Ionics

, Volume 20, Issue 7, pp 949–955

Study of the electrochemical performance of VO2+/VO2+ redox couple in sulfamic acid for vanadium redox flow battery

Authors

  • Zhangxing He
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
  • Yaoyi He
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
  • Chen Chen
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
  • Shuai Yang
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
  • Jianlei Liu
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
  • Zhen He
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
    • Key Laboratory of Resources Chemistry of Nonferrous Metals, Ministry of Education, School of Chemistry and Chemical EngineeringCentral South University
Original Paper

DOI: 10.1007/s11581-013-1051-6

Cite this article as:
He, Z., He, Y., Chen, C. et al. Ionics (2014) 20: 949. doi:10.1007/s11581-013-1051-6

Abstract

The present work was performed in order to evaluate sulfamic acid as the supporting electrolyte for VO2+/VO2+ redox couple in vanadium redox flow battery. The oxidation process of VO2+ has similar electrochemical kinetics compared with the reduction process of VO2+. The exchange current density and standard rate constant of VO2+/VO2+ redox reaction on a graphite electrode in sulfamic acid are determined as 7.6 × 10−4 A cm−2 and 7.9 × 10−5 cm s−1, respectively. The energy efficiency of the cell employing sulfamic acid as supporting electrolyte in the positive side can reach 75.87 %, which is adequate for redox flow battery applied in energy storage. The addition of NH4+ to the positive electrolyte can enhance the electrochemical performance of the cell, with larger discharge capacity and energy efficiency. The preliminary exploration shows that the vanadium sulfamate electrolyte is promising for vanadium redox flow battery and is worthy of further study.

Keywords

Vanadium redox flow batterySupporting electrolyteSulfamic acidElectrochemical kinetics

Introduction

With a large scale of consumption of fossil fuels, environmental pollution and limited reserves have led to an increasing use of clean and renewable energy, such as solar energy, wind energy, etc. However, the intermittent varied nature of these clean and renewable resources makes it difficult to integrate these valuable energies into electrical supply grids [1]. One solution to this problem would be to employ large-scale electrical energy storage. There are several energy-storage technologies which have been applied in electrical energy storage, such as flywheels, compressed air, superconducting magnetic energy storage, and flow battery [24]. Among these technologies, redox flow battery systems have attracted much attention, such as Br/Br2 vs. Zn2+/Zn [5], Fe3+/Fe2+ vs. Cr2+/Cr3+ [6, 7], V5+/V4+ vs. V2+/V3+ [6], etc. [710].

Vanadium redox flow battery (VRFB), originally proposed by Skyllas-Kazacos et al., has been considered as one of the most promising energy storage systems for intermittently renewable energy because of its long cycle life, high energy efficiency, and environmental friendship [1114]. A VRFB uses a V2+/V3+ sulfate solution at the negative side and a VO2+/VO2+ sulfate solution at the positive side. Cross-contamination effect can be eliminated between the electrolytes through the ion-exchange membrane due to the same element of active species. A standard voltage of 1.26 V is produced by the VRFB system through the following reactions:
$$ \begin{array}{ll}\mathrm{Cathode}:{\mathrm{VO}}^{2+}+{\mathrm{H}}_2\mathrm{O}-{\mathrm{e}}^{-}\underset{\mathrm{Discharge}}{\overset{\mathrm{Charge}}{\leftrightarrow }}{\mathrm{VO}}_2^{+}+2{\mathrm{H}}^{+}\hfill & {E}^0=1.00\mathrm{V}\hfill \end{array} $$
(1)
$$ \begin{array}{ll}\mathrm{Anode}:{\mathrm{V}}^{3+}+{\mathrm{e}}^{-}\underset{\mathrm{Discharge}}{\overset{\mathrm{Charge}}{\leftrightarrow }}{\mathrm{V}}^{2+}\hfill & {E}^0=-0.26\mathrm{V}\hfill \end{array} $$
(2)
$$ \begin{array}{ll}\mathrm{Cell}:{\mathrm{V}\mathrm{O}}^{2+}+{\mathrm{H}}_2\mathrm{O}+{\mathrm{V}}^{3+}\underset{\mathrm{Discharge}}{\overset{\mathrm{Charge}}{\leftrightarrow }}{\mathrm{V}\mathrm{O}}_2^{+}+2{\mathrm{H}}^{+}+{\mathrm{V}}^{2+}\hfill & {E}^0=1.26\mathrm{V}\hfill \end{array} $$
(3)

