazine monohydrate, molybdenum and vanadium in aqueous sulfuric acid solution can be

summarized as follows.

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The oxidation of hydrazine monohydrate:

N2H4 → N2 + 4H+ + 4e- (2a)

The reduction of Mo6+ to Mo5+:

4MoO22+ + 4e- → 2Mo2O4

2+ (2b)

The oxidation and reduction of Mo5+/Mo6+ and VO2+/V3+:

2Mo2O42+ + 4VO2+ + 8H+ → 4MoO2

2+ + 4V3+ +4H2O (2c)

Overall reaction:

N2H4 +4VO2+ + 4H+ → 4V3+ + 4H2O + N2 (2d)

The effect of MoO3 on the one-pot preparation of V3.5+ electrolytes was investigated more

thoroughly, the results of which are represented in Fig. 2. The concentration of MoO3 in the

electrolytes was varied from 1 mM to 16 mM. As can be found in Fig. 2a, the VOS measured

at the end of sulfuric acid addition (or reaction time, t = 0) decreased in proportion to the

amount of MoO3 used in the reaction. This indicated that 15.4% (MoO3 1 mM) to 30.8 %

(MoO3 16 mM) of the VO2+ was already reduced to V3+. The rapid decrease of VOS observed

at the early stage of reaction gradually diminished as the reaction proceeded due to the

depletion of hydrazine monohydrate, and the final VOS of electrolytes ranged from 3.47 to

3.49. There was an obvious relationship between the concentration of MoO3, which depicted

the reaction rate, and the final VOS of the electrolytes. The more MoO3 used in the reaction,

the lower the VOS observed. Since hydrazine monohydrate was used in 4 mol% excess, the

electrolyte should have the VOS of 3.44 if all of the hydrazine monohydrate was utilized in the

reduction of VO2+. This discrepancy in concentration can be interpreted to indicate that

additional side reactions such as thermal decomposition or reaction with oxygen likely occur,

besides the electrochemical oxidation of hydrazine monohydrate47. The final VOS of the

electrolyte can be adjusted more precisely by controlling the amount of excess hydrazine

monohydrate (Supplementary Fig. 2). With regards to the reaction mechanism suggested from

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Eq. 2a to Eq. 2d, the following reaction rate equation can be considered for hydrazine

monohydrate and VO2+.

d[VO2+]dt = k[VO2+][N2H4] (3)

As shown in Fig. 2b, the reduction of VO2+ with hydrazine monohydrate can be assumed to be

a second-order reaction, and the reaction rate constant can be estimated from a linear

relationship between the reaction time and the concentration of the reactant (Fig. 2c). The linear

relationship was satisfactory until the conversion of hydrazine monohydrate reached

approximately 70 to 80%. Fig. 2d shows that the molybdenum concentration in the electrolytes

measured by ICP-OES was linearly proportional to the amount of MoO3 used in the reaction.

From this, we can understand that MoO3 is completely dissolved in the electrolyte and plays a

role as a homogeneous catalyst, which is beneficial to secure the reproducibility of the

electrolyte manufacturing process (Supplementary Fig. 3). Due to the simplicity of the process,

scale-up of electrolyte production can be accomplished without many difficulties

(Supplementary Fig. 4). The effect of molybdenum on the physicochemical properties of the

electrolytes, such as kinematic viscosity and electrical conductivity, was not appreciable

(Supplementary Table 2). In addition, it was confirmed that all electrolytes had the target

composition of 1.6 M vanadium and 4.0 M sulfate.

Electrochemical properties of electrolytes.

We investigated the electrochemical properties of electrolytes containing molybdenum using

cyclic voltammetry (CV) and electrochemical impedance spectroscopy (EIS). There is a lot of

controversy about the effect of impurities in electrolyte48,49. It is not rare to find that an impurity

showing a negative effect in one study led to an improvement in performance in another study.

High purity electrolytes may be desirable to ensure long-term durability of VRFB systems.

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However, if one considers the economics of the entire system, it is recommended to identify

and distinguish between impurities that have a relatively insignificant effect and those that

cause fatal performance degradation50. Accordingly, the influence of molybdenum on the

electrochemical performance of electrolytes should be carefully evaluated.

CV was performed at scan rates ranging from 2 mV s-1 to 30 mV s-1 (Supplementary Fig. 5).

