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Spontaneous hydrolysis of sodium borohydride in harsh conditions
i n t e r n a t i o n a l j o u r n a l o f h y d r o g e n e n e r g y 3 6 ( 2 0 1 1 ) 2 2 4 e2 3 3
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Spontaneous hydrolysis of sodium borohydride in harsh conditions Je´roˆme Andrieux a, Umit Bilge Demirci a,*, Julien Hannauer a, Christel Gervais b, Christelle Goutaudier a, Philippe Miele a a
Universite´ Lyon 1, CNRS, UMR 5615, Laboratoire des Multimate´riaux et Interfaces, 43 boulevard du 11 Novembre 1918, F-69622 Villeurbanne, France b Universite´ Paris 06, Laboratoire de Chimie de la Matie`re Condense´e Paris, CNRS, UMR 7574, Colle`ge de France 11, place Marcelin Berthelot F-75231 Paris Cedex 05, France
article info
abstract
Article history:
The present study reports fundamental results about the spontaneous hydrolysis of
Received 25 May 2010
sodium borohydride NaBH4, which is a potential hydrogen storage material for small,
Received in revised form
portable applications. The reaction (without stabilizing agent or catalyst) was carried out at
15 October 2010
temperatures of 30e80 C, initial NaBH4 concentrations of 0.63e6.19 mol L1 (2.3e18.9 wt%),
Accepted 19 October 2010
and for unbuffered solutions, which are harsher experimental conditions than those reported so far. The H2 evolution and the subsequent pH variation were observed to determine the reaction kinetic parameters and characterize the hydrolysis intermediates,
Keywords:
i.e. the hydroxyborates BH4z(OH) z , by XRD, IR and
Kinetics
apparent activation energy of the reaction was 98 10 kJ mol1 and the reaction order
Hydroxyborate anions
versus the initial NaBH4 concentration was 0; (ii) all of the reactions BH 4 / BH3(OH) ,
11
B NMR. It was found that: (i) the
Sodium borohydride
BH3(OH) / BH2(OH) 2 , BH2(OH)2 / BH(OH)3 and BH(OH)3 / B(OH)4 took place simulta-
Sodium metaborate
neously; (iii) only 25 mol% of B(OH) 4 and 75 mol% of BH4 were found at 25% of conversion;
Spontaneous hydrolysis
(iii) the hydroxyborates are very short-lived intermediates and only traces of BH3(OH) were detected. ª 2010 Professor T. Nejat Veziroglu. Published by Elsevier Ltd. All rights reserved.
1.
Introduction
Sodium borohydride NaBH4 was first synthesized in the 1940se1950s by Schlesinger et al. [1] and soon attracted much attention [2]. Stable in dry conditions, it spontaneously (exothermically) hydrolyzes in the presence of water while generating hydrogen: BH 4 ðaqÞ þ 4H2 OðlÞ/BðOHÞ4 ðaqÞ þ 4H2 ðgÞ
(1)
In the 1950se1960s, most of the studies focused on improving the knowledge of the kinetics and mechanisms of
this reaction (1) [3e14]. Today, interest in the hydrolysis of NaBH4 is focused on applications, such that the objective is typically to accelerate H2 release with the aid of metal catalysts [15e22]. The spontaneous hydrolysis is less investigated and it is even avoided by virtue of stabilizing agents such as NaOH [23]. The spontaneous hydrolysis of NaBH4 is an acid-catalyzed reaction (in-situ formed H3Oþ ions) [5,6]. Mochalov et al. [10,11] proposed a detailed mechanism based on the substitution of BeH bonds by BeOH bonds. The hydrolysis would be a multistep process involving hydroxyborate intermediates
* Corresponding author. Tel.: þ33372448403; fax: þ33372440618. E-mail address: [email protected] (U.B. Demirci). 0360-3199/$ e see front matter ª 2010 Professor T. Nejat Veziroglu. Published by Elsevier Ltd. All rights reserved. doi:10.1016/j.ijhydene.2010.10.055
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BH4z(OH) z (with z from 1 to 3). For example, the reactions (2) and (3) show the H3Oþ-catalyzed hydrolysis of BH 4 and the formation of the first intermediate BH3(OH) (i.e. z ¼ 1), respectively: þ BH 4 ðaqÞ þ H3 O ðaqÞ/BH3 ðaqÞ þ H2 OðlÞ þ H2 ðgÞ
(2)
BH3 ðaqÞ þ OH ðaqÞ/BH3 OH ðaqÞ
(3) BH 4
The formation of BH3 from involves an activated or even BH5 (or H2BH3) complex, which is HþBH 4 [5,6,8,10,11,14]. The reactions (2) and (3) are both rate determining steps. The spontaneous hydrolysis of NaBH4 depends on experimental conditions, e.g. temperature, solution pH and ratio H2O/NaBH4 [5,6,24]. For example, a well-known empirical relation (4) shows the dependence of the half-reaction time t1/2 (in min) on the pH and the temperature T (in K) [10,11]: log t1=2 ; NaBH4 ¼ pH ð0:034 T 1:92Þ
(4)
To illustrate, based on this equation, hydrolysis should achieve a conversion of 50% in 134 min at pH 10 at 15 C, 28 min at pH 10 at 35 C, 13271 min (i.e. >9 days) at pH 12 at 15 C and 2773 min (i.e. about 2 days) at pH 12 at 35 C. This equation was established in the following experimental conditions: 7 < pH < 10 and 15 < T ( C) < 35. Even though it does not cover pH ranges >10 and temperatures >35 C, it is an indicative tool that gives a satisfactory approximation of the effects of both pH and temperature [15]. From this brief literature survey, it is clear that the kinetic parameters reported so far were obtained at specific experimental conditions. Most of the studies were carried out with aqueous NaBH4 solutions at low concentrations (e.g. 0.01 mol L1), which besides could be buffered [10,11,24]. Another point is that the kinetic data were calculated on the few first percentages of NaBH4 converted [10,11,14]. Accordingly, we studied the kinetics of the spontaneous hydrolysis of NaBH4 at higher concentrations (up to about 6 mol L1) and high temperatures (30e80 C). Furthermore, the solutions were not buffered. Our objective was to contribute to a better understanding of the reaction in different (more severe) experimental conditions while attempting to identify the reaction intermediates upon the evolution of 1 equiv H2. Note that such severe conditions are much closer to the conditions that should be used in technological applications, even though catalysts will be used to enhance kinetics of H2 release.
2.
225
overall system was considered as a non-steady state slurry batch reactor. The H2O/NaBH4 molar ratio was equal to 9 (18.9 wt% of NaBH4 or about 6.18 mol L1). To determine the reaction order with respect to the initial NaBH4 concentration, molar ratios of 13 (13.9 wt%), 18 (10.4 wt%), 27 (7.2 wt%), 36 (5.5 wt%) and 90 (2.3 wt%) were considered by varying the NaBH4 mass. The NaBH4 content was below its solubility limit (55 g per 100 g H2O) on the temperature range we used, namely 30e80 C. Prior to any experiment, the NaBH4 solid was transferred into the tube. In a typical experiment, water was injected into the reactor with a needle placed directly inside the bed of NaBH4. The injection needle was immediately removed following water injection. An automated burette was used to control the amount of water. The H2 produced was collected and evacuated through a second needle, also inserted in the tube through the stopper and connected to an inverted water-filled graduated cylinder. The H2 volumes were measured as a function of time by video monitoring the displaced water as the reaction proceeded. The hydrolysis tests were stopped when the H2 generation rate (HGR) was lower than 0.0001 mL s1, which is equivalent to a maximum conversion value of 50e60%, depending on the experiment conditions (the HGRs decreased because of the pH increase as suggested by Equation (4)). A thermal probe placed inside the NaBH4 bed measured the temperature during the reaction. The NaBH4 conversion, 3 (%), was calculated as being the ratio of the experimental volume of H2 released at a given time to the expected maximum volume. The HGR (r, mL s1) was obtained by linear regression of the hydrogen evolution curves. Because HGR changed with the reaction conversion for spontaneous hydrolysis, it was determined by taking the slope of the curve at a given NaBH4 conversion range defined by 3 3%. For example, at 3 ¼ 20%, the range 17e23% was considered and the data over this range were linearized. From the as-obtained line (with R2 0.997), the HGR (i.e. the line slope) was calculated. The hydrolysis by-products were analyzed by X-ray diffraction (Bruker D5005 powder diffractometer, CuKa radi˚ ), Fourier transform infrared spectroscopy ation (l ¼ 1.5406 A
Experimental
Commercial NaBH4 (ACROS organics, 98% purity, powder, n CAS: 16940-66-2, average particle size of 200 mm) was used as received. It was stored and handled in an argon-filled glove box in order to avoid moisture contact and subsequent hydrolysis. Distilled water was purged with argon to remove oxygen. The hydrolysis experiments were performed in a 20 mL tube closed by a silicon stopper and placed in a thermostatic bath. For the determination of the kinetic parameters the
Fig. 1 e Hydrogen evolutions obtained at various temperatures (30, 50, 60, 70 and 80 C) for an H2O/NaBH4 molar ratio of 9.