Sulfuric acid is a common supporting electrolyte in the VRFB. It was mentioned that trifluoromethanesulfonic acid (CF3SO3H), methanesulfonic acid (CH3SO3H), perchloric acid (HClO4), nitric acid (HNO3), hydrochloric acid (HCl), and sulfamic acid (NH2SO3H) could be chosen as bifunctional acid electrolytes for the redox flow battery [1, 2, 15, 16]. In recent years, other supporting electrolytes such as methanesulfonic acid, hydrochloric acid, and mixed sulfate-chloride electrolyte have been investigated in VRFB [1, 17, 18]. Kim et al. [18] reported that >6 M hydrochloric acid supporting electrolyte for VRFB has better thermal stability than the sulfuric acid supporting electrolyte, capable of dissolving more than 2.3 M vanadium at varied valence states and remained stable at 0–50 °C. The improved thermal stability is attributed to the formation of a vanadium dinuclear [V2O3·4H2O]4+ or a dinuclear-chloro complex (V2O3Cl·3H2O)3+ in the solution over a wide temperature range. Moreover, VRFB with chloride electrolytes demonstrated excellent reversibility and fairly high efficiency. Meanwhile, the thermal stability of the vanadium electrolyte at various valence states and the electrochemical behavior of a VRFB using mixed sulfate-chloride electrolyte were studied by Li et al. [1]. Peng et al. [17] reported the investigation of using CH3SO3H-H2SO4 mixed acid as a supporting electrolyte for VRFB. Compared to that in the sulfuric acid solution system, the redox reaction kinetics of VO2+/VO2+ couple in the CH3SO3H-H2SO4 mixed solution is increasing, with stable cycling performance and higher energy density. However, as a supporting electrolyte, hydrochloric acid has strong corrosivity, and chlorine evolution may occur at the electrode in the anode process due to the inconsistency of the cells. Nitrous oxide evolution also occurs at the surface of the electrode with nitric acid as supporting electrolyte. Trifluoromethanesulfonic acid is the strongest organic acid and has strong hydroscopicity, which may impose some safety problems and strict requirements to equipment.

To the best of our knowledge, the electrochemical performance of VO2+/VO2+ redox couple in the sulfamic acid solution has not been reported. Sulfamic acid represents a borderline between the organic and inorganic acids [1921], which is a nonvolatile, noncorrosive, environmental friendly, low-cost, and commercially available reagent, with strong acidity like sulfuric acid and hydrochloric acid [22]. Single-crystal X-ray and neutron diffraction of sulfamic acid [23] show that sulfamic acid exists in the form of a zwitterion:
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Figa_HTML.gif

Sulfamic acid has been widely used as an efficient and environmental friendly heterogeneous catalyst for acid-catalyzed reactions as an alternative to metal catalyst [24, 25]. It is also used in the electroplating field [2]. Under consideration of the sustainable development, sulfamic acid as a new supporting electrolyte in VRFB is studied. This paper reports the kinetic characteristics of VO2+/VO2+ redox couple in sulfamic acid solution and the charge-discharge performance of a VRFB.

Experimental

Preparation of vanadium sulfamate solution

Electrolyte of vanadium(IV) sulfamate was prepared by electrolytic dissolution of V2O5 (Jishou Huifeng Mining Industry Co., Ltd., China) in NH2SO3H (Shanghai Shanpu Chemical Industry Co., Ltd., China) solution in a two-compartment electrolysis cell [26]. The NH2SO3H solution at an appropriate concentration was employed as the anolyte. The NH2SO3H solution with an appropriate weight of V2O5 was employed as the catholyte. NH2SO3H was employed as anion-matched V(IV) ions and supporting electrolyte. A graphite plate with an area of 49 cm2 was employed as the electrode. Electrolysis was performed with DC Power Supply System (Ming Shing Engineering Co., Hong Kong) providing constant current of 3 A. At the end of electrolysis, the final vanadium concentration was analyzed by redox titration. The vanadium electrolytes in other valence states were prepared by electrolysis with vanadium(IV) sulfamate as the original electrolyte and terminated by controlling the electrolysis time strictly. Vanadium(IV) sulfamate solutions containing NH4+ at different concentrations were prepared by adding NH4HCO3 and NH2SO3H into the vanadium(IV) sulfamate solution. All reagents in the experiment were of analytical reagent grade.

