Article pubs.acs.org/JPCC
Hydrogen Generation from Hydrolysis and Methanolysis of Guanidinium Borohydride Leigang Li,†,‡ Shaofeng Li,‡ Yingbin Tan,‡ Ziwei Tang,‡ Wanyu Cai,§ Yanhui Guo,‡ Qian Li,*,† and Xuebin Yu*,‡ †
Shanghai Key Laboratory of Modern Metallurgy & Materials Processing, Shanghai University, Shanghai, 200072, China Department of Materials Science, Fudan University, Shanghai, 200433, China § Shanxi Rock New Materials Co., Ltd., Huizhongli, Chaoyang District, China ‡
S Supporting Information *
ABSTRACT: Metal-catalyzed hydrolysis and methanolysis of guanidinium borohydride (C(NH2)3BH4 or GBH) for hydrogen generation are reported. GBH is comparatively stable in water with only 0.3 equiv of H2 liberated in 24 h at 25 °C while it reacts vigorously with methanol, releasing more than 3.2 equiv of H2 within only 17 min. Even at 0 °C, there was still nearly 2.0 equiv of H2 released after 2 h, but no H2 liberation was observed for hydrolysis under the same conditions. Various metal chlorides were adopted to enhance the reaction kinetics of the hydrolysis and methanolysis, of which CoCl2 exhibits the highest activity in both cases. With the addition of 2.0 mol % CoCl2 at 25 °C, the methanolysis of GBH could generate 4 equiv of H2 within 10 min with a maximum hydrogen generation rate of 9961.5 mL·min−1·g−1 while only 1.8 equiv of H2 was obtained under the same conditions at a maximum hydrogen generation rate of 692.3 mL·min−1·g−1 for hydrolysis. Compared with hydrolysis, methanolysis of GBH possesses much faster reaction kinetics, rendering it an advantage for hydrogen generation, especially at subzero areas. It was proposed that the faster reaction kinetics of methanolysis of BH4− containing compounds is ascribed to the more electron donating methoxy group than that of hydroxyl group. Moreover, a comparison between hydrolysis and methanolysis of GBH indicates that the loss of the first H from BH4− controls the hydrolysis kinetics instead of the cleavage of the O−H bond.
1. INTRODUCTION The continuously rising concerns over dwindling energy resources and the environmental impact of burning fossil fuels (petroleum, for example) have triggered intensive attention on the exploration of sustainable and renewable alternative energy carriers, such as nuclear energy, solar energy, and hydrogen energy. Among the alternative energy carriers, hydrogen has been in the limelight for the past several years from the standpoint of the so-called hydrogen energy economy, primarily due to its clean burning nature, environmental friendliness and high energy content.1 Hydrogen can be readily used in fuel cells in vehicles or in portable electronic devices with high efficiency. However, the main obstacle in the way of transition to a hydrogen economy is the absence of a practical means for hydrogen storage.2 In view of the futuristic aspect of the use of hydrogen as a clean fuel and the hurdles in its safe and efficient storage, various solutions have been investigated to solve the problem of hydrogen storage during the past several decades, including tanks of compressed and liquefied H2,3 molecular clathrates,4 metal hydrides,5 and chemical complexes.6−8 In this last category, boron-based chemical hydrides (e.g., NH3BH3,9,10 LiBH4,11 NH3B3H712,13) cover a central role owing to their high gravimetric and volumetric hydrogen content. They can generate hydrogen to supply fuel cells via © 2012 American Chemical Society