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Table 1 e HGRs (r, mL sL1) obtained for the H2O/NaBH4 molar ratio 9 (18.9 wt%), at various temperatures and determined for various conversions (3 ± 3%). HGR (r, mL s1) at various conversions
Temperature ( C) 5% 30 50 60 70 80
0.006 0.03 0.08 0.24 0.53
10% 0.003 0.016 0.03 0.10 0.33
15% 0.0025 0.009 0.021 0.05 0.15
20% 0.0015 0.006 0.014 0.04 0.14
25% 0.001 0.004 0.008 0.03 0.07
30% 0.0005 0.004 0.006 0.021 0.06
35% a
e 0.002 0.005 0.012 0.043
40% a
e 0.0012 0.003 0.007 0.04
45% a
e 0.0008 0.003 0.007 0.024
50% ea 0.0008 0.0015 0.005 0.017
a The hydrolysis was stopped because of HGRs <0.0001 mL s1.
(FT-IR, FTIR Nicolet 380) and 11B nuclear magnetic resonance (either Bruker Avance500 MHz for solid compounds or Bruker DRX 500 400 MHz for ions in aqueous solution). 11B MAS NMR spectra of solid were recorded at 11.7 T on a Bruker Avance500 wide-bore spectrometer operating at 160.49 MHz, using
a Bruker 4 mm probe and a spinning frequency of the rotor of 14 kHz. Spectra were acquired using a spin-echo q-t-2q pulse sequence with q ¼ 90 to overcome problems of probe signal and recycle delay of 1s was used. Chemical shifts were referenced to BF3(OEt)2 (d ¼ 0 ppm).
Fig. 2 e Plots of ln(r) as a function of TL1 for various conversions, i.e. (a) 5%, (b) 10%, (c) 30%, and (d) 50%, and (e) apparent activation energies as a function of conversion.
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3.
Results and discussion
3.1.
Kinetic study
The NaBH4 spontaneous hydrolysis was studied from 30 to 80 C. The molar ratio H2O/NaBH4 was kept at 9. Fig. 1 shows the H2 evolutions obtained at 30, 50, 60, 70 and 80 C. As expected, the higher the temperature, the more the hydrolysis extent (in terms of HGRs and conversions at a given time). This is also shown by the data given in Table 1. Moreover, the curves had similar profiles; for a given temperature, the HGRs decreased with time, i.e. with the reaction conversion. As suggested by Equation (4), this was due to an increase of the solution pH because of the formation of the basic hydroxyborates BH4z(OH) z [10,11]. In fact, it was observed that at conversions of around 50% the HGRs were very low, at this point the experiments were stopped. The data in Table 1 were then exploited to plot InðrÞ ¼ f ð1=TÞ and the apparent activation energy was calculated. This was done for HGRs obtained at various conversions, i.e. 5 3%, 10 3%, 15 3%, 20 3%, 25 3%, 30 3%, 35 3%, 40 3%, 45 3%, and 50 3%. For example, Fig. 2 (aed) depicts some of the as-obtained linear regressions. It is important to note that for a given conversion the solution characteristics (pH, products concentration and so on) should have been similar, with the only different parameter being the temperature. Fig. 2(e) shows the apparent activation energies as a function of the conversion. The apparent activation energy is 98 10 kJ mol1 over the range 30e80 C (lit. [4e9,12,13]: 100 35 kJ mol1 at low conversions, low temperatures, i.e. < 35 C, and for buffered solutions). The reaction order versus the initial NaBH4 concentration was determined for various H2O/NaBH4 molar ratios (i.e. 9, 13, 18, 27, 36 and 90; 6.18, 4.27, 3.09, 2.08, 1.54 and 0.63 mol L1; 18.9, 13.9, 10.4, 7.2, 5.5 and 2.3 wt%) at 80 C. The H2 evolution curves were analyzed at both 20 3% and 40 3% (Fig. 3(a,b)); at these conversions, the H2 evolution was quite linear, and the HGRs were deduced by linear regression. Table 2 reports the as-determined HGRs. Three observations stood out: (i) the HGRs were quite similar, i.e. about 0.17 mL s1, at a conversion of 20% and irrespective of the molar ratios; (ii) a similar observation was reported at a conversion of 40% (HGRs of 0.04e0.07 mL s1); (iii) at a given reaction time, the lower the H2O/NaBH4 molar ratio (i.e. the higher the NaBH4 concentration), the higher the volume of H2 released. The HGR data were then used to plot their logarithm as a function of the logarithm of the initial NaBH4 concentration (Fig. 3(c)). A reaction order versus the initial NaBH4 concentration of zero was obtained at 20 3% and 40 3%. In other words, the spontaneous hydrolysis is independent on the initial NaBH4 concentration for concentrations up to 6.18 mol L1. To the best of our knowledge, reaction orders versus the initial NaBH4 concentration in these experimental conditions have not been reported to date, which does not make any comparison possible. It is noteworthy that a reaction order of 0 was obtained for lower conversions and temperatures and for buffered solutions [4e9,12,13]. To summarize the results discussed heretofore, it mainly stands out that the apparent activation energy of the
227
spontaneous hydrolysis determined over the range 30e80 C is 98 10 kJ mol1 and is consistent with the energies obtained at lower temperatures. Given that the apparent activation energy determined experimentally includes the energies of every occurring reaction, one may assume that all of the / BH3(OH), BH3(OH) / BH2(OH) reactions BH 4 2, BH2(OH)2 / BH(OH) 3 and BH(OH)3 / B(OH)4 take place simultaneously. Another assumption might be that the reac tion BH 4 / B(OH)4 preferentially takes place and predominates. To highlight this, the hydrolysis slurry was analyzed upon the evolution of 1 equiv H2.
3.2.
Effect of pH on hydrolysis
According to Mochalov et al. [10,11], the solution pH has a significant effect on the spontaneous hydrolysis kinetics. We therefore followed the pH evolution in the course of the spontaneous hydrolysis of NaBH4 at 30 C (Fig. 4(a)). The pH first increased quite rapidly from about 9 to 10.8 (in about 16 min).
Fig. 3 e Hydrogen evolution at conversions of 20% (a) and 40% (b) for various H2O/NaBH4 molar ratios, at 80 C, and (c) determination of the reaction order versus the NaBH4 concentration at 80 C and for conversions of 20 ± 3% and 40 ± 3%.
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Table 2 e HGRs (r) for various molar ratios and for conversions of 20 ± 3% and 40 ± 3%. H2O/NaBH4 molar ratio
[NaBH4]0 (mol L1)
NaBH4]0 (wt%)
Conversion of 20 3% 1
90 36 27 18 13 9
0.63 1.54 2.08 3.09 4.27 6.18
2.3 5.5 7.2 10.4 13.9 18.9
Conversion of 40 3%
r (mL s )
Dr
r (mL s1)
Dra
0.20 0.16 0.16 0.16 0.17 0.14
0.06 0.05 0.03 0.02 0.04 0.02
0.07 0.07 0.08 0.07 0.06 0.04
0.03 0.03 0.02 0.01 0.01 0.05
a
a With Dr, the experimental error on r.