UV-Vis spectrometry

UV-Vis spectrometry of the vanadium electrolyte were measured with a UNIC 3802 UV/Vis spectrophotometer (Shanghai, China) in the range of 400–900 nm using a 1.0-cm quartz cell. The measured electrolytes are 0.04 M VO2+ in 3 M H2SO4 and 1.0 M NH2SO3H, respectively. The reference solutions used were 3.0 M H2SO4 and 1.0 M NH2SO3H.

Viscosity and electrical conductivity tests

A Ubbelohde viscometer was used to measure the viscosity of vanadium(IV) sulfamate solutions containing NH4+ at different concentrations. The measurements were carried out at 293 K, unless otherwise specified. Viscosity can be determined by timing the solution with a specified volume flowing through a capillary tube. The solution viscosity was calculated according to the equation as follows:
$$ \mu =\frac{\pi \rho g{d}^4t}{128V} $$
(4)
where μ is the viscosity, ρ is the solution density, d is the diameter of the capillary tube, t is the average time of the solution flowing through the capillary tube for three times, and V is the solution volume.
The electrical conductivity of electrolytes was determined by AC impedance method using a conductivity electrode (Shanghai REX Instrument Factory, China) connected with CHI660C electrochemical workstation (Shanghai Chenhua Instrument Co., Ltd., China). The electrical conductivity of electrolytes can be figured out with the tested ohmic resistance according to the equation as follows:
$$ \kappa =\frac{L}{ RS}=\frac{1}{0.74R} $$
(5)
where κ is the electrical conductivity, L is the distance between two platinum sheets, S is the surface area of the platinum sheet, and R is the ohmic resistance of the solution. The L/S of the conductivity electrode is 0.74.

Cyclic voltammetry and linear sweep voltammetry

Cyclic voltammetry measurements of the electrolyte with and without NH4+ were carried out from 0.3 to 1.2 V vs. saturated calomel electrode (SCE) at a scan rate range of 5–200 mV s−1. Linear sweep voltammetry measurements of the electrolyte were performed at a scan rate of 1 mV s−1. The cyclic voltammetry and linear sweep voltammetry were performed using the CHI660C electrochemical workstation (Shanghai Chenhua Instrument Co., Ltd., China) with a three-electrode system. A 1-cm2 graphite plate and 4-cm2 platinum sheet were used as working electrode and counter electrode, respectively. SCE along with a double salt bridge full of saturated potassium chloride solution was used as reference electrode. Prior to each measurement, the working electrode was polished with 600- and 1,200-grit SiC paper and then washed with distilled water as described in the literature [27]. The reference electrode was washed with distilled water, and the solution in the salt bridge was replaced before use. Unless otherwise specified, the electrode potential is the potential relative to the SCE electrode.

Charge-discharge test

A single static cell was assembled, which consisted of two pieces of polyacrylonitrile (PAN)-based graphite felts with an area of 9 cm2 (Shenhe Carbon Fiber Materials Co., Ltd.), two current collectors, and a perfluorinated ion-exchange membrane (Best Industrial & Trade Co., Ltd., China). The membrane was treated in 3 % H2O2 solution for 1 h and in 1 M H2SO4 solution for 30 min at 80 °C afterward before use. The graphite felt was oxidized in air at 400 °C for 6 h to enhance electrochemical activity and hydrophilicity [12]. Before the cell was assembled, graphite felts were soaked in the original electrolyte for 24 h at an ambient temperature. The original electrolyte used in the tests includes 0.6 M VO2+ with NH4+ at different concentrations in 1.0 M NH2SO3H as positive electrolyte and 0.6 M V3+ in 3.0 M H2SO4 as negative electrolyte. The galvanostatic charge-discharge test was performed between 0.7 and 1.7 V at a current density of 20 mA cm−2 using CT2001C-10V/2A (Wuhan Land Co., China).

