Electrochemical hydrogen production system from ammonia borane in methanol solution

Electrochimica Acta 82 (2012) 392–396

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Electrochemical hydrogen production system from ammonia borane in methanol solution Hiroshi Inoue ∗,1 , Toshiki Yamazaki, Takayuki Kitamura, Motoki Shimada, Masanobu Chiku, Eiji Higuchi 1 Department of Applied Chemistry, Graduate School of Engineering, Osaka Prefecture University, Sakai, Osaka 599-8531, Japan

a r t i c l e

i n f o

Article history: Received 1 December 2011 Received in revised form 29 May 2012 Accepted 29 May 2012 Available online 5 June 2012 Keywords: Ammonia borane Methanolysis Hydrogen production Electrolysis Electrochemical oxidation

a b s t r a c t An electrochemical hydrogen production system based on ammonia borane was constructed in a 2 M methanol solution of lithium perchlorate. Cyclic voltammograms of ammonia borane at the Pt anode demonstrated the production of hydrogen by the electrochemical oxidation of ammonia borane in methanol solution for the first time. Hydrogen and NH4 B(OCH3 )4 were identified as products of the potentiostatic electrolyses by mass spectroscopy, 11 B NMR spectroscopy, and gas chromatography. Hydrogen was produced at both electrodes and the current efficiency for hydrogen production at the Pt anode was greater than 100%. The electrochemical data was used to propose a mechanism for the electrochemical production of hydrogen from ammonia borane. © 2012 Elsevier Ltd. All rights reserved.

1. Introduction Ammonia borane, NH3 BH3 , can desorb a maximum of three moles of hydrogen, and has a high gravimetric and volumetric hydrogen density (19.6 wt% and 0.145 kg(H2 ) L−1 , respectively). One and two moles of hydrogen (6.5 and 13 wt%) are produced by the thermal dehydrogenation of ammonia borane at 107–117 ◦ C and 150 ◦ C, respectively [1,2]. Such production has not been achieved using any other light-weight high-density hydrogen storage material, such as lithium borohydride. In addition, ammonia borane is suitable for secure hydrogen transportation, because it is a stable, nontoxic, eco-friendly material [1,2]. Thus, ammonia borane is promising as a next-generation high-density hydrogen storage material. However, the recovery of ammonia borane through hydrogen absorption is an important consideration. Solvolysis, such as hydrolysis and methanolysis, is also a promising method for producing hydrogen from ammonia borane. Solvolysis reactions can proceed under appropriate catalysis even at room temperature, although the gravimetric hydrogen density is decreased by the weight of solvents; by 7.1 wt% for hydrolysis and by 4.7 wt% for methanolysis [1–12]. Hydrolysis is the most widely used type of solvolysis, and there have been many studies of catalysts for hydrolysis at room temperature [3–7]. The purity of the

∗ Corresponding author. Tel.: +81 72 254 9283; fax: +81 72 254 9283. E-mail address: [email protected] (H. Inoue). 1 ISE member. 0013-4686/$ – see front matter © 2012 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.electacta.2012.05.091

hydrogen was high, but ammonia was also produced in concentrated aqueous ammonia borane solutions. In addition, borate, a hydrolysis product, was too stable to regenerate ammonia borane. However, methanolysis produced high-purity hydrogen irrespective of the ammonia borane concentration and ammonia borane was regenerated from the methanolysis product [8], which is different from the hydrolysis product. Electrolysis consists of an oxidation reaction at the anode and a reduction reaction at the cathode, which generates at least two products. For example, during brine electrolysis, chlorine is produced at the anode, and sodium hydroxide and hydrogen are produced at the cathode. During water electrolysis hydrogen and oxygen are produced at the cathode and anode, respectively. If hydrogen is the intended product, the oxygen produced at the anode is wasted. Thus, an electrolysis system where hydrogen is produced at both the anode and the cathode is desirable. Recently, the electrochemical oxidation of ammonia borane [13] and dimethylamine borane [14] has been investigated in order to elucidate the reaction mechanism of borane-based electroless plating, where hydrogen production was a side reaction. In addition, the electrochemical characterization of ammonia borane in aqueous solutions was also carried out for direct ammonia borane fuel cells [15–17]. However, the electrochemical production of hydrogen from ammonia borane in methanol has not yet been explored. In this study, we report the use of the electrochemical methanolysis of ammonia borane as an electrochemical hydrogen production system, and demonstrate that hydrogen was produced at both electrodes.