Then, the pH still increased but less rapidly. Actually the pH increase per unit of time was less and less high as the hydrolysis proceeded. It increased up to 11.7 after 30000 s of hydrolysis (i.e. >8 h). The profile of the pH evolution was similar to that of the H2 evolution. This clearly suggests that the solution pH strongly influences the hydrolysis kinetics. Nevertheless, the apparent activation energy is constant within the conversion range 0e50% and the hydrolysis is independent on the initial NaBH4 concentration. Hence, one may assume that the pH has no effect on the reaction mechanisms. The NaBH4 spontaneous hydrolysis is a complex reaction involving at least 5 different reaction species, from BH 4 to B(OH)4 , which play an important role in the evolution of the pH solution. Mochalov et al. suggested a stepwise hydrolysis of NaBH4 and showed that the half-reaction times of NaBH3(OH), NaBH2(OH)2 and NaBH(OH)3 were lower than that of NaBH4 [10,11]:
Fig. 4 e Evolution of the solution pH (experimental and calculated) as a function of time for an H2O/NaBH4 molar ratio of 9 at 30 C (a) and difference between the experimental and calculated pH values (b).
log t1=2 ; NaBH3 ðOHÞ ¼ pH ð0:027 T 0:357Þ
(5)
log t1=2 ; NaBH2 ðOHÞ2 ¼ pH ð0:027 T 0:384Þ
(6)
log t1=2 ; NaBHðOHÞ3 ¼ pH ð0:024 T 4:00Þ
(7)
The equations (4)e(7) were established under the following experimental conditions: 7 < pH < 10, 15 < T ( C) < 35, and buffered solutions. The rate constants of the reactions / BH2(OH) BH 4 / BH3(OH) , BH3(OH) 2 , BH2(OH)2 / BH (OH)3 and BH(OH)3 / B(OH)4 were found as 5.3 107, 3.6 1011, 10.1 107, 5.7 1010 mol1 min1, respectively. A
Fig. 5 e XRD patterns of the solids recovered at 5, 10, 15, 20, 25% of conversion from top to bottom (a) and XRD pattern of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (b); the stars show the peaks attributed to NaBH4 (ICDD 00-009-0386) while the others are assigned to NaB (OH)4 (ICDD 01-081-1512); the broad peaks within the range 2q [ 13 e28 is due to the amorphous film used to prevent the sample from moisture.
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constant rate of 5.2 107 mol1 min1 was also determined for the reaction BH 4 / B(OH)4 . In fact, BH4 can hydrolyze according to two routes: a stepwise one via the into B(OH) 4 formation of BH3(OH) or a direct one, the rate constants of the respective reactions being equivalent. It has been also reported that the concentrations of the BH3(OH) and BH (OH) 3 ions were small (0.03e0.05 mol%) in the course of the hydrolysis whereas the concentration of BH2(OH) 2 conversion w26 mol%. Molina Concha et al. [25] showed the formation of BH3(OH) through heterogeneous hydrolysis of NaBH4 and the presence of both BH3 and BH2 as stable intermediates by in-situ IR measurements. According to Ruman et al., it should be possible to synthesize BH3(OH) through the hydrolysis under ambient conditions of one BH 4 by one H2O in an excess of THF [26]. Its formation was evidenced by 1H NMR and the spectrum was similar to that of NaBH4 as it showed similar BeH coupling region but it was different in the region containing the BH resonances. An interesting point in Ruman et al’s conclusion
229
is that they assumed all of the BH 4 ions underwent a one-step hydrolysis and thus all led to the formation of BH3(OH). Unlike Mochalov et al’s studies [10,11], the NaBH4 solution was not buffered. If the discrepancies in the experimental conditions are not considered, Ruman et al’s work is in a way in contradiction with Mochalov et al’s. In order to highlight the discrepancy reported above, we tentatively calculated a theoretical pH evolution on the basis of the pKa of the couple B(OH)3/B(OH) 4 , which is 9.2 [27]. This couple is involved in the stepwise hydrolysis of NaBH4 [10,11]: þ BðOHÞ3 ðaqÞ þ H2 OðlÞ/BðOHÞ 4 ðaqÞ þ H ðaqÞ
(8)
Even though considering B(OH)3 as the only, first acid will overestimate the pH, it was assumed that the pH evolution was only dependent on the formation of B(OH) 4 from B(OH)3 and the following relation was applied, with B(OH)3 being a weak acid: h i pH ¼ 7 þ 0:5 pKa þ 0:5log BðOHÞ 4
(9)
Fig. 4(a,b) shows the calculated evolution. Three observations stood out: (i) the profile of the calculated evolution was quite similar to that of the experimental one; (ii) consistent with the acid/base couple considered and its pKa, the calculated pH was overestimated (DpH of about 1) at the beginning of the hydrolysis (1e3% of conversion) but interestingly DpH decreased to 0.1 at 17e18%; (iii) for conversions higher than 22% (i.e. t > 20000 s), the calculated pH was almost equal to the experimental pH. Observations (ii) and (iii) suggest therefore that the acid/base couple(s) involving the first hydrolysis intermediates such as e.g. BH3(OH) would have a pKa value lower than that of the couple B(OH)3/B(OH) 4 , and that the solution pH would be controlled by B(OH) 4 at conversion
Fig. 6 e IR spectra of NaBH4 and solids recovered at 5, 10, 15, 20, 25% of conversion from top to bottom (a) and IR spectrum of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (b).
Fig. 7 e Solid 11B RMN spectrum of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (a) and proton coupled 11B RMN spectrum of this product in aqueous (90 wt% H2O D 10 wt% D2O) solution.
230
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22%. In other words, it is important to identify the reaction intermediates that are involved in the beginning of the hydrolysis.
3.3.
Hydrolysis intermediates identification
To make clear the pH evolution, we focused our work on characterizing the hydrolysis intermediates at conversions up to 25%. The following tests were performed. The NaBH4 spontaneous hydrolysis was stopped at various conversions, i.e. 5, 10, 15, 20 and 25%, by dipping the tube in liquid air (<140 C). The slurry froze within 10 s, which is negligible in comparison to the timescale required for completion of spontaneous NaBH4 hydrolysis. The reaction was thus effectively quenched. The frozen slurry was dried under vacuum at
48 C for 48 h (to remove all water and prevent any hydrolysis when the sample was warmed) and then handled in an argon-filled glove box. The solids were analyzed by XRD (Fig. 5(a)). NaBH4 (ICDD 00-009-0386) was clearly identified at each conversion. It was the sole phase identified whereas a likely peak at around 2q ¼ 32 , characteristic of borates, had been especially regarded. No borate was found, which could be due an amorphous state [28]. Since the NaBH4 hydrolysis products have been reported as being stable up to 170 C [29], our solids were heated at 150 C for 4 h under Ar but such process was unsuccessful in crystallizing the borates. As a further test, the solids were then analyzed by IR (Fig. 6(a)). Three observations stood out. The spectrum recorded for NaBH4 was confirmed [30,31]. The NaBH4 print, i.e. nmax/ cm1 1114, 2220, 2293, 2341, 2357 and 2438 (BeH), was clearly
Fig. 8 e Proton decoupled (top) and coupled (bottom) 11B NMR spectra of the hydrolysis by-product recovered at a conversion of 25%.
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visible in each of the other spectra, showing that most of the reactant had not reacted. New bands, at nmax/cm1 500 d (BeO), 700e900 n(BeO), 1200 d(BeOeH), 1330 n(BeO), 1550e1650 d(HeOeH) and 3300e3500 n(OeH) [25,32], were observed for the partly hydrolyzed NaBH4, their respective contribution appearing to increase with the NaBH4 conversion. From such observations, it is in fact quite difficult to assert the formation of one hydroxyborate compound over another one. Tentatively our IR spectra might be compared to spectra previously reported and some/most of the bands might be assigned to e.g. NaBH3(OH) [29], NaBO2 [33,34], NaB(OH)4 [35], or NaBO2.2H2O [29,33,36]. In Ref. [25], bands at 970 and 1180 cm1 were attributed to BH2 and BH3, respectively. Nevertheless, none clearly matched and one may propose that the IR spectra indicate the formation of a tetra-coordinated boron but the presence of both BeH (especially from unreacted NaBH4) and BeOeH stretching modes impedes a clear identification of BH4z(OH) z or B(OH)4 . To succeed in identifying the compounds NaBH4z(OH)z that could be found in the hydrolysis medium at a conversion of 25%, we attempted to synthesize NaBH3(OH) according to the procedure recently reported by Ruman et al. [26]. Typically 1 equiv H2O per 1 NaBH4 were mixed in extra dry THF under Ar and the slurry was aged for 40 days while being stirred. Then, it was dried under vacuum for 24 h, transferred in the glove box, ground and analyzed by XRD. The XRD pattern is given in Fig. 5(b). The peaks could be assigned to both NaBH4 and NaB (OH)4. It is noteworthy that, some synthesis attempts showed unreacted NaBH4 and the presence of peaks assigned to NaB (OH)4. Nevertheless, as NaBH3(OH) is not listed in our XRD patterns database, we did not discard its formation. The synthesized product was then analyzed by IR (Fig. 6(b)). The spectrum supported the formation of NaB(OH)4, matching IR bands reported elsewhere [29]. As our observations contradict those of Ruman et al. [26], we completed the analysis by performing 11B solid state MAS NMR. The spectrum is given in Fig. 7(a). Two signals were observed at 2.2 and 41.8 ppm (difference of 44.0 ppm), corresponding to NaB(OH)4 and NaBH4, respectively [37]. Furthermore, we prepared an aqueous (90 wt% H2O and 10 wt% D2O) solution of this compound to be analyzed by 11B NMR. Water was injected into the sample-filled vial only few seconds before the analysis. Despite the H2 evolution, a spectrum not decoupled from proton in Fig. 7(b) was obtained. Once more, two signals were observed at 5.1 and 38.5 ppm (difference of 43.6 ppm), cor responding to B(OH) 4 and BH4 , respectively. The spectrum consisted of a singlet (NaB(OH)4) and a quintet (NaBH4). The JBH spinespin coupling constant for BH 4 was 80.6 Hz (lit. [12,13,26]: 82.0, 81.0 Hz). By focusing on d z 10 ppm, it was possible to distinguish within the background noise a negligibly small signal, likely a quartet that could be assigned to BH3(OH), which could have formed by hydrolysis of the dissolved solid. As a result, XRD, IR and NMR strongly suggested the formation of NaB(OH)4 rather than that of NaBH3(OH). In other words, when 1 equiv H2O is intended to hydrolyze 1 NaBH4, it is in fact 25% of the NaBH4 that hydrolyzes whereas 75% does not hydrolyze. This implies that at a conversion of 25% the specie controlling the solution pH is NaB(OH)4. 11 B RMN was also used to analyze the slurry recovered at a conversion of 25 1%. In this case, despite a low H2 evolution,
231
Fig. 9 e Spontaneous hydrolysis of NaBH4 and intermediates effect on pH over the conversion range 0e25%.