either pyrolysis, or hydrolysis, and the latter one prevails over the former route from several aspects. Hydrogen generation from hydrolysis of chemical hydrides (NaBH4, for example) can proceed at room temperature with no external heat input, and its byproduct, borax, is environmentally benign. Another advantage of hydrogen generation by hydrolysis of chemical hydrides is that half of the generated hydrogen originates from water, and the slightly humid hydrogen can be directly used in a polymer electrolyte membrane fuel cell (PEMFC). These distinct advantages of hydrogen generation from hydrolysis of boron hydrides make it a promising on-board hydrogen generation method for portable PEMFC. In the field of hydrolysis research, NH3BH314−17 and NaBH418−22 have attracted more and more attention. In the meanwhile, the hydrolysis of several other boron-containing compounds was also investigated, such as NaB 3 H 8 , 2 3 NH 4 B 3 H 8 , 2 4 NH3B3H7,12,13 LiBH4,25 KBH4,26 and N2H4BH3.27 Guanidinium borohydride C(NH2)3BH4 (denoted as GBH hereafter), another boron-containing chemical hydride, was first reported by Schechte in 1954.28 Lately, Groshens29 and Received: April 6, 2012 Revised: June 13, 2012 Published: June 13, 2012 14218
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Rieger30 prepared pure GBH and investigated its thermal decomposition properties, respectively, where GBH was proposed as a hydrogen storage material. However, a problem with GBH as a solid state hydrogen storage material is that much ammonia is released during its thermal decomposition, which makes it unsuitable to supply hydrogen for PEMFC. More recently, GBH was prepared by ball milling in our group, and its dehydrogenation properties were improved by adding LiBH4 or Ca(BH4)2.31,32 Resembling NaBH4, GBH also contains BH4−, possessing high hydrogen content. However, there have been no previous explorations of hydrogen generation from hydrolysis of GBH and therefore hydrolysis of GBH was investigated in this manuscript to investigate its hydrogen generation properties. Additionally, methanolysis of GBH was also investigated to generate hydrogen more quickly33 and to curb the disadvantages of direct hydrogen generation from methanol cracking, which is carried out under harsh conditions of high temperature and high pressure and accompanied with environmentally detrimental byproduct of CO2 and toxic CO. A comparison of the hydrolysis and methanolysis provides a possible evidence for clarifying the rate-controlling step of the hydrolysis of BH4−. Methanolysis of GBH outperforming its hydrolysis, offers a promise for lowtemperature hydrogen generation.
2.3. Characterizations. Powder X-ray diffraction (XRD) patterns were obtained on a D8 Advanced Bruker diffractometer (Cu Kα radiation, 80KV). During the XRD measurement, samples were mounted in a glovebox, and an amorphous polymer tape was used to cover the surface of the powder to avoid oxidation and moisture. Solid-state Fourier transform infrared (FT-IR) measurements of the samples (as KBr pellets) were recorded with a Nicolet Nexus 470. During the IR measurements (KBr pellets), samples were loaded into a single closed tube with CaF2 windows. Liquid 11B NMR spectra were collected at room temperature using Bruker DMX500 spectrometer. D2O and CD3OD were used as lockers for the filtrates of hydrolysis and methanolysis, respectively. During the liquid 11B NMR measurements, the filtrates were referenced externally to BF3.(C2H5)2O at 0 ppm. All sample manipulations were performed in an argon-filled glovebox that kept both water and oxygen concentrations below 0.1 ppm.