H. Inoue et al. / Electrochimica Acta 82 (2012) 392–396

2. Experimental

3nNH3 BH3

.

3. Results and discussion 3.1. Cyclic voltammograms of ammonia borane in methanol electrolyte solution Fig. 1 shows the CVs of 1 mM ammonia borane in the N2 saturated methanol electrolyte solution at different sweep rates. The oxidation of methanol was not observed in this potential range. The oxidation of ammonia borane began around −0.25 V; an oxidation peak was observed in the forward sweep, and a reduction peak was observed in the reverse sweep. In addition, the plot of the peak oxidation current (ip ) against the square root of the sweep rate (v) was linear with an intercept at the origin (Fig. 2), indicating the oxidation reaction was diffusion-controlled. In the cyclic voltammetric response to a reversible heterogeneous electron transfer process, the Randles–Sevcik equation (Eq. (2)) usually holds between ip and v1/2 [18]:

 nFvD 1/2 RT

.

(2)

Here, n is the number of electrons, F is the Faraday constant [C mol−1 ], A is the area [cm2 ], c is the concentration [mol cm−3 ], D is the diffusion coefficient [cm2 s−1 ], R is the gas constant [J K−1 mol−1 ], and T is the temperature [K]. If n = 1, the diffusion coefficient of ammonia borane in the methanol electrolyte solution was 8.2 × 10−5 cm2 s−1 , which is larger than that reported value for 1 M NaOH aqueous solution, measured by chronoamperometry with a microelectrode (1.68 × 10−5 cm2 s−1 ) [13]. Usually, the diffusion coefficient depends on the solvent. For example, the

0.4 0.2 0 -0.2

-0.4

-0.3

-0.2

-0.1

0

0.1

0.2

Potential vs. (Ag/AgCl) / V

Current density / mA cm-2

2.5 — 500 mV s-1 — 200 mV s-1 — 100 mV s-1

2.0 1.5 1.0 0.5 0 -0.5

-

-1.0 -0.5

(1)

Here, nH2 is the amount of hydrogen produced [mmol], and nNH3 BH3 is the amount of ammonia borane starting material [mmol]. The conversion of 100% means a maximum of three moles of hydrogen is produced from a mole of ammonia borane [1,2,8–12].

— 50 mV s-1 — 20 mV s-1 — 10 mV s-1 — 5 mV s-1

-0.4 -0.5

-0.4

-0.3

-0.2

-0.1

0

0.1

0.2

Potential vs. (Ag/AgCl) / V Fig. 1. Cyclic voltammograms of 1 mM ammonia borane in the N2 -saturated methanol electrolyte solution at different sweep rates.

diffusion coefficient of ferrocene is 2.17 × 10−5 cm2 s−1 in acetonitrile containing 0.2 M tetrabutylammonium perchlorate [19] and 7 × 10−6 cm2 s−1 in aqueous solution containing 0.2 M lithium sulfate [20]. This difference in diffusion coefficients can be attributed to differences in viscosity. At sweep rates of 100 mV s−1 or more, additional redox peaks were observed around 0.1 V (Fig. 1(b)). At faster sweep rates, the second redox couple was observed. The following mechanisms were proposed to explain the results. Ammonia borane may be