the 11B NMR collecting was done for the solution to keep the slurry state unchanged. The temperature was set at 5 C to minimize the H2 evolution. Fig. 8(a) shows the proton decoupled spectrum. It showed 2 signals at 38.4 and 4.8 ppm (difference of 43.2 ppm), characteristic of BH 4 and B(OH)4 [37]. A very small signal was observed at 9.9 ppm, ascribed to BH3(OH). Assuming a conversion of NaBH4 of 25%, the integration gave a relative content in BH3(OH) of 0.4 mol%, rather in agreement with Mochalov et al.’s observations [10,11]. No other hydroxyborate was detected. Guella et al. [37] investigated the hydrolysis intermediates for Pd-catalyzed hydrolysis of NaBH4. In a typical 11B proton decoupled NMR spectrum, no compound other than BH 4 and B(OH)4 was present, what was attributed to the low symmetry of the hydroxyborates structure leading to a broadening of 11B signals. Another reason could be the presence of the Pd catalyst, which accelerated the hydrolysis and so decreased the lifetime of very short-lived intermediates such as BH3(OH) [10,11]. It is noteworthy that, for uncatalyzed hydrolysis of NaBH4, the main thermodynamically stable borate is NaB(OH)4 in experimental conditions different from ours [28,38,39]. Fig. 8(b) shows the 11B spectrum not decoupled from proton, which is rather consistent with that reported by Gardiner and Collat [12,13]. The spectrum consisted of a singlet (NaB(OH)4), a quartet (NaBH3(OH)) and a quintet (NaBH4). The JBH spinespin coupling constants were calculated. The JBH value for NaBH3(OH) was 87.6 Hz (lit. [12,13,26,37]: 87.5, 82.0, 81.0 Hz). The JBH value for NaBH4 was 80.6 Hz (lit. [12,13,26]: 82.0, 81.0 Hz). The presence of the hydroxyborates BH2(OH) 2 and BH(OH)3 was especially scrutinized by exploring the shifts within the range defined by the signals of the borates BH3(OH) and B(OH) 4 , a shielding being expected. As shown in Fig. 8, no BH2(OH) 2 or BH(OH)3 were detected. In other words, when NaBH3(OH) was synthesized by reacting 1:1H2O with NaBH4 or by quenching the NaBH4 hydrolysis at a conversion of 25%, mainly NaB(OH)4 was formed with traces of NaBH3(OH).
4.
Conclusion
The spontaneous hydrolysis of NaBH4 was studied under experimental conditions that are much more severe than those
232
i n t e r n a t i o n a l j o u r n a l o f h y d r o g e n e n e r g y 3 6 ( 2 0 1 1 ) 2 2 4 e2 3 3
reported to date (i.e. T ¼ 30e80 C, [NaBH4]0 ¼ 0.63e6.19 mol L1, solutions not buffered). Under these conditions, an apparent activation energy of 98 10 kJ mol1 and a reaction order versus the initial NaBH4 concentration of 0 were found. Besides, a parallelity in the evolution of the solution pH with that of the NaBH4 conversion was observed. The hydrolysis intermediates at conversions up to 25% were recovered by dipping the hydrolysis slurry in liquid air (<140 C) and then dried under vacuum. The as-obtained solids were analyzed by XRD, IR and 11B NMR. The three characterization techniques gave the same results; they evidenced the formation of 25 mol% of B(OH) 4 and 75 mol% of BH 4 at a conversion of 25%. This was consistent with the kinetic data. It was besides confirmed that BH3(OH), which is the first hydrolysis intermediate, is a very short-lived intermediate. Traces of it were detected. No other hydroxyborate was detected by 11B NMR. As a result, one concludes that (i) the direct reaction preferentially occurred under our experiBH 4 / BH3(OH) mental conditions and (ii) the B(OH) 4 anions controlled the solution pH at conversions 22%. Fig. 9 depicts schematically our observations. Notably, our study has confirmed that at basic pH, the H2 release through spontaneous hydrolysis is harshly hindered.
Acknowledgements The present study was mainly funded by the ‘Cluster Energies’ of ‘Re´gion Rhoˆne-Alpes, France’. It was also funded by the ANR project BoraHCx. The authors are grateful to Mrs. Caroline TOPPAN (Univ. Lyon 1, CCRMN) and Mr. Olivier Majoulet (Univ. Lyon 1, LMI) for some of the 11B NMR spectra and the subsequent valuable discussions.
references
[1] Schlesinger HI, Brown HC, Finholt AE, Gilbreath JR, Hoekstra HR, Hyde EK. Sodium borohydride, its hydrolysis and its use as a reducing agent and in the generation of hydrogen. J Am Chem Soc 1953;75:215e9. [2] Demirci UB, Miele P. Sodium tetrahydroborate as hydrogen/ energy carrier, its history. C R Chim 2009;12:943e50. [3] Kilpatrick M, McKinney CD. Kinetics of the reaction of lithium borohydride in aqueous solution. J Am Chem Soc 1950;72:5474e6. [4] Pecsok RL. Polarographic studies of the oxidation and hydrolysis of sodium borohydride. J Am Chem Soc 1953;75: 2862e4. [5] Davis RE, Swain CG. General acid catalysis of the hydrolysis of sodium borohydride. J Am Chem Soc 1960;82:5949e50. [6] Davis RE, Bromels E, Kibby CL. Boron hydrides. III. Hydrolysis of sodium borohydride in aqueous solution. J Am Chem Soc 1962;84:885e92. [7] Stockmayer WH, Miller RR, Zeto RJ. Kinetics of borohydride hydrolysis. J Phys Chem A 1961;61:1076e7. [8] Jolly WL, Mesmer RE. Exchange of hydrogen between water and hydroborate ion. J Am Chem Soc 1961;83:4470e1. [9] Mesmer RE, Jolly WL. The hydrolysis of aqueous hydroborate. Inorg Chem 1962;1:608e12.