3. RESULTS AND DISCUSSION 3.1. Hydrolysis of GBH. GBH was prepared via a liquidphase synthesis route and characterized by XRD, FT-IR (Figure S1 and S2, Supporting Information) and 11B NMR test (see Figure 3a below). Figure 1 shows the hydrogen generation
2. EXPERIMENTAL SECTION 2.1. Reagents and Sample Preparation. All the raw materials were obtained commercially. NaBH 4 (95%), guanidinium chloride (CN3H5·HCl, 98%), PtCl2 (98%) were all purchased from Sigma-Aldrich. NiCl2 (99.9%), CoCl2 (99.7%), ZnCl2 (99.8%), CuCl2 (98%), MgCl2 (99.9%), and LiCl (99.9%), were all purchased from Alfa-Aesar. GBH was synthesized by the reaction of NaBH4 and guanidinium chloride (mole ratio 1:1) in tetrahydrofuran (THF) following a literature procedure. THF and methanol were dried by distillation over metallic sodium and magnesium ribbon, respectively, before using. 2.2. Hydrogen Generation Measurements. In a typical experiment, GBH and metal catalysts were placed and mixed in a sealed two-necked flask fitted with an outlet tube for collecting the evolved H2. The outlet tube exhaust was placed under an inverted, water-filled graduated cylinder that was situated in a water-filled vessel. The flask was immersed in a water bath to maintain the temperature at relatively constant value. Then water or methanol is injected into the mixture of solid GBH and catalyst, and the solution was stirred by magnetic stirrer at a fixed rate during the hydrogen generation measurement. The reaction time was calculated starting from the time water or methanol was injected into the mixture, and the progress of the reaction was monitored by following the amount of hydrogen generated using a volumetric technique, which allowed tracking of the reaction in real time. During the methanolysis test, methanol vapor in the product stream was anticipated because of its significant vapor pressure at experimental conditions, but it can be absorbed by water to ensure that the volume change in the inverted graduated cylinder was due only to hydrogen. At the end of the reactions of hydrolysis and methanolysis, the resulting solutions were filtered, and the filtrates were collected for the liquid 11B NMR test. All the hydrogen generation rates (mL·min‑1·g‑1) for hydrolysis and methanolysis were calculated by dividing the hydrogen volume (mL) with the corresponding time (min) and the weight of catalyst (g) according to the common method.
Figure 1. Hydrogen generation plots of mol H2/mol GBH versus time (s) for the catalytic hydrolysis of GBH using different metal chlorides (2.0 mol % with reference to GBH).
plots for hydrolysis of GBH. Unlike the solutions of sodium borohydride, ammonia triborane, and dimethylamine borane, GBH solution shows high stability at room temperature against self-hydrolysis, as proved by the hydrogen evolution results, which shows a lag time of 20 min, and the reaction was far from complete even after 24 h with 93% of GBH unreacted, while 50% of sodium borohydride hydrolyzed.23 The relatively stable nature of GBH aqueous solution shows great advantage over NaBH4, as NaOH is needed to stabilize the NaBH4 aqueous solution.34 Since GBH is fairly stable in water, catalysts should be used to accelerate the reaction kinetics of hydrolysis. Five different metal chlorides (2.0 mol % with reference to GBH) were adopted to accelerate the hydrolysis, among which CoCl2 exhibits the highest activity, in accordance with the literature results.23 When deionized water was injected into the mixtures of GBH and CoCl2, black powder formed instantly, which is entirely possible to be Co−B compound, as indicated by the literature result.23 PtCl2 shows a catalytic effect similar to that of CoCl2, but with the H2 evolution rate decreasing after 8 min. 14219
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exhibits one sharp peak centered at 16.31 ppm resulting from a fast-exchanging equilibrium mixture of B(OH)3/B(OH)4−, along with a weak broad signal situated at around 12.92 ppm, which is associated with B3O3(OH)4−.12,13,35−37 The 11B NMR results indicated that complete hydrolysis of GBH solution can be achieved with 6.0 mol % CoCl2 loading. Hydrolysis kinetics is not only dependent on the catalyst concentration, but also on other factors (for example, reaction temperature). Generally, the chemical reaction rate increases with the increasing reaction temperature. Kinetic studies at varied temperatures were further carried out using the optimized concentration of CoCl2 (6.0 mol %). Figure 4
The other three chlorides, ZnCl2, NiCl2, and CuCl2, showed similar catalytic effects, but were much less effective than was CoCl2. Although 2.0 mol % CoCl2 exhibits the best catalytic effect among the above-mentioned metal chlorides, the hydrolysis of GBH was still not complete, and the H2 generation rate was not fast enough to meet the on-board application requirements. Then different mole percent of CoCl2 (with reference to GBH) was added into GBH to further accelerate the hydrogen generation rate. The hydrogen generation rate was largely increased with the increase of CoCl2, as illustrated by Figure 2.
Figure 2. Hydrogen evolution plots versus time (s) at 25 °C from aqueous solution containing 2.0, 4.0, 6.0, and 8.0 mol % CoCl2. Inset shows the relationship between hydrogen generation rate and concentration of CoCl2.