Results and discussion

Viscosity and electrical conductivity of the electrolyte

The viscosity and electrical conductivity of VO2+ electrolytes containing NH4+ at different concentrations are shown in Table 1. It can be known that the viscosity increases with increasing NH4+ concentration, as the intermolecular interactions between ions become stronger at a high concentration. The electrical conductivity also increases with increasing NH4+ concentration, as the addition of concentrated NH4+ increases the number of charged ions in the solution. With the further increase of NH4+ concentration, the degree of increase in electrical conductivity becomes smaller. The electrical conductivity increases by 0.0126 S cm−1 with the addition of NH4+ from 2 to 3 M, while it increases by 0.0074 S cm−1 with the addition of NH4+ from 3 to 4 M. The phenomenon may be attributed to the stronger interactions between ions due to the increased viscosity that make the ions’ migration more difficult.
Table 1

The viscosity and electrical conductivity of vanadium(IV) sulfamate solutions containing NH4+ at different concentrations

Concentration of NH4+ (M)

ρ (g mL−1)

t (s)

η (m Pas)

RΩ (Ω)

κ (S cm−1)

0

1.197

23.78

1.902

17.39

0.0777

0.5

1.218

24.58

2.000

14.20

0.0952

1

1.245

25.90

2.154

11.78

0.1147

2

1.287

29.35

2.524

10.35

0.1306

3

1.326

32.02

2.837

9.44

0.1432

4

1.370

39.94

3.656

8.97

0.1506

V(IV) 1.0 mol L−1, NH2SO3H 1.0 mol L−1

UV-Vis spectra of vanadium(IV) sulfamate

Figure 1 shows the UV-Vis spectra of the VO2+ electrolyte for H2SO4 and NH2SO3H systems. The characteristic absorption peak of UV-Vis for VO2+ appears at about 770 nm [28]. As observed, neither a new absorption peak nor a wavelength shift is observed in the sulfamic acid system compared with the sulfuric acid system. It indicates that sulfamic acid as a new supporting electrolyte does not change the bonding style between V and O and the effective concentration of VO2+.
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig1_HTML.gif
Fig. 1

UV-Vis absorption spectra of the VO2+ electrolyte for H2SO4 and NH2SO3H systems

Linear sweep voltammetry

Figure 2 shows the anodic and cathodic polarization curves on the graphite plate for the electrolyte containing 0.1 M VO2+, 0.1 M VO2+, and 1.0 M NH2SO3H. As observed, the anodic and cathodic polarization current densities are close in the same overpotential. It indicates that the oxidation of VO2+ to VO2+ and the reduction of VO2+ to VO2+ have similar electrochemical kinetics. The current density platform appears in both anodic and cathodic polarization curves when the overpotential η > 0.10 V, which means that the oxidation of V(IV) and the reduction of V(V) are mainly controlled by diffusion. It can be known from Fig. 2 that the diffusion-limited current densities in anode and cathode processes for the electrolyte (0.1 M VO2+, 0.1 M VO2+, and 1.0 M NH2SO3H) at a scan rate of 1 mV s−1 are 1.89 × 10−3 and 1.84 × 10−3 A cm−2, respectively.
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig2_HTML.gif
Fig. 2

Anodic and cathodic polarization curves on the graphite plate. a Anodic polarization curve. b Cathodic polarization curve. Scan rate 1 mV s−1, temperature 293 K

Figure 3 shows that there is a linear relationship between the overpotential and current density at a relatively low overpotential area. The polarization resistance, exchange current density, and electrochemical reaction rate constant can be figured out from the equations as follows:
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig3_HTML.gif
Fig. 3

The linear relationship between the overpotential and current density

$$ \begin{array}{lll}{R}_{\mathrm{ct}}=\frac{\eta }{i}\hfill & {i}_0=\frac{ RT}{ nF{R}_{\mathrm{ct}}}\hfill & {k}_0=\frac{i_0}{ nF{C}_{\mathrm{O}}^{\beta }{C}_{\mathrm{R}}^{\alpha }}\hfill \end{array} $$
(6)
where Rct is the charge transfer resistance, η is the overpotential, i is the current density, R is the universal gas constant, T is the Kelvin temperature, n is the number of electrons transferred in the electrode reaction, F is the Faraday constant, k0 is the rate constant, CO is the solution concentration in the oxidation state, CR is the solution concentration in reduction state, β is the transfer coefficient corresponding to the oxidation reaction, and α is the transfer coefficient corresponding to the reduction reaction [9, 29]. From Fig. 3, Rct is estimated as 33.2 Ω cm2; j0, exchange current density, is 7.6 × 10−4 A cm−2; and the corresponding k0 is 7.9 × 10−5 cm s−1. Compared with the literature [30], k0 is larger than that in the sulfuric acid system in which k0 is 1.5 × 10−5 cm s−1. It indicates that the VO2+/VO2+ redox couple in sulfamic acid has better kinetics than that in the sulfuric acid system, which is due to the influence of the complexation and hydration reaction involving VO2+ in different electrolytes. Sulfuric acid is a dibasic acid, and VO2+ and SO42+ can be integrated into a chain-like molecule (see Fig. 4a) in vanadium sulfate electrolyte. Sulfamic acid is a monatomic acid, and VO2+ can only combine two NH2SO3 to generate VO(NH2SO3)2 (see Fig. 4b) in vanadium sulfamate electrolyte. This structure of VO2+ in vanadium sulfamate electrolyte benefits for the decrease of resistance of mass transfer and charge transfer.
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig4_HTML.gif
Fig. 4