[10] Mochalov KN, Khain VS, Gil’manshin GG. Kinetic studies on intermediate steps of BH 4 hydrolysis. Kinet Katal 1965;6: 541e4. [11] Mochalov KN, Khain VS, Gil’manshin GG. Generalized scheme for the hydrolysis of the borohydride ion and diborane. Dokl Akad Nauk 1965;162:402e6. [12] Gardiner JA, Collat JW. The hydrolysis of sodium tetrahydroborates. Identification of an intermediate. J Am Chem Soc 1964;86:3165e6. [13] Gardiner JA, Collat JW. Kinetics of the stepwise hydrolysis of tetrahydroborate ion. J Am Chem Soc 1965;87:1692e700. [14] Kreevoy MM, Hutchins JEC. H2BH3 as an intermediate in tetrahydroborate hydrolysis. J Am Chem Soc 1972;94: 6371e6. [15] Amendola SC, Sharp-Goldman SL, Janjua MS, Spencer NC, Kelly MT, Petillo PJ, et al. A safe, portable, hydrogen gas generator using aqueous borohydride solution and Ru catalyst. Int J Hydrogen Energy 2000;25:969e75. ¨, O ¨ zkar S. Hydrogen generation from the hydrolysis [16] Metin O of sodium borohydride by using water dispersible, hydrogenphosphate-stabilized nickel(0) nanoclusters as catalyst. Int J Hydrogen Energy 2007;32:1707e15. [17] Hung AJ, Tsai SF, Hsu YY, Ku JR, Chen YH, Yu CC. Kinetics of sodium borohydride hydrolysis reaction for hydrogen generation. Int J Hydrogen Energy 2008;33:6205e15. [18] Liu BH, Li Q. A highly active CoeB catalyst for hydrogen generation from sodium borohydride hydrolysis. Int J Hydrogen Energy 2008;33:7385e91. [19] Oronzio R, Monteleone G, Pozio A, De Francesco M, Galli S. New reactor design for catalytic sodium borohydride hydrolysis. Int J Hydrogen Energy 2009;34:4555e60. [20] Akdim O, Demirci UB, Miele P. More reactive cobalt chloride in the hydrolysis of sodium borohydride. Int J Hydrogen Energy 2009;34:9444e9. [21] Prosini PP, Gislon P. Water consumption during solid state sodium borohydride hydrolysis. Int J Hydrogen Energy; 2010. doi:10.1016/j.ijhydene.2010.07.094. [22] Ding XL, Yuan X, Jia C, Ma ZF. Hydrogen generation from catalytic hydrolysis of sodium borohydride solution using CobalteCoppereBoride (CoeCueB) catalysts. Int J Hydrogen Energy; 2010. doi:10.1016/j.ijhydene.2010.07.030. [23] Hua D, Hanxi Y, Xinping A, Chuansin C. Hydrogen production from catalytic hydrolysis of sodium borohydride solution using nickel boride catalyst. Int J Hydrogen Energy 2003;28:1095e100. [24] Jensen EH. A study on sodium borohydride with special reference to its analytical application in organic chemistry. Copenhagen: Nyt. Nordisk Forlag; 1954. [25] Molina Concha B, Chatenet M, Maillard F, Ticianelli EA, Lima FHB, de Lima RB. In situ infrared (FTIR) study of the mechanism of the borohydride oxidation reaction. Phys Chem Phys Chem 2010;12:11507e16. [26] Ruman T, Anna Ku snierz A, Jurkiewicz A, Le s A, Rode W. The synthesis, reactivity and 1H NMR investigation of the hydroxyborohydride anion. Inorg Chem Commun 2007;10: 1074e8. [27] CRC handbook of chemistry and physics. 81st ed. Boca Raton: CRC Press; 2000. [28] Stepanov N, Uvarov V, Popov N, Sasson Y. Study of byproduct of NaBH4 hydrolysis and its behavior at a room temperature. Int J Hydrogen Energy 2008;33:7378e84. [29] Goubeau VJ, Kallfass H. Die reaktion natriumborhydrid and wasser. Z Anorg Allg Chem; 1959:160e9. [30] Harvey KT, McQuaker NR. Infrared and Raman spectra of potassium and sodium borohydride. Can J Chem 1971;49: 3272e81. [31] Somer M, Acar S, Koz C, Kokal I, Ho¨hn P, Cardoso-Gil R, et al. a- and b-Na2[NH4][NH2]: two modifications of a complex
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hydride in the system NaNH2-NaBH4; syntheses, crystal structures, thermal analyses, mass and vibrational spectra. J Alloys Compd 2009;491:98e105. Jun L, Shuping X, Shiyang G. TF-IR and Raman spectroscopic study of hydrated borates. Spectrochim Acta 1995;51A: 519e32. Miller FA, Wilkins CH. Infrared spectra and characteristics frequencies of inorganic ions. Anal Chem 1952;24:1253e61. Molina Concha B, Chatenet M, Coutanceau C, Hahn F. In situ infrared (FTIR) study of the borohydride oxidation reaction. Electrochem Commun 2009;11:223e6. Janda R, Heller G. IR- and Ramanspektren isotop markierter tetra- und pentaborate? Spectrochim Acta 1982;36: 997e1001.
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[36] Kanturk A, Sari M, Piskin S. Synthesis, crystal structure and dehydration kinetics of NaB(OH)4$2H2O. Korean J Chem Eng 2008;25:1331e7. [37] Guella G, Zanchetta C, Patton B, Miotello A. New insights on the mechanism of palladium-catalyzed hydrolysis of sodium borohydride from 11B NMR measurements. J Phys Chem B 2006;110:17024e33. [38] Marrero-Alfonso EY, Davis TA, Matthews MA. Minimizing water utilization in hydrolysis of sodium borohydride: the role of sodium metaborate hydrates. Int J Hydrogen Energy 2007;32:4717e22. [39] Marrero-Alfonso EY, Gray JR, Davis TA, Matthews MA. Hydrolysis of sodium borohydride with steam. Int J Hydrogen Energy 2007;32:4723e8.
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Spontaneous hydrolysis of sodium borohydride in harsh conditions Je´roˆme Andrieux a, Umit Bilge Demirci a,*, Julien Hannauer a, Christel Gervais b, Christelle Goutaudier a, Philippe Miele a a
Universite´ Lyon 1, CNRS, UMR 5615, Laboratoire des Multimate´riaux et Interfaces, 43 boulevard du 11 Novembre 1918, F-69622 Villeurbanne, France b Universite´ Paris 06, Laboratoire de Chimie de la Matie`re Condense´e Paris, CNRS, UMR 7574, Colle`ge de France 11, place Marcelin Berthelot F-75231 Paris Cedex 05, France
article info
abstract
Article history:
The present study reports fundamental results about the spontaneous hydrolysis of
Received 25 May 2010
sodium borohydride NaBH4, which is a potential hydrogen storage material for small,
Received in revised form
portable applications. The reaction (without stabilizing agent or catalyst) was carried out at
15 October 2010
temperatures of 30e80 C, initial NaBH4 concentrations of 0.63e6.19 mol L1 (2.3e18.9 wt%),
Accepted 19 October 2010
and for unbuffered solutions, which are harsher experimental conditions than those reported so far. The H2 evolution and the subsequent pH variation were observed to determine the reaction kinetic parameters and characterize the hydrolysis intermediates,
Keywords:
i.e. the hydroxyborates BH4z(OH) z , by XRD, IR and
Kinetics
apparent activation energy of the reaction was 98 10 kJ mol1 and the reaction order
Hydroxyborate anions
versus the initial NaBH4 concentration was 0; (ii) all of the reactions BH 4 / BH3(OH) ,
11
B NMR. It was found that: (i) the
Sodium borohydride
BH3(OH) / BH2(OH) 2 , BH2(OH)2 / BH(OH)3 and BH(OH)3 / B(OH)4 took place simulta-
Sodium metaborate
neously; (iii) only 25 mol% of B(OH) 4 and 75 mol% of BH4 were found at 25% of conversion;
Spontaneous hydrolysis
(iii) the hydroxyborates are very short-lived intermediates and only traces of BH3(OH) were detected. ª 2010 Professor T. Nejat Veziroglu. Published by Elsevier Ltd. All rights reserved.
1.