Figure 4. Hydrogen generation plots versus time (s) at different solution temperatures with the same concentration of GBH solution and CoCl2 concentration (6.0 mol %). Inset shows the Arrhenius plot of the H2 generation rates.
The inset of Figure 2 presents the relationship between the H2 generation rate and the concentration of CoCl2. It is revealed that the hydrogen generation rate increased exponentially with the increase of CoCl2 loading. With 6.0 mol % and 8.0 mol % CoCl2 loading, respectively, nearly the same volume of H2 was obtained, which may be an indicator of complete hydrolysis of GBH. In order to test whether the hydrolysis of GBH solution with 6.0 mol % CoCl2 loading was completed, liquid 11B NMR analysis was performed on the filtrate of hydrolysis solution (Figure 3). As illustrated by the spectra of Figure 3a,b, peaks related to GBH disappeared while a sharp and broad peak appeared, confirming the complete conversion of the starting material GBH (−39.75 ppm) to guanidinium metaborate. For the hydrolysis of GBH solution, the 11B NMR spectrum
presents the hydrogen generation kinetic curves at a solution temperature ranging from 20 to 40 °C. As expected, the hydrogen generation rate increased significantly with the increment of temperatures. The influence of temperature is clearly shown in Figure 4, by both the increasing slope values of hydrogen generation plots and the decreasing reaction completing time. For example, the hydrogen generation rate increased from 47.3 mL·min−1·g−1 at 20 °C to 168.9 mL·min−1·g−1 at 40 °C, and the completed reaction time decreased from 19 min at 20 °C to 6 min at 40 °C. In order to evaluate the efficacy of CoCl2, the hydrogen generation rate was treated through the Arrhenius equation, and the apparent activation energy of GBH hydrolysis catalyzed by CoCl2 was determined to be 51.5 kJ/mol (inset in Figure 4), which is in the range of 30−90 kJ/mol found for metal catalyzed hydrolysis of sodium borohydride, ammonia borane, and ammonia triborane.27 3.2. Methanolysis of GBH. Besides the hydrolysis of GBH, hydrogen generation from methanolysis of GBH was also investigated. Methanol was preferred owing to its lightest weight and highest reactivity among primary alcohols.33,38,39 As it can be seen from Figure 5, methanolysis of GBH without addition of catalysts was violent and 3.2 equiv of H2 was generated while it was only 0.3 equiv of H2 for hydrolysis of GBH under the same conditions (see Figure 8 below). Although methanolysis of GBH outperforms its hydrolysis in the reaction kinetics, it was not fully reacted. Several metal chlorides were screened to catalyze the methanolysis of GBH. As illustrated by Figure 5, LiCl has no effect on the kinetics or the conversion extent of methanolysis. MgCl2 and ZnCl2 can
Figure 3. Liquid 11B NMR spectra for GBH in THF (a), hydrolysis product (b), and methanolysis product (c) at room temperature. 14220
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NaBH4 methanolysis without catalyst can be achieved at 60 °C with 4.0 equiv of H2 released.38 So it seems that the methanolysis of GBH at low temperatures without catalyst can be classified into two phases, i.e., the two stages before and after full precipitation of the byproduct on unreacted GBH. In order to obtain the activation energy of the first step of GBH methanolysis without catalyst, the hydrogen generation rate was treated by the Arrhenius equation, and the apparent activation energy of GBH methanolysis was determined to be 11.2 kJ/mol (inset in Figure 6), which is much smaller than that of GBH hydrolysis catalyzed by CoCl2. In addition, we also investigated the concentration dependence of GBH methanolysis without catalyst. As shown in Figure 7, the concentration of GBH has a slight influence on the rate Figure 5. Hydrogen generation plots of mol H2/mol GBH versus time (s) for the catalytic methanolysis of GBH using different metal chlorides (2.0 mol %).