Structure of VO2+ species in H2SO4 (a) and NH2SO3H (b)

Cyclic voltammetry

Figure 5 shows the cyclic voltammograms of the electrolyte with and without NH4+ on the graphite plate at a range of scan rates. The anodic peak corresponds to the oxidation of VO2+, and the cathodic peak corresponds to the reduction of VO2+. The anodic and cathodic peaks appear at 0.7–0.9 and 0.5–0.7 V (vs. SCE), respectively. As for both electrolytes with and without NH4+, the peak current increases with the increase of scan rates. The effect of scan rate on the reduction of VO2+ is similar to that on the oxidation of VO2+, which reflects the approximate diffusion performance of VO2+ and VO2+ in vanadium sulfamate electrolyte. The anodic peak current is close to the cathodic peak current, which is similar to that in vanadium sulfate electrolyte. The formal potential E0 of the VO2+/VO2+ redox couple in 1.0 M NH2SO3H is approximately 0.72 V, which could be figured out from the cyclic voltammograms by taking the average of the anodic and the cathodic peak potential, Epa and Epc [2, 31], i.e.,
$$ {E}^0={\displaystyle \sum_{i=1}^m\frac{E_{\mathrm{pa}}+{E}_{\mathrm{pc}}}{2m}} $$
(7)
where m is the total number of scans. It can be known from the cyclic voltammograms that the redox peak potential interval is larger than 59 mV, suggesting that the electrochemical reaction of VO2+/VO2+ redox couple is not completely reversible. Moreover, the anodic and cathodic peak shapes are symmetrical, especially at low scan rates, which suggests that the redox reaction of VO2+/VO2+ couple is quasi-reversible. Earlier literature showed a quasi-reversible behavior of VO2+/VO2+ redox couple in common supporting electrolytes, such as sulfuric acid [30]. For a quasi-reversible redox couple, the peak current ip is given in the equation as follows [32]:
$$ {i}_{\mathrm{p}}=0.4463{\left(\frac{F^3}{ RT}\right)}^{1/2}{n}^{3/2}A{D}^{1/2}{C}_0{v}^{1/2}K\left(\varLambda, \alpha \right) $$
(8)
where ip is the peak current, F is the Faraday constant, T is the Kelvin temperature, R is the gas constant, n is the electron transfer number, A is the surface area of the electrode, D is the diffusion coefficient, C0 is the concentration of active species, v is the scan rate, and K(Λ, α) is related to the degree of irreversibility [33]. A plot of the anodic peak current as a function of the square root of scan rates (v1/2) for the electrolyte with and without NH4+ is shown in Fig. 6. Generally, K(Λ, α) is dependent on the scan rate, and ip is not proportional to v1/2 for a quasi-reversible reaction. However, in the sulfamic acid system, ip is obviously proportional to v1/2 as shown in Fig. 6. Thereby, K(Λ, α) can be treated as a constant in the range of scan rates [10]. We can assume that α is 0.5 and K(Λ, α) equals to 0.8 under a wide range of Λ from the results of linear sweep voltammetry [33].
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig5_HTML.gif
Fig. 5

Cyclic voltammograms of the electrolyte with and without NH4+ on graphite electrode at a range of scan rates. a Pristine. b With NH4+. Scan rate (mV s−1): a 5, b 10, c 20, d 50, e 100, f 200. VO2+ 0.1 mol L−1, NH2SO3H 1.0 mol L−1, NH4+ 0.2 mol L−1

https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig6_HTML.gif
Fig. 6

Plot of the anodic peak current as a function of the square root of scan rates (v1/2) for the vanadium sulfamate solution with and without NH4+

In the VO2+/VO2+ quasi-reversible reaction, single-electron transfer process, under 293 K, the equation can be simplified as follows [32]:
$$ {i}_p=\left(2.171\times {10}^5\right)A{D}^{1/2}{C}_0{v}^{1/2} $$
(9)