Introduction
Sodium borohydride NaBH4 was first synthesized in the 1940se1950s by Schlesinger et al. [1] and soon attracted much attention [2]. Stable in dry conditions, it spontaneously (exothermically) hydrolyzes in the presence of water while generating hydrogen: BH 4 ðaqÞ þ 4H2 OðlÞ/BðOHÞ4 ðaqÞ þ 4H2 ðgÞ
(1)
In the 1950se1960s, most of the studies focused on improving the knowledge of the kinetics and mechanisms of
this reaction (1) [3e14]. Today, interest in the hydrolysis of NaBH4 is focused on applications, such that the objective is typically to accelerate H2 release with the aid of metal catalysts [15e22]. The spontaneous hydrolysis is less investigated and it is even avoided by virtue of stabilizing agents such as NaOH [23]. The spontaneous hydrolysis of NaBH4 is an acid-catalyzed reaction (in-situ formed H3Oþ ions) [5,6]. Mochalov et al. [10,11] proposed a detailed mechanism based on the substitution of BeH bonds by BeOH bonds. The hydrolysis would be a multistep process involving hydroxyborate intermediates
* Corresponding author. Tel.: þ33372448403; fax: þ33372440618. E-mail address: [email protected] (U.B. Demirci). 0360-3199/$ e see front matter ª 2010 Professor T. Nejat Veziroglu. Published by Elsevier Ltd. All rights reserved. doi:10.1016/j.ijhydene.2010.10.055
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BH4z(OH) z (with z from 1 to 3). For example, the reactions (2) and (3) show the H3Oþ-catalyzed hydrolysis of BH 4 and the formation of the first intermediate BH3(OH) (i.e. z ¼ 1), respectively: þ BH 4 ðaqÞ þ H3 O ðaqÞ/BH3 ðaqÞ þ H2 OðlÞ þ H2 ðgÞ
(2)
BH3 ðaqÞ þ OH ðaqÞ/BH3 OH ðaqÞ
(3) BH 4
The formation of BH3 from involves an activated or even BH5 (or H2BH3) complex, which is HþBH 4 [5,6,8,10,11,14]. The reactions (2) and (3) are both rate determining steps. The spontaneous hydrolysis of NaBH4 depends on experimental conditions, e.g. temperature, solution pH and ratio H2O/NaBH4 [5,6,24]. For example, a well-known empirical relation (4) shows the dependence of the half-reaction time t1/2 (in min) on the pH and the temperature T (in K) [10,11]: log t1=2 ; NaBH4 ¼ pH ð0:034 T 1:92Þ
(4)
To illustrate, based on this equation, hydrolysis should achieve a conversion of 50% in 134 min at pH 10 at 15 C, 28 min at pH 10 at 35 C, 13271 min (i.e. >9 days) at pH 12 at 15 C and 2773 min (i.e. about 2 days) at pH 12 at 35 C. This equation was established in the following experimental conditions: 7 < pH < 10 and 15 < T ( C) < 35. Even though it does not cover pH ranges >10 and temperatures >35 C, it is an indicative tool that gives a satisfactory approximation of the effects of both pH and temperature [15]. From this brief literature survey, it is clear that the kinetic parameters reported so far were obtained at specific experimental conditions. Most of the studies were carried out with aqueous NaBH4 solutions at low concentrations (e.g. 0.01 mol L1), which besides could be buffered [10,11,24]. Another point is that the kinetic data were calculated on the few first percentages of NaBH4 converted [10,11,14]. Accordingly, we studied the kinetics of the spontaneous hydrolysis of NaBH4 at higher concentrations (up to about 6 mol L1) and high temperatures (30e80 C). Furthermore, the solutions were not buffered. Our objective was to contribute to a better understanding of the reaction in different (more severe) experimental conditions while attempting to identify the reaction intermediates upon the evolution of 1 equiv H2. Note that such severe conditions are much closer to the conditions that should be used in technological applications, even though catalysts will be used to enhance kinetics of H2 release.
2.
225
overall system was considered as a non-steady state slurry batch reactor. The H2O/NaBH4 molar ratio was equal to 9 (18.9 wt% of NaBH4 or about 6.18 mol L1). To determine the reaction order with respect to the initial NaBH4 concentration, molar ratios of 13 (13.9 wt%), 18 (10.4 wt%), 27 (7.2 wt%), 36 (5.5 wt%) and 90 (2.3 wt%) were considered by varying the NaBH4 mass. The NaBH4 content was below its solubility limit (55 g per 100 g H2O) on the temperature range we used, namely 30e80 C. Prior to any experiment, the NaBH4 solid was transferred into the tube. In a typical experiment, water was injected into the reactor with a needle placed directly inside the bed of NaBH4. The injection needle was immediately removed following water injection. An automated burette was used to control the amount of water. The H2 produced was collected and evacuated through a second needle, also inserted in the tube through the stopper and connected to an inverted water-filled graduated cylinder. The H2 volumes were measured as a function of time by video monitoring the displaced water as the reaction proceeded. The hydrolysis tests were stopped when the H2 generation rate (HGR) was lower than 0.0001 mL s1, which is equivalent to a maximum conversion value of 50e60%, depending on the experiment conditions (the HGRs decreased because of the pH increase as suggested by Equation (4)). A thermal probe placed inside the NaBH4 bed measured the temperature during the reaction. The NaBH4 conversion, 3 (%), was calculated as being the ratio of the experimental volume of H2 released at a given time to the expected maximum volume. The HGR (r, mL s1) was obtained by linear regression of the hydrogen evolution curves. Because HGR changed with the reaction conversion for spontaneous hydrolysis, it was determined by taking the slope of the curve at a given NaBH4 conversion range defined by 3 3%. For example, at 3 ¼ 20%, the range 17e23% was considered and the data over this range were linearized. From the as-obtained line (with R2 0.997), the HGR (i.e. the line slope) was calculated. The hydrolysis by-products were analyzed by X-ray diffraction (Bruker D5005 powder diffractometer, CuKa radi˚ ), Fourier transform infrared spectroscopy ation (l ¼ 1.5406 A
Experimental
Commercial NaBH4 (ACROS organics, 98% purity, powder, n CAS: 16940-66-2, average particle size of 200 mm) was used as received. It was stored and handled in an argon-filled glove box in order to avoid moisture contact and subsequent hydrolysis. Distilled water was purged with argon to remove oxygen. The hydrolysis experiments were performed in a 20 mL tube closed by a silicon stopper and placed in a thermostatic bath. For the determination of the kinetic parameters the
Fig. 1 e Hydrogen evolutions obtained at various temperatures (30, 50, 60, 70 and 80 C) for an H2O/NaBH4 molar ratio of 9.
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Table 1 e HGRs (r, mL sL1) obtained for the H2O/NaBH4 molar ratio 9 (18.9 wt%), at various temperatures and determined for various conversions (3 ± 3%). HGR (r, mL s1) at various conversions
Temperature ( C) 5% 30 50 60 70 80
0.006 0.03 0.08 0.24 0.53
10% 0.003 0.016 0.03 0.10 0.33
15% 0.0025 0.009 0.021 0.05 0.15
20% 0.0015 0.006 0.014 0.04 0.14
25% 0.001 0.004 0.008 0.03 0.07
30% 0.0005 0.004 0.006 0.021 0.06
35% a
e 0.002 0.005 0.012 0.043
40% a
e 0.0012 0.003 0.007 0.04
45% a
e 0.0008 0.003 0.007 0.024
50% ea 0.0008 0.0015 0.005 0.017
a The hydrolysis was stopped because of HGRs <0.0001 mL s1.
(FT-IR, FTIR Nicolet 380) and 11B nuclear magnetic resonance (either Bruker Avance500 MHz for solid compounds or Bruker DRX 500 400 MHz for ions in aqueous solution). 11B MAS NMR spectra of solid were recorded at 11.7 T on a Bruker Avance500 wide-bore spectrometer operating at 160.49 MHz, using
a Bruker 4 mm probe and a spinning frequency of the rotor of 14 kHz. Spectra were acquired using a spin-echo q-t-2q pulse sequence with q ¼ 90 to overcome problems of probe signal and recycle delay of 1s was used. Chemical shifts were referenced to BF3(OEt)2 (d ¼ 0 ppm).
Fig. 2 e Plots of ln(r) as a function of TL1 for various conversions, i.e. (a) 5%, (b) 10%, (c) 30%, and (d) 50%, and (e) apparent activation energies as a function of conversion.
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3.
Results and discussion
3.1.