As an example, voltammograms of the positive (VO2+/VO2+) and negative (V2+/V3+)

electrolytes measured at a scan rate of 5 mV s-1 are shown in Fig. 3a and b, respectively. The

results for electrolyte prepared by chemical reduction with oxalic acid and electrolysis (Ox-EL)

were also included for comparison. In the case of positive electrolytes, peak potential

separation (ΔEp = Ec – Ea) and anodic and cathodic peak currents (ia and ic) showed little

difference between each electrolyte when the molybdenum concentration was lower than 4 mM

(Fig. 3a). Increased peak separation and decreased peak currents were observed with increasing

molybdenum. On the other hand, for negative electrolytes, the effect of molybdenum was more

appreciable. As the molybdenum concentration increased from 1 mM to 16 mM, ΔEp increases

from 0.23 V to 0.35 V, and ia and ic decreases from 2.17 mA to 1.40 mA and -2.41 mA to -1.61

mA, respectively, were observed (Fig. 3b). The effect of molybdenum on the electrochemical

performance was further examined by CV at different scan rates. From Fig. 3c and d, an

increase in ΔEp with increasing scan rate can be found, which is consistent with many studies

showing that the redox reaction of vanadium ions on the electrode is a quasi-reversible

reaction51-53. As the scan rate increases, ΔEp increases because a larger overpotential is required

due to the enhanced polarization effect, which suggests that the reaction kinetics come into

competition with mass transfer. In addition, we can find that ia and ic deviated from linearity in

plots of these versus the square root of the scan rate for both positive (Fig. 3e) and negative

(Fig. 3f) electrolytes, especially with increase of the scan rate or amount of molybdenum in the

electrolytes, which provides additional evidence of quasi-reversibility of the redox reaction.

12

According to the aforementioned CV results, molybdenum has a larger effect on the

reversibility of the redox reaction of negative electrolytes than positive electrolytes.

Nevertheless, it is worthwhile to mention that the CV behavior of Ox-EL containing a

negligible amount of molybdenum (1.9 ppm) is almost the same as electrolytes prepared by

molybdenum-catalyzed one-pot synthesis. For negative electrolyte, similar values of ΔEp, ia

and ic with electrolyte containing 16 mM of molybdenum were observed.

Fig. 4 represents the EIS results for positive and negative electrolytes. All Nyquist plots are

constituted by a semicircle part in the high frequency region and a linear part in the low

frequency region, indicating the total impedance is governed by both the diffusion of vanadium

ions and charge transfer at the electrode/electrolyte interface. With regard to the molybdenum

concentration, there was no difference between positive electrolytes at concentrations below 4

mM, but the radius of the semicircle increased slightly at 8 mM, with a more evident increase

at 16 mM (Fig. 4a). However, the slope of the linear line of the low frequency region was

independent of the molybdenum concentration. In contrast, for negative electrolytes, the radius

of the semicircle increased consistently according to the molybdenum concentration (Fig. 4b).

Therefore, charge transfer between electrode and electrolyte may be influenced to some extent

by molybdenum in these cases. Fig. 4c and d show two parameters, Rs and Rct, evaluated by

fitting experimental results to the equivalent circuit consisting of ohmic resistance of bulk

electrolyte (Rs), the charge transfer resistance across electrode/electrolyte (Rct), the constant-

phase-element (CPE) describing the electrochemical double layer, and the Warburg coefficient

(Z) resulting from vanadium ion diffusion54. For both the positive and negative electrolytes, Rs

values were independent of molybdenum concentration. This result is consistent with the

observation of no significant difference in the electrical conductivity of the electrolytes

containing different amounts of molybdenum. On the other hand, the influence of molybdenum

on the Rct was not negligible. For positive electrolyte, Rct increased by 25% and 50% compared

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to Ox-EL when the molybdenum concentration was increased to 8 mM and 16 mM,

respectively (Fig. 4c). A more pronounced effect of molybdenum can be observed in Fig. 4d,

in which estimated Rct values for negative electrolytes are presented. Compared to electrolytes

containing less than 2 mM of molybdenum, a more than 2-fold increase of Rct could be found

when 16 mM of molybdenum was added to the electrolyte. It was also found that the increase

of Rct was linearly proportional to the molybdenum concentration. Other parameters provided

by the equivalent circuit, CPE and Warburg coefficients are listed in Supplementary Table 3.

VRFB single cell performance.

In order to examine the performance of electrolytes, a charge/discharge cycle test was

conducted using a VRFB cell at different current densities of 50, 100, 150 and 200 mA cm-2.