Figure 6 shows the linear relationship between the anodic peak current and the square root of scan rates for the electrolyte with and without NH4+. The diffusion coefficient of VO2+ in sulfamic acid with and without NH4+ can be obtained by the slope value in Fig. 6 and by Eq. 9. The diffusion coefficient of VO2+ in sulfamic acid without NH4+ was estimated as 2.71 × 10−6 cm2 s−1, higher than that with NH4+ (2.10 × 10−6 cm2 s−1), which is due to NH4+ that hinders the diffusion of VO2+. The diffusion coefficient of VO2+ in sulfamic acid is higher than that in sulfuric acid compared with the literature [30]. The better diffusion performance of vanadium ions in sulfamic acid may be related to the low complexation ability of NH2SO3 with V ions. The differences in the diffusion performance may reflect differences in the state of complexation and hydration in each solution.

Charge-discharge test

A single static cell was assembled. The performance of the cell was measured at a constant current density of 20 mA cm−2 between 0.7 and 1.7 V. Figure 7 shows the 20th charge-discharge curves with NH4+ at different concentrations. It can be known that the cell employing the vanadium sulfamate electrolyte with 3.0 M NH4+ as positive electrolyte exhibits better electrochemical performance compared with other cells, with a larger charge-discharge plateau difference and longer charge-discharge time.
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig7_HTML.gif
Fig. 7

The 20th charge-discharge curves of the cells employing the electrolyte with NH4+ at different concentrations at a current density of 20 mA cm−2

The discharge capacity and energy efficiency of the cells employing the electrolyte with NH4+ at different concentrations as a function of cycle number are shown in Fig. 8. It can be known that the discharge capacity presents decrease, than increase trend, which is related to the difference of osmotic pressure of the electrolyte and the transfer of cations between the positive and negative sides on the test. The discharge capacity increases with the increase of the concentration of NH4+. The average discharge capacity of the cells with NH4+ of 1.0 and 3.0 mol L−1 for 50 cycles is 28.91 and 30.98 mAh, respectively, both higher than that without NH4+ in the positive electrolyte (27.14 mAh). The energy efficiency of the cell with sulfamic acid as supporting electrolyte for VO2+/VO2+ redox couple in the positive side can reach 75.87 %, which is adequate for the energy storage system. When the concentration of the added NH4+ is 1.0 and 3.0 mol L−1, the average energy efficiency is 79.54 and 84.57 %, respectively, 3.67 and 8.70 % higher than that without NH4+. Discharge capacity and energy efficiency both increase with the addition of NH4+, which indicates that NH4+ can increase the electrochemical performance of the cell employing the vanadium sulfamate electrolyte.
https://static-content.springer.com/image/art%3A10.1007%2Fs11581-013-1051-6/MediaObjects/11581_2013_1051_Fig8_HTML.gif
Fig. 8

a Discharge capacity and b energy efficiency of the cells employing the electrolyte with NH4+ at different concentrations as a function of cycle number

Conclusions

Environmental friendly and commercially available sulfamic acid as supporting electrolyte for vanadium redox flow battery has been investigated. Compared with the sulfuric acid system, the bonding style between V and O is not changed. The electrical conductivity of the vanadium(IV) sulfamate solution can be improved significantly by adding NH4+. The VO2+/VO2+ couple in sulfamic acid has better kinetics than that in the sulfuric acid system. The exchange current density and standard rate constant of the VO2+/VO2+ redox reaction on the graphite electrode in sulfamic acid are determined as 7.6 × 10−4 A cm−2 and 7.9 × 10−5 cm s−1, respectively. The diffusion coefficients of VO2+ in sulfamic acid with and without NH4+ are 2.10 × 10−6 and 2.71 × 10−6 cm2 s−1, respectively. Vanadium redox flow battery employing sulfamic acid as supporting electrolyte in the positive side was assembled and evaluated. The energy efficiency of the cell employing pure sulfamic acid in the positive electrolyte can reach 75.87 %, and when the concentration of the added NH4+ is 1.0 and 3.0 mol L−1, the energy efficiency increases by 3.66 and 8.69 % compared with that without NH4+, respectively. The preliminary exploration shows that the vanadium sulfamate electrolyte is promising for vanadium redox flow battery and is worthy of further study.

Acknowledgments

This work was financially supported by the Major State Basic Research Development Program of China (973 Program, No. 2010CB227201) and Innovation Project of College Students of Central South University (No. CL12136).

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© Springer-Verlag Berlin Heidelberg 2014