Oxidation peak current / mA

Conversion [%] =

100nH2

Current density / mA cm-2

0.6

Ammonia borane (97%, Sigma–Aldrich) was used as purchased. A two-compartment cell divided by a glass filter was used. A Pt disk anode (5 mm diameter) was used for cyclic voltammetry and a Pt plate (16 mm × 10 mm) was used for potentiostatic electrolysis, whereas a Ni plate (40 mm × 20 mm) cathode was used for all electrochemical measurements. The reference electrode was a laboratory-made Ag/AgCl/sat. LiCl electrode. The electrolyte was a 2 M solution of lithium perchlorate in dry methanol (Wako Pure Chemical Co.). The lithium perchlorate was dried in a vacuum at 120 ◦ C for 3 h before use. Various concentrations of ammonia borane were dissolved in the methanol electrolyte solution. The redox potential of 1 mM ferrocene in the methanol electrolyte solution measured by cyclic voltammetry against the reference electrode was 0.400 V. The cyclic voltammograms (CVs) were recorded at various sweep rates in nitrogen-saturated methanol electrolyte solutions, which contained various concentrations of ammonia borane. The methanol electrolyte solution which contained 10 mM ammonia borane was used in the potentiostatic electrolyses at various potentials. All the electrochemical measurements were performed at 30 ◦ C. The hydrogen produced at both electrodes was qualitatively identified by gas chromatography and mass spectroscopy, and the amount of hydrogen produced in the potentiostatic electrolyses was determined by a volumetric method [8]. The product in the reaction bath was identified by 11 B NMR. The conversion of ammonia borane to hydrogen is defined by the following equation:

ip = 0.446nFAc

393

0.4

0.3

0.2

0.1

0 0

0.2

0.4

0.6

0.8

1.0

v1/2 / V1/2 s-1/2 Fig. 2. Relationship between the oxidation peak current and the square root of the sweep rate.

H. Inoue et al. / Electrochimica Acta 82 (2012) 392–396

8 6 4

— 20 mV s-1 — 10 mV s-1 — 5 mV s-1

2 0

Intensity

Current density / mA cm-2

394

-2 -4 -6 -0.6

-0.4

-0.2

0

0.2

0.4

NH3BH3

After electrolysis

Potential vs. (Ag/AgCl) / V 50

Current density / mA cm-2

15

30

40

20

10

10

— 200 mV — 100 mV s-1 — 50 mV s-1

0

-20

-30

-40 -50

current was much smaller than the second, which is different from the CVs of 1 mM ammonia borane (see Fig. 1). The methanolysis of ammonia borane proceeds at room temperature with various catalysts including Pt [9–12]. The mechanism of methanolysis is not yet known, although in the mechanism of electroless plating with borohydride [21], the decomposition of ammonia borane probably occurs at the first step (Eq. (7)).

-5 -10 -0.4

-10

Fig. 4. 11 B NMR spectra of ammonia borane and the liquid phase product after the potentiostatic electrolysis of a 0.5 M methanol electrolyte solution, which contained 100 mM ammonia borane, for 35 h at 0.5 V vs. Ag/AgCl.

5

-15 -0.6

0

(ppm)

s-1

-0.2

0

0.2

0.4

•

NH3 BH3 → NH3 BH2 (III) + H

Potential vs. (Ag/AgCl) / V Fig. 3. Cyclic voltammograms of 10 mM ammonia borane in the N2 -saturated methanol electrolyte solution at different sweep rates.

converted to a radical cation (I) by the first electrochemical oxidation (Eq. (3)). •+

NH3 BH3  NH3 B H3 (I) + e−

(3)

At slower sweep rates, atomic hydrogen would leave the radical cation (I) to produce cation (II) (Eq. (4)). +

Radical cation (I) → NH3 BH2 (II) + H

•

(5)

In the reverse scan, the reverse reactions of Eqs. (3) and (5) would proceed at higher sweep rates, while cation (II) would be reduced to form ammonia borane at lower sweep rates (Eq. (6)). Cation (II) + H+ + 2e− → NH3 BH3