For each current density, five charge/discharge cycles were carried out as shown in

Supplementary Fig. 6, and the charge and discharge voltages of the fifth cycle versus capacity

are represented in Fig. 5 for comparison. The performance of all electrolytes seemed to be

identical in the viewpoint of overpotential and charge/discharge capacities when the current

density was not too high. Supplementary Fig. 7 shows the average coulombic (CE), voltage

(VE), and energy (EE) efficiencies of five charge/discharge cycles at different current densities.

CE usually increases with current density since the charging/discharging occurs at a faster rate,

and thus, the permeation of ions through the separator decreases. CE reflects the decrease in

capacity caused by the crossover of vanadium ions, which is mainly governed by the physical

properties of the ion exchange membrane55, and may not be influenced much by the existence

of molybdenum in electrolytes. Therefore, it can be expected that all electrolytes exhibit similar

performance in terms of CE. VE is a factor that reflects the resistance of the electrode56. As the

current density increased, a decrease in VE due to an increase in overpotential was observed

for all electrolytes. However, when comparing VE at the same current density, it is difficult to

14

find a consistent trend resulting from the molybdenum contained in the electrolyte. All

electrolytes showed a similar discharge energy, as shown in Supplementary Fig. 8. The long-

cycle performance of V3.5+ electrolytes was evaluated by carrying out 200 cycle-

charge/discharge experiments at the current density of 150 mA cm-2. There was a slight

variation in EE in the early stage of charge/discharge cycles, but eventually, all electrolytes

revealed the same efficiencies (Fig. 5e). In addition, the discharge capacity retention, which is

defined as the percentage of discharge capacity of the Nth cycle compared to that of first cycle,

showed similar tendency for all electrolytes (Fig. 5f). Accordingly, it can be concluded that

molybdenum provides little effect on the quality of the electrolyte up to 16 mM, the amount

considered in this study.

Recently, it has been reported that MoO3 deposited on carbon electrodes can act as an

electrocatalyst enhancing the reversibility of VO2+/VO2+ and V2+/V3+ redox couples57,58. This

was responsible for the improvement of voltage efficiency of the VRFB cell. In one study57, it

was also shown that addition of 5 mM of Na2MoO4 directly to the electrolytes modified the

surface of the electrodes during cycling and produced similar effects on VRFB cell

performance as MoO3 deposited on electrodes. In our experiments, the deposition of

molybdenum compounds after cycling was not observed, and the dissolved molybdenum

appeared to slightly worsen the reversibility of redox couples, though not seriously. Therefore,

the enhancement of VRFB cell performance via electrode modification with MoO3 particles

observed in the literature may be attributed to the increased surface wettability of the electrodes

more than to electrocatalytic effects, for which further studies are required to understand this

controversial phenomenon.

Discussion

In this study, we have reported the simplest process ever described for the preparation of high

15

performance V3.5+ electrolyte. It was found that hydrazine monohydrate is quite effective not

only at the reductive dissolution of V2O5, but also at the reduction of VO2+ to V3+ with the

assistance of molybdenum. Consequently, it is possible to prepare V3.5+ electrolyte by a one-

step process without using a complex electrolysis process. The molybdenum concentration in

the electrolytes is coincident with the amount of MoO3 used as a catalyst. Molybdenum slightly

deteriorates the reversibility of redox reaction of the negative electrolyte by increasing charge

transfer resistance, but this seems not to be a critical demerit according to the charge/discharge

test conducted using a VRFB single cell. Conversely, the electrochemical performance of

positive electrolytes was found to be less sensitive to the molybdenum concentration.

When considering the mass production of V3.5+ electrolytes, the current one-pot process is

attractive because only one jacketed reaction vessel equipped with an agitator and a condenser

is necessary, which is common equipment that can be found in ordinary chemical plants.

Moreover, no consumable materials are necessary in the process developed herein. In contrast,

electrolysis requires regular replacement of electrode materials, and catalytic processes

consume expensive Pt catalyst35. Impurities in V2O5 depend on the source of the vanadium and

the manufacturing process. At present, considerable amounts of V2O5 are recycled from spent

hydrodesulfurization catalysts59. In this case, 50 ~ 150 ppm of molybdenum is usually

contained in V2O5 raw materials. Therefore, no additional MoO3 is actually required in order

to prepare V3.5+ electrolyte from V2O5. All aforementioned aspects are favorable for low-cost

V3.5+ electrolyte production (Supplementary Note 1).