Kinetic study
The NaBH4 spontaneous hydrolysis was studied from 30 to 80 C. The molar ratio H2O/NaBH4 was kept at 9. Fig. 1 shows the H2 evolutions obtained at 30, 50, 60, 70 and 80 C. As expected, the higher the temperature, the more the hydrolysis extent (in terms of HGRs and conversions at a given time). This is also shown by the data given in Table 1. Moreover, the curves had similar profiles; for a given temperature, the HGRs decreased with time, i.e. with the reaction conversion. As suggested by Equation (4), this was due to an increase of the solution pH because of the formation of the basic hydroxyborates BH4z(OH) z [10,11]. In fact, it was observed that at conversions of around 50% the HGRs were very low, at this point the experiments were stopped. The data in Table 1 were then exploited to plot InðrÞ ¼ f ð1=TÞ and the apparent activation energy was calculated. This was done for HGRs obtained at various conversions, i.e. 5 3%, 10 3%, 15 3%, 20 3%, 25 3%, 30 3%, 35 3%, 40 3%, 45 3%, and 50 3%. For example, Fig. 2 (aed) depicts some of the as-obtained linear regressions. It is important to note that for a given conversion the solution characteristics (pH, products concentration and so on) should have been similar, with the only different parameter being the temperature. Fig. 2(e) shows the apparent activation energies as a function of the conversion. The apparent activation energy is 98 10 kJ mol1 over the range 30e80 C (lit. [4e9,12,13]: 100 35 kJ mol1 at low conversions, low temperatures, i.e. < 35 C, and for buffered solutions). The reaction order versus the initial NaBH4 concentration was determined for various H2O/NaBH4 molar ratios (i.e. 9, 13, 18, 27, 36 and 90; 6.18, 4.27, 3.09, 2.08, 1.54 and 0.63 mol L1; 18.9, 13.9, 10.4, 7.2, 5.5 and 2.3 wt%) at 80 C. The H2 evolution curves were analyzed at both 20 3% and 40 3% (Fig. 3(a,b)); at these conversions, the H2 evolution was quite linear, and the HGRs were deduced by linear regression. Table 2 reports the as-determined HGRs. Three observations stood out: (i) the HGRs were quite similar, i.e. about 0.17 mL s1, at a conversion of 20% and irrespective of the molar ratios; (ii) a similar observation was reported at a conversion of 40% (HGRs of 0.04e0.07 mL s1); (iii) at a given reaction time, the lower the H2O/NaBH4 molar ratio (i.e. the higher the NaBH4 concentration), the higher the volume of H2 released. The HGR data were then used to plot their logarithm as a function of the logarithm of the initial NaBH4 concentration (Fig. 3(c)). A reaction order versus the initial NaBH4 concentration of zero was obtained at 20 3% and 40 3%. In other words, the spontaneous hydrolysis is independent on the initial NaBH4 concentration for concentrations up to 6.18 mol L1. To the best of our knowledge, reaction orders versus the initial NaBH4 concentration in these experimental conditions have not been reported to date, which does not make any comparison possible. It is noteworthy that a reaction order of 0 was obtained for lower conversions and temperatures and for buffered solutions [4e9,12,13]. To summarize the results discussed heretofore, it mainly stands out that the apparent activation energy of the
227
spontaneous hydrolysis determined over the range 30e80 C is 98 10 kJ mol1 and is consistent with the energies obtained at lower temperatures. Given that the apparent activation energy determined experimentally includes the energies of every occurring reaction, one may assume that all of the / BH3(OH), BH3(OH) / BH2(OH) reactions BH 4 2, BH2(OH)2 / BH(OH) 3 and BH(OH)3 / B(OH)4 take place simultaneously. Another assumption might be that the reac tion BH 4 / B(OH)4 preferentially takes place and predominates. To highlight this, the hydrolysis slurry was analyzed upon the evolution of 1 equiv H2.
3.2.
Effect of pH on hydrolysis
According to Mochalov et al. [10,11], the solution pH has a significant effect on the spontaneous hydrolysis kinetics. We therefore followed the pH evolution in the course of the spontaneous hydrolysis of NaBH4 at 30 C (Fig. 4(a)). The pH first increased quite rapidly from about 9 to 10.8 (in about 16 min).
Fig. 3 e Hydrogen evolution at conversions of 20% (a) and 40% (b) for various H2O/NaBH4 molar ratios, at 80 C, and (c) determination of the reaction order versus the NaBH4 concentration at 80 C and for conversions of 20 ± 3% and 40 ± 3%.
228
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Table 2 e HGRs (r) for various molar ratios and for conversions of 20 ± 3% and 40 ± 3%. H2O/NaBH4 molar ratio
[NaBH4]0 (mol L1)
NaBH4]0 (wt%)
Conversion of 20 3% 1
90 36 27 18 13 9
0.63 1.54 2.08 3.09 4.27 6.18
2.3 5.5 7.2 10.4 13.9 18.9
Conversion of 40 3%
r (mL s )
Dr
r (mL s1)
Dra
0.20 0.16 0.16 0.16 0.17 0.14
0.06 0.05 0.03 0.02 0.04 0.02
0.07 0.07 0.08 0.07 0.06 0.04
0.03 0.03 0.02 0.01 0.01 0.05
a
a With Dr, the experimental error on r.
Then, the pH still increased but less rapidly. Actually the pH increase per unit of time was less and less high as the hydrolysis proceeded. It increased up to 11.7 after 30000 s of hydrolysis (i.e. >8 h). The profile of the pH evolution was similar to that of the H2 evolution. This clearly suggests that the solution pH strongly influences the hydrolysis kinetics. Nevertheless, the apparent activation energy is constant within the conversion range 0e50% and the hydrolysis is independent on the initial NaBH4 concentration. Hence, one may assume that the pH has no effect on the reaction mechanisms. The NaBH4 spontaneous hydrolysis is a complex reaction involving at least 5 different reaction species, from BH 4 to B(OH)4 , which play an important role in the evolution of the pH solution. Mochalov et al. suggested a stepwise hydrolysis of NaBH4 and showed that the half-reaction times of NaBH3(OH), NaBH2(OH)2 and NaBH(OH)3 were lower than that of NaBH4 [10,11]:
Fig. 4 e Evolution of the solution pH (experimental and calculated) as a function of time for an H2O/NaBH4 molar ratio of 9 at 30 C (a) and difference between the experimental and calculated pH values (b).
log t1=2 ; NaBH3 ðOHÞ ¼ pH ð0:027 T 0:357Þ
(5)
log t1=2 ; NaBH2 ðOHÞ2 ¼ pH ð0:027 T 0:384Þ
(6)
log t1=2 ; NaBHðOHÞ3 ¼ pH ð0:024 T 4:00Þ
(7)
The equations (4)e(7) were established under the following experimental conditions: 7 < pH < 10, 15 < T ( C) < 35, and buffered solutions. The rate constants of the reactions / BH2(OH) BH 4 / BH3(OH) , BH3(OH) 2 , BH2(OH)2 / BH (OH)3 and BH(OH)3 / B(OH)4 were found as 5.3 107, 3.6 1011, 10.1 107, 5.7 1010 mol1 min1, respectively. A
Fig. 5 e XRD patterns of the solids recovered at 5, 10, 15, 20, 25% of conversion from top to bottom (a) and XRD pattern of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (b); the stars show the peaks attributed to NaBH4 (ICDD 00-009-0386) while the others are assigned to NaB (OH)4 (ICDD 01-081-1512); the broad peaks within the range 2q [ 13 e28 is due to the amorphous film used to prevent the sample from moisture.
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constant rate of 5.2 107 mol1 min1 was also determined for the reaction BH 4 / B(OH)4 . In fact, BH4 can hydrolyze according to two routes: a stepwise one via the into B(OH) 4 formation of BH3(OH) or a direct one, the rate constants of the respective reactions being equivalent. It has been also reported that the concentrations of the BH3(OH) and BH (OH) 3 ions were small (0.03e0.05 mol%) in the course of the hydrolysis whereas the concentration of BH2(OH) 2 conversion w26 mol%. Molina Concha et al. [25] showed the formation of BH3(OH) through heterogeneous hydrolysis of NaBH4 and the presence of both BH3 and BH2 as stable intermediates by in-situ IR measurements. According to Ruman et al., it should be possible to synthesize BH3(OH) through the hydrolysis under ambient conditions of one BH 4 by one H2O in an excess of THF [26]. Its formation was evidenced by 1H NMR and the spectrum was similar to that of NaBH4 as it showed similar BeH coupling region but it was different in the region containing the BH resonances. An interesting point in Ruman et al’s conclusion
229
is that they assumed all of the BH 4 ions underwent a one-step hydrolysis and thus all led to the formation of BH3(OH). Unlike Mochalov et al’s studies [10,11], the NaBH4 solution was not buffered. If the discrepancies in the experimental conditions are not considered, Ruman et al’s work is in a way in contradiction with Mochalov et al’s. In order to highlight the discrepancy reported above, we tentatively calculated a theoretical pH evolution on the basis of the pKa of the couple B(OH)3/B(OH) 4 , which is 9.2 [27]. This couple is involved in the stepwise hydrolysis of NaBH4 [10,11]: þ BðOHÞ3 ðaqÞ þ H2 OðlÞ/BðOHÞ 4 ðaqÞ þ H ðaqÞ
(8)
Even though considering B(OH)3 as the only, first acid will overestimate the pH, it was assumed that the pH evolution was only dependent on the formation of B(OH) 4 from B(OH)3 and the following relation was applied, with B(OH)3 being a weak acid: h i pH ¼ 7 þ 0:5 pKa þ 0:5log BðOHÞ 4
(9)
Fig. 4(a,b) shows the calculated evolution. Three observations stood out: (i) the profile of the calculated evolution was quite similar to that of the experimental one; (ii) consistent with the acid/base couple considered and its pKa, the calculated pH was overestimated (DpH of about 1) at the beginning of the hydrolysis (1e3% of conversion) but interestingly DpH decreased to 0.1 at 17e18%; (iii) for conversions higher than 22% (i.e. t > 20000 s), the calculated pH was almost equal to the experimental pH. Observations (ii) and (iii) suggest therefore that the acid/base couple(s) involving the first hydrolysis intermediates such as e.g. BH3(OH) would have a pKa value lower than that of the couple B(OH)3/B(OH) 4 , and that the solution pH would be controlled by B(OH) 4 at conversion
Fig. 6 e IR spectra of NaBH4 and solids recovered at 5, 10, 15, 20, 25% of conversion from top to bottom (a) and IR spectrum of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (b).