(7)

If so, the first oxidation (Eq. (3)) would be suppressed, and the oxidation of radical (III) (Eq. (8)) would proceed instead. Radical (III) → Cation (II) + e−

(8)

Therefore the oxidation peak at 0.05 V in Fig. 3 can be assigned to the oxidation of radical (III) (Eq. (8)), and the reverse reaction can be represented by Eq. (6). Whether the decomposition of ammonia borane occurs before its electrochemical oxidation depends on the concentration of ammonia borane. However, for any route, the electrochemical oxidation of ammonia borane accelerates hydrogen production.

(4)

In contrast, at faster sweep rates, radical cation (I) would be oxidized further before the release of atomic hydrogen, leading to the formation of cation (II) (Eq. (5)), which explains the appearance of the second oxidation peak. Radical cation (I)  Cation (II) + H+ + e−

•

(6)

This explains why the reduction peak current is higher than the oxidation peak current at lower sweep rates. Fig. 3 shows the CVs of 10 mM ammonia borane in the N2 saturated methanol electrolyte solution at different sweep rates. At lower sweep rates, a small shoulder and a peak from the oxidation of ammonia borane were observed in the forward scan around −0.1 V and 0.05 V, respectively. These peaks shifted in the positive direction as the sweep rate increased. However, the first oxidation

3.2. Potentiostatic electrolysis of ammonia borane It has previously been reported that 3 mole of hydrogen and 1 mole of NH4 B(OCH3 )4 per 1 mole of ammonia borane were produced in the methanolysis of ammonia borane [1,2,8–12]. The products of the reaction were identified by mass spectroscopy, 11 B NMR, and gas chromatography. Fig. 4 shows the 11 B NMR spectra of ammonia borane and the liquid phase product after potentiostatic electrolysis of the 0.5 M methanol electrolyte solution, which contained 100 mM ammonia borane, for 35 h at 0.5 V vs. Ag/AgCl. The single peak at −23.4 ppm for ammonia borane was assigned to BH3 [8]. After the electrolysis, a new singlet peak appeared at −8.6 ppm in the 11 B NMR spectrum, which was assigned to the tetramethoxyborate [B(OCH3 )4 ]− anion [8,10]. Mass spectroscopy and gas chromatography showed that hydrogen was the sole gas phase product. Figs. 5–7 show the amount of hydrogen produced and current efficiency for hydrogen production against the reaction time for the potentiostatic electrolyses of 10 mM ammonia borane at 0.1, 0.5, and −0.1 V vs. Ag/AgCl, respectively. Hydrogen was

60

0.8

Anode

0.6

40

0.4 20

Cathode

0.2

0

0 0

2

4

6

0.20

395

(a)

0.15

0.10 5

Cathode

0.05

0

0 0

1

Time / h

(b)

500

Anode

300

200

Cathode

100

2

4

(b) Anode

Cathode 100 0 0

1

2

3

4

5

Time / h Fig. 7. Time courses for: (a) the amount of hydrogen produced and (b) the current efficiency of hydrogen production at both electrodes for the potentiostatic electrolyses of ammonia borane at −0.1 V vs. Ag/AgCl.

100

(a) Anode

75

1.0 50 0.5 Cathode

25

0

Conversion (%)

Amount of produced H2 / mmol

5

200

6

Fig. 5. Time courses for: (a) the amount of hydrogen produced and (b) the current efficiency of hydrogen production at both electrodes for the potentiostatic electrolyses of ammonia borane at 0.1 V vs. Ag/AgCl.