significantly accelerate the kinetics, especially at the initial stage, but also have no efficacy on the reaction extent. Meanwhile, CoCl2 and CuCl2 exhibit the similar and best catalytic effect, not only on the reaction kinetics but also on the reaction extent. In the presence of 2.0 mol % CoCl2, the methabolysis of GBH could undergo vigorously and complete within 10 min, generating 4.0 equiv of H2. At the end of methanolysis catalyzed by 2.0 mol % CoCl2, the filtrate of reaction was collected for 11B NMR analysis (Figure 3c). The single 11B NMR chemical shift of 2.76 ppm revealed the formation of guanidinium tetramethoxyborate40 and the full conversion of GBH methanolysis. After the methanolysis reaction, we once attempted to regenerate the methanolysis byproduct of guanidinium tetramethoxyborate according to the method described by Ramachandran et al.40 However, no product was obtained, which may result from the instantaneous reaction of the possibly regenerated GBH and in situ formed methanol. Figure 6 shows the influence of temperature on the reaction kinetics of methanolysis without the addition of metal catalyst.
Figure 7. Plots for hydrogen generation from methanolysis with different concentrations of GBH (mole ratio of GBH to methanol equals to 1:2, 1:4, and 1:8, respectively). Inset shows the relationship between H2 generation rate and concentration of GBH in methanol.
of hydrogen generation at the initial stage, and the amount of H2 generated increases nearly linearly with time before the methanolysis reaction slowly approaches its end. However, with the reaction proceeding, the concentration influence became more apparent, i.e., with more methanol added, the reaction kinetics became faster, and the reaction was more complete. The inset in Figure 7 shows the relationship between the hydrogen generation rate and the concentration of GBH, which shows a good linearity trend between the H2 generation rate and the concentration of GBH. As shown by the inset result in Figure 7, the hydrogen generation rate and concentration of GBH display a disproportionate relationship. The reason for this lies in the fact that with more methanol added, there are more chances for GBH to react with methanol. 3.3. Comparison between Hydrolysis and Methanolysis of GBH. The comparison of hydrogen generation between hydrolysis and methanolysis of GBH without catalyst is shown in Figure 8. As illustrated in Figure 8, methanolysis of GBH is extraordinarily violent with 2.0 equiv of H2 liberated within only 4 min, but GBH is comparatively stable in water. In detail, a lag time of about 20 min was observed for the hydrolysis with only 0.3 equiv of H2 liberated after 24 h. On the contrary, 3.2 equiv of H2 was generated in only 17 min, with a maximum hydrogen generation rate of 226.7 mL·min−1·g−1 while it is only 10.7 mL·min‑1·g‑1 for hydrolysis (seeing the inset in Figure 8). Even at 0 °C, nearly 2.0 equiv of H2 was generated after 2 h. The low freezing point of methanol is a significant advantage for generating hydrogen at subzero temperature under which
Figure 6. Hydrogen generation plots of mol H2/mol GBH versus time (s) at different temperatures with the same concentration of GBH in methanol. Inset shows the Arrhenius plot of the H2 generation rates.
In the tested temperature range (25−35 °C), no full reaction was achieved, but the reaction extent increased with the increasing temperatures. The reason for the above-mentioned phenomenon is that the methanolysis byproduct precipitates on the surface of unreacted GBH. In fact, complete reaction of 14221
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Figure 9. Schematic of reaction pathways of GBH with H2O and CH3OH, respectively.
What’s more, a comparison of methanolysis and hydrolysis may shed some light on the understanding of the ratecontrolling step of hydrolysis of BH4− ion. With respect to the rate-controlling step of hydrolysis of BH4− ion, there are two different viewpoints, that is, the loss of the first H from BH4− and the cleavage of water O−H bond. Guella et al.20,43 proposed that the cleavage of the water O−H bond controls the rate of BH4− hydrolysis, while, according to the results of ref 44, it is the cleavage of B−H that determines the reaction rate. On the basis of the discussion of the last paragraph, it may be concluded that it is the loss of the first H from BH4− that controls the rate of hydrolysis but not the cleavage of O−H bond, as it is the same value for the O−H bond in H2O and CH3OH.