Methods

Preparation of electrolytes.

A series of V3.5+ electrolytes were prepared in the presence of MoO3 as a catalyst. Total

vanadium and sulfate concentrations were 1.6 M and 4.0 M, respectively. Deionized water (760

16

g), V2O5 (0.8 mole, 145.6 g) and the appropriated amount of MoO3 (Easchem, 99.95%) were

added to a 2 L round flask equipped with a mechanical agitator, condenser, and thermocouple

connected to the heating mantle. To this suspension, hydrazine monohydrate (0.624 mole,

Samchum Chem, 80% aqueous solution) was added while agitating the reaction mixture.

Finally, concentrated sulfuric acid (4.0 mole, Daejung Chem, 98%) was pumped for 1 h by a

peristaltic pump. During the first 30 min, a severe exothermic reaction proceeded and vigorous

gas evolution was observed. After that point, the reaction mixture was heated externally and

the vanadium oxidation state was measured regularly using a UV-Vis spectrophotometer

(Thermo Fisher Scientific BioMate 160). After cooling, the electrolytes were filtered using a

0.45 μm hydrophilic PTFE membrane filter. For the purpose of comparison, one electrolyte

with the same composition was prepared by a conventional process. V2O5 (0.8 mole) and oxalic

acid dihydrate (0.8 mole, Daejung Chem) were added to deionized water and then sulfuric acid

was added. The reaction mixture was mechanically stirred for 4 h at 90 ℃. After filtration, the

resulting VO2+ solution was electrolyzed in a flow cell. Each half-cell reservoir was filled with

110 mL of the VO2+ electrolyte and charged galvanostatically at a current density of 100 mA

cm-2 until the cell potential reached the value of 1.9 V, in order to produce V3+ in the negative

compartment. At the end of the electrolysis, V3+ electrolyte was mixed with the same volume

of VO2+ electrolyte to obtain a working V3.5+ electrolyte, which is referred to as Ox-EL. The

molybdenum concentration in the electrolytes was determined by ICP-OES (Agilent 5100).

Electrochemical analysis of electrolytes.

The electrochemical performance of electrolytes was evaluated using an electrochemical

workstation (PARSTAT MC PMC-2000A, AMETEK Princeton Allied Research). A standard

three-electrode cell configuration was used employing a platinum counter electrode and a

saturated Ag/AgCl reference electrode. A glassy carbon working electrode with 5 mm diameter

17

was polished manually with 800-grit SiC sandpaper. V2+ and VO2+ electrolytes were obtained

by reducing or oxidizing the V3.5+ electrolytes using a VRFB cell. The positive (VO2+/VO2+)

and negative (V2+/V3+) electrolytes were prepared by mixing VO2+ and V2+ electrolytes at a

volume ratio of 5:1 or 1:5, respectively. Cyclic voltammograms were obtained at different scan

rates. At least three scans were carried out for each experiment and the last scan was taken.

Electrochemical impedance spectroscopy measurements were carried out at open-circuit

potential with an a.c. amplitude of 10 mV from 20 kHz to 0.2 Hz, and data were analyzed using

a Randles equivalent circuit so as to calculate charge transfer resistances.

VRFB single cell performance.

A proton exchange membrane, Nafion 117 (10 cm ⅹ 10 cm), was soaked in deionized water

before experiments for at least 24 h60. Carbon felts (GFD4.6, SGL Carbon Group) were used

for positive and negative electrodes. The active area was 49 cm2. The cell was connected to

two electrolyte reservoirs containing 80 mL of the V3.5+ electrolytes for both the positive and

negative sides. The electrolyte was circulated through the cell via a peristaltic pump with two

heads at a constant flow rate of 80 mL min-1. The electrolytes were pre-charged

galvanostatically at a constant current density of 100 mA cm-2 up to 1.6 V and then discharged

to 0.8 V. The charge-discharge experiments were performed at current densities of 50, 100,

150 and 200 mA cm-2 within a potential window of 0.8 ~ 1.6 V at room temperature using a

test station (WBCS3000, WonATech).

Data availability

The authors declare that data supporting the findings of this study are available within the

article and its Supplementary Information and also are available from the corresponding author

upon reasonable request.