0 0

4

300

Time / h

1.5

3

400

0 0

2

Time / h Current efficiency (%)

Current efficiency (%)

400

10

Anode

Conversion (%)

(a)

Amount of produced H2 / mmol

1.0

Conversion (%)

Amount of produced H2 / mmol

H. Inoue et al. / Electrochimica Acta 82 (2012) 392–396

2

4

6

produced at both electrodes irrespective of the applied potential; although the amount of hydrogen produced at the Pt anode was much greater than that at the Ni cathode. In addition, the hydrogen production rate and the amount of hydrogen produced at the Pt cathode increased as the applied potential or the overpotential was increased (Fig. 8). Without electrolysis, the hydrogen was produced by the methanolysis of ammonia borane at both electrodes, although the hydrogen production rate was very slow because the surface area of both electrodes was small (Fig. 8). Thus, hydrogen evolution from ammonia borane at the anode can be electrochemically accelerated and Pt is an active electrocatalyst.

8

(b)

300

Anode

200

Cathode

100

0

0

2

4

6

8

Time / h Fig. 6. Time courses for: (a) the amount of hydrogen produced and (b) the current efficiency of hydrogen production at both electrodes for the potentiostatic electrolyses of ammonia borane at 0.5 V vs. Ag/AgCl.

1.5

100

0.5 V

75

1.0 50

0.1 V 0.5

25

Conversion (%)

Current efficiency (%)

400

Amount of produced H2 / mmol

Time / h

-0.1 V No electrolysis 0

0 0

2

4

6

8

Time / h Fig. 8. Time courses for amount of hydrogen produced at the Pt anode in the potentiostatic electrolyses of ammonia borane at 0.1, 0.5, and −0.1 V vs. Ag/AgCl and without electrolysis.

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At the Ni cathode, hydrogen is produced by the following twoelectron reduction reaction (Eq. (9)): 2CH3 OH + 2e− → 2CH3 O− + H2

(9)

The current efficiency for hydrogen evolution at each potential was equal to or greater than 60%, as shown in Figs. 5–7(b). The current efficiency at the Pt anode was evaluated by assuming that two electrons were consumed during hydrogen evolution. At the Pt anode, the current efficiency was greater than 100%, and was 3- or 4-fold greater than that at the Ni cathode. This suggests that some hydrogen-producing chemical reactions occur during the electrochemical oxidation of ammonia borane. The methoxide produced at the cathode may also have caused the unusually high hydrogen production at the anode. The methanolysis of the 20 mM ammonia borane in the presence of 20 mM sodium methoxide was carried out to investigate the contribution of methoxide. The hydrogen production rate increased by 20% upon addition of sodium methoxide, indicating that methoxide had a small catalytic effect on the methanolysis. During the electrolyses at 0.1 and 0.5 V, cation (II), which was formed through various routes (Eqs. (3) and (5) and/or Eqs. (7) and (8)) could react with methanol to form ammonia monomethoxyborane (IV) (Eq. (10)). Cation(II) + CH3 OH → NH3 BH2 (OCH3 )(IV) + H+

(10)

Compound (IV) could then consecutively react with methanol at the Pt anode to produce hydrogen. NH3 BH2 (OCH3 )(IV) + CH3 OH → NH3 BH(OCH3 )2 + H2

(11)

NH3 BH(OCH3 )2 + CH3 OH → NH3 B(OCH3 )3 + H2

(12)

NH3 B(OCH3 )3 + CH3 OH → NH4 B(OCH3 )4

(13) 11 B

Only NH4 B(OCH3 )4 was detected in the NMR spectrum, suggesting that Eqs. (11)–(13) proceed quickly. These reactions electrochemically produce cation (II) or compound (IV), which means the current efficiency for hydrogen production at the Pt anode exceeds 100%. During electrolysis at −0.1 V, hydrogen was produced at a slower rate than at 0.1 and 0.5 V. Atomic hydrogen produced by cation (II) in Eq. (4) could recombine to give hydrogen (Eq. (14)). H• + H• → H2

(14)

A detailed investigation of our proposed mechanism of the electrochemical methanolysis is required, particularly the identification of intermediates. In addition, the investigation of anode materials for improving the hydrogen production rate is also underway. 4. Conclusions An electrochemical hydrogen production system based on ammonia borane was constructed in a 2 M methanol solution of