Figure 8. Comparison of hydrogen generation plots versus time (s) for hydrolysis and methanolysis of GBH at 25 °C. Inset shows the comparison of H2 generation rate for hydrolysis and methanolysis, respectively.
conditions the reactant of water for hydrolysis reaction would exist as a solid. Therefore, hydrogen generation from methanolysis of GBH outperforms its hydrolysis, showing promise of providing a method for low temperature hydrogen generation. In fact, Karan et al. has already reported that the reaction kinetics of methanolysis of sodium borohydride was much faster than that of hydrolysis.33 However, no reason was given for the distinct kinetics. Considering the same reagent of GBH but different solvents of methanol and water in methanolysis and hydrolysis, respectively, the large difference in the rate of reaction must be related to the different solvents, and the solvents have affected the reaction kinetics in some manner. Hydrolysis or methanolysis of BH4− ion can be characterized as a process of hydride ion transfer and the formation of a dative B−O bond (as indicated by the 11B NMR results in Figure 3b,c) by donation of an electron pair of the hydroxyl or methoxy group into the empty orbital of boron, resembling the reaction of amine and borane in which a B−N bond formed.41 In terms of this hypothesis, the slow reaction kinetics of hydrolysis would be ascribed to the poor donor properties of oxygen atoms in the hydroxyl group, whereas the high reactivity of GBH with methanol would be attributable to the more electron donating properties of the oxygen atoms in methoxy group, which is owing to the higher electron density of the methoxy group than that of hydroxyl group; thereby more electrons transfer to the B atoms (δ2>δ1), leading to higher hydridic character and reactivity of hydrogen in the borohydride group during the methanolysis.38 Therefore, methanolysis of BH4− containing compounds possess faster reaction kinetics than that of hydrolysis. The discussion above could be illustrated by Figure 9. A ready example corroborating the above assumption lies in two alkali-metal amidoboranes, LiNH 2 BH 3 and NaNH 2 BH 3 (Li/NaAB),42 which were prepared through substituting one H in the NH3 group with the more electron donating elements of Li/Na. The introduction of Li/Na brings about considerable changes in the electronic state of N with concomitant stronger B−N interactions in Li/NaAB than in NH3BH3 (AB). As a consequence, the chemical bonding of B−H and N−H is affected, leading to significantly enhanced dehydrogenation kinetics of Li/NaAB compared with AB.
4. CONCLUSIONS In summary, we report for the first time, metal-catalyzed hydrolysis and methanolysis of GBH for hydrogen generation. With the addition of cobalt chloride, hydrolysis and methanolysis of GBH both exhibit fast reaction kinetics, especially for the methanolysis, which outperforms its hydrolysis at low-temperature hydrogen generation. For example, the maximum hydrogen generation rate at 25 °C for methanolysis of GBH is 226.7 mL·min‑1·g‑1 while it is only 10.7 mL·min‑1·g‑1 for hydrolysis. Even at 0 °C, nearly 2.0 equiv of H2 was generated after 2 h, but water was solid at this ice point. What’s more, we propose a potential mechanism to clarify the faster reaction kinetics of methanolysis than that of hydrolysis. Finally, a comparison between methanolysis and hydrolysis of GBH may shed some light on the understanding of the ratecontrolling step of BH4− hydrolysis, that is, the loss of the first H from BH4− controls the rate of hydrolysis instead of the cleavage of the O−H bond.
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ASSOCIATED CONTENT
S Supporting Information *
Preparation and characterization of GBH are given. This information is available free of charge via the Internet at http:// pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Phone and Fax: +86-21-5566 4581. E-mail: yuxuebin@fudan. edu.cn (X.Y.);
[email protected] (Q.L.). Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was partially supported by the Ministry of Science and Technology of China (2010CB631302), the National 14222
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Natural Science Foundation of China (Grant No. 51071047), the Ph.D. Programs Foundation of Ministry of Education of China (20110071110009), the Shanghai Rising-Star Program (11QH1400900), and the Science and Technology Commission of Shanghai Municipality (11JC1400700, 11520701100).
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