18

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Author contributions

D.H., J.-W.H. and Y.S.P. conceived the idea of the project. Y.S.P. supervised the project. D.H.

and J.-W.H. designed and performed the experiments. All authors discussed and analyzed the

data. J.-W.H. and Y.S.P. wrote the manuscript.

Competing interests

The authors declare no competing interests.

24

Fig. 1 Electrolyte preparation using hydrazine monohydrate as a chemical reducing agent. a

Temperature profiles during the addition of H2SO4. b, c UV-Vis spectra measured after addition of half

of (b) and the whole amount of (c) H2SO4. d VOS of electrolytes prepared from different V2O5 raw

materials with or without MoO3 as a function of reaction time. e UV-Vis spectra of the electrolyte

prepared from V2O5-A and 16 mM of MoO3 at different reaction times at 100 °C.

25

Fig. 2 One-pot preparation of V3.5+ electrolytes. a Vanadium oxidation state versus reaction time with

different amounts of MoO3. b Second-order kinetic plots for the reduction of VO2+ (a = [VO2+]0 and b =

[N2H4]0). c Second-order reaction rate constants versus the amount of MoO3 added. d Molybdenum

concentration in the electrolytes measured by ICP-OES.

26

Fig. 3 Cyclic voltammograms (CV) of electrolytes. a, b CVs measured at a scan rate of 5 mV s-1 for

V4.5+ (a) and V2.5+ (b) electrolytes. c, d Peak potential separation (ΔEp) versus scan rate for V4.5+ (c) and

V2.5+ (d) electrolytes. e, f Anodic (ia) and cathodic (ic) peak currents versus the square root of scan rates

for V4.5+ (e) and V2.5+ (f) electrolytes.

27

Fig. 4 Electrochemical impedance spectroscopy (EIS) of electrolytes. a, b Nyquist plots of V4.5+ (a) and

V2.5+ (b) electrolytes. c, d Ohmic resistance (Rs) and charge transfer resistance (Rct) obtained from

fitting EIS results for V4.5+ (c) and V2.5+ (d) electrolytes.

28

Fig. 5 VRFB single cell performance. a-d Charge/discharge curves at current densities of 50 mA cm-2

(a), 100 mA cm-2 (b), 150 mA cm-2 (c) and 200 mA cm-2 (d). e, f Cycling stability of electrolytes in terms

of coulombic and energy efficiencies (e) and discharge capacity retention (f) at 150 mA cm-2.

Figures

Figure 1

Electrolyte preparation using hydrazine monohydrate as a chemical reducing agent. a Temperaturepro�les during the addition of H2SO4. b, c UV-Vis spectra measured after addition of half of (b) and thewhole amount of (c) H2SO4. d VOS of electrolytes prepared from different V2O5 raw materials with orwithout MoO3 as a function of reaction time. e UV-Vis spectra of the electrolyte prepared from V2O5-Aand 16 mM of MoO3 at different reaction times at 100 °C.

Figure 2

One-pot preparation of V3.5+ electrolytes. a Vanadium oxidation state versus reaction time with differentamounts of MoO3. b Second-order kinetic plots for the reduction of VO2+ (a = [VO2+]0 and b = [N2H4]0).c Second-order reaction rate constants versus the amount of MoO3 added. d Molybdenum concentrationin the electrolytes measured by ICP-OES.

Figure 3

Cyclic voltammograms (CV) of electrolytes. a, b CVs measured at a scan rate of 5 mV s-1 for V4.5+ (a)and V2.5+ (b) electrolytes. c, d Peak potential separation (ΔEp) versus scan rate for V4.5+ (c) and V2.5+(d) electrolytes. e, f Anodic (ia) and cathodic (ic) peak currents versus the square root of scan rates forV4.5+ (e) and V2.5+ (f) electrolytes.

Figure 4

Electrochemical impedance spectroscopy (EIS) of electrolytes. a, b Nyquist plots of V4.5+ (a) and V2.5+(b) electrolytes. c, d Ohmic resistance (Rs) and charge transfer resistance (Rct) obtained from �tting EISresults for V4.5+ (c) and V2.5+ (d) electrolytes.

Figure 5

VRFB single cell performance. a-d Charge/discharge curves at current densities of 50 mA cm-2 (a), 100mA cm-2 (b), 150 mA cm-2 (c) and 200 mA cm-2 (d). e, f Cycling stability of electrolytes in terms ofcoulombic and energy e�ciencies (e) and discharge capacity retention (f) at 150 mA cm-2.

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