LiClO4 . Cyclic voltammograms and potentiostatic electrolyses of ammonia borane in a 2 M methanol solution of LiClO4 at the Pt anode demonstrated the production of hydrogen by the electrochemical oxidation of ammonia borane in methanol for the first time. Hydrogen and NH4 B(OCH3 )4 were identified as products of the potentiostatic electrolyses by mass spectroscopy, 11 B NMR spectroscopy, and gas chromatography. In particular, hydrogen was produced at both electrodes, and the current density for hydrogen production at the Pt anode was greater than 100%. A mechanism for the electrochemical production of hydrogen from ammonia borane was proposed based on the electrochemical data. Acknowledgments This work was partly supported by a Grant-in-Aid for Scientific Research on Challenging Exploratory Research (No. 22656157) from the Ministry of Education, Culture, Sports, Science and Technology of Japan. References [1] C.W. Hamilton, R.T. Baker, A. Staubitz, I. Manners, Chemical Society Reviews 38 (2009) 279. [2] P. Wang, X.-D. Kang, Dalton Transactions 40 (2008) 5400. [3] H.-L. Jiang, Q. Xu, Catalysis Today 170 (2011) 56. [4] O. Metin, V. Mazumder, S. Ozkar, S. Sun, Journal of the American Chemical Society 132 (2010) 1468. [5] F. Durap, M. Zhamakiran, S. Ozker, Applied Catalysis A: General 369 (2009) 53. [6] T.W. Graham, C.-W. Tsang, X. Chen, R. Guo, W. Jia, S.-M. Lu, C. Sui-Seng, C.B. Ewart, A. Lough, D. Amoroso, K. Abdur-Rashid, Angewandte Chemie International Edition 49 (2010) 8708. [7] T. Umegaki, J.-M. Yan, X.-B. Zhang, H. Shioyama, N. Kuriyama, Q. Xu, Journal of Power Sources 195 (2010) 8209. [8] P.V. Ramachandran, P.D. Gagare, Inorganic Chemistry 46 (2007) 7810. [9] H.-B. Dai, X.-D. Kang, P. Wang, International Journal of Hydrogen Energy 35 (2010) 10317. [10] H. Erdgan, O. Metin, S. Ozkar, Physical Chemistry Chemical Physics 11 (2009) 10519. [11] S.B. Kalidindi, U. Sanyal, B.R. Jagirdar, Physical Chemistry Chemical Physics 10 (2008) 5870. [12] S.B. Kalidindi, U. Sanyal, B.R. Jagirdar, Physical Chemistry Chemical Physics 11 (2009) 770. [13] L.C. Nagle, J.F. Rohan, Journal of the Electrochemical Society 153 (2006) C773. [14] L.C. Nagle, J.F. Rohan, Electrochemical and Solid-State Letters 8 (2005) C77. [15] X.B. Zhang, S. Han, J.M. Yan, H. Shioyama, N. Kuriyama, T. Kobayashi, Q. Xu, International Journal of Hydrogen Energy 34 (2009) 174. [16] V. Kiran, S.B. Kalidindi, B.R. Jagirdar, S. Sampath, Electrochimica Acta 56 (2011) 10493. [17] L.C. Nagle, J.F. Rohan, Journal of the Electrochemical Society 158 (2011) B772. [18] F. Marken, A. Neudeck, A.M. Bond, in: F. Scolz (Ed.), Electroanalytical Methods, 2nd edition, Springer, Heidelberg, 2010 (Chapter II.1). [19] J.E. Baur, R.M. Wightman, Journal of Electroanalytical Chemistry 305 (1991) 73. [20] Y. Ohsawa, S. Aoyagui, Journal of Electroanalytical Chemistry 136 (1982) 353. [21] J.E.A.M. van den Meerakker, Journal of Applied Electrochemistry 11 (1981) 395.