Fig. 7 e Solid 11B RMN spectrum of the product obtained by the reaction of 1 equiv. H2O with 4 NaBH4 (a) and proton coupled 11B RMN spectrum of this product in aqueous (90 wt% H2O D 10 wt% D2O) solution.
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22%. In other words, it is important to identify the reaction intermediates that are involved in the beginning of the hydrolysis.
3.3.
Hydrolysis intermediates identification
To make clear the pH evolution, we focused our work on characterizing the hydrolysis intermediates at conversions up to 25%. The following tests were performed. The NaBH4 spontaneous hydrolysis was stopped at various conversions, i.e. 5, 10, 15, 20 and 25%, by dipping the tube in liquid air (<140 C). The slurry froze within 10 s, which is negligible in comparison to the timescale required for completion of spontaneous NaBH4 hydrolysis. The reaction was thus effectively quenched. The frozen slurry was dried under vacuum at
48 C for 48 h (to remove all water and prevent any hydrolysis when the sample was warmed) and then handled in an argon-filled glove box. The solids were analyzed by XRD (Fig. 5(a)). NaBH4 (ICDD 00-009-0386) was clearly identified at each conversion. It was the sole phase identified whereas a likely peak at around 2q ¼ 32 , characteristic of borates, had been especially regarded. No borate was found, which could be due an amorphous state [28]. Since the NaBH4 hydrolysis products have been reported as being stable up to 170 C [29], our solids were heated at 150 C for 4 h under Ar but such process was unsuccessful in crystallizing the borates. As a further test, the solids were then analyzed by IR (Fig. 6(a)). Three observations stood out. The spectrum recorded for NaBH4 was confirmed [30,31]. The NaBH4 print, i.e. nmax/ cm1 1114, 2220, 2293, 2341, 2357 and 2438 (BeH), was clearly
Fig. 8 e Proton decoupled (top) and coupled (bottom) 11B NMR spectra of the hydrolysis by-product recovered at a conversion of 25%.
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visible in each of the other spectra, showing that most of the reactant had not reacted. New bands, at nmax/cm1 500 d (BeO), 700e900 n(BeO), 1200 d(BeOeH), 1330 n(BeO), 1550e1650 d(HeOeH) and 3300e3500 n(OeH) [25,32], were observed for the partly hydrolyzed NaBH4, their respective contribution appearing to increase with the NaBH4 conversion. From such observations, it is in fact quite difficult to assert the formation of one hydroxyborate compound over another one. Tentatively our IR spectra might be compared to spectra previously reported and some/most of the bands might be assigned to e.g. NaBH3(OH) [29], NaBO2 [33,34], NaB(OH)4 [35], or NaBO2.2H2O [29,33,36]. In Ref. [25], bands at 970 and 1180 cm1 were attributed to BH2 and BH3, respectively. Nevertheless, none clearly matched and one may propose that the IR spectra indicate the formation of a tetra-coordinated boron but the presence of both BeH (especially from unreacted NaBH4) and BeOeH stretching modes impedes a clear identification of BH4z(OH) z or B(OH)4 . To succeed in identifying the compounds NaBH4z(OH)z that could be found in the hydrolysis medium at a conversion of 25%, we attempted to synthesize NaBH3(OH) according to the procedure recently reported by Ruman et al. [26]. Typically 1 equiv H2O per 1 NaBH4 were mixed in extra dry THF under Ar and the slurry was aged for 40 days while being stirred. Then, it was dried under vacuum for 24 h, transferred in the glove box, ground and analyzed by XRD. The XRD pattern is given in Fig. 5(b). The peaks could be assigned to both NaBH4 and NaB (OH)4. It is noteworthy that, some synthesis attempts showed unreacted NaBH4 and the presence of peaks assigned to NaB (OH)4. Nevertheless, as NaBH3(OH) is not listed in our XRD patterns database, we did not discard its formation. The synthesized product was then analyzed by IR (Fig. 6(b)). The spectrum supported the formation of NaB(OH)4, matching IR bands reported elsewhere [29]. As our observations contradict those of Ruman et al. [26], we completed the analysis by performing 11B solid state MAS NMR. The spectrum is given in Fig. 7(a). Two signals were observed at 2.2 and 41.8 ppm (difference of 44.0 ppm), corresponding to NaB(OH)4 and NaBH4, respectively [37]. Furthermore, we prepared an aqueous (90 wt% H2O and 10 wt% D2O) solution of this compound to be analyzed by 11B NMR. Water was injected into the sample-filled vial only few seconds before the analysis. Despite the H2 evolution, a spectrum not decoupled from proton in Fig. 7(b) was obtained. Once more, two signals were observed at 5.1 and 38.5 ppm (difference of 43.6 ppm), cor responding to B(OH) 4 and BH4 , respectively. The spectrum consisted of a singlet (NaB(OH)4) and a quintet (NaBH4). The JBH spinespin coupling constant for BH 4 was 80.6 Hz (lit. [12,13,26]: 82.0, 81.0 Hz). By focusing on d z 10 ppm, it was possible to distinguish within the background noise a negligibly small signal, likely a quartet that could be assigned to BH3(OH), which could have formed by hydrolysis of the dissolved solid. As a result, XRD, IR and NMR strongly suggested the formation of NaB(OH)4 rather than that of NaBH3(OH). In other words, when 1 equiv H2O is intended to hydrolyze 1 NaBH4, it is in fact 25% of the NaBH4 that hydrolyzes whereas 75% does not hydrolyze. This implies that at a conversion of 25% the specie controlling the solution pH is NaB(OH)4. 11 B RMN was also used to analyze the slurry recovered at a conversion of 25 1%. In this case, despite a low H2 evolution,
231
Fig. 9 e Spontaneous hydrolysis of NaBH4 and intermediates effect on pH over the conversion range 0e25%.
the 11B NMR collecting was done for the solution to keep the slurry state unchanged. The temperature was set at 5 C to minimize the H2 evolution. Fig. 8(a) shows the proton decoupled spectrum. It showed 2 signals at 38.4 and 4.8 ppm (difference of 43.2 ppm), characteristic of BH 4 and B(OH)4 [37]. A very small signal was observed at 9.9 ppm, ascribed to BH3(OH). Assuming a conversion of NaBH4 of 25%, the integration gave a relative content in BH3(OH) of 0.4 mol%, rather in agreement with Mochalov et al.’s observations [10,11]. No other hydroxyborate was detected. Guella et al. [37] investigated the hydrolysis intermediates for Pd-catalyzed hydrolysis of NaBH4. In a typical 11B proton decoupled NMR spectrum, no compound other than BH 4 and B(OH)4 was present, what was attributed to the low symmetry of the hydroxyborates structure leading to a broadening of 11B signals. Another reason could be the presence of the Pd catalyst, which accelerated the hydrolysis and so decreased the lifetime of very short-lived intermediates such as BH3(OH) [10,11]. It is noteworthy that, for uncatalyzed hydrolysis of NaBH4, the main thermodynamically stable borate is NaB(OH)4 in experimental conditions different from ours [28,38,39]. Fig. 8(b) shows the 11B spectrum not decoupled from proton, which is rather consistent with that reported by Gardiner and Collat [12,13]. The spectrum consisted of a singlet (NaB(OH)4), a quartet (NaBH3(OH)) and a quintet (NaBH4). The JBH spinespin coupling constants were calculated. The JBH value for NaBH3(OH) was 87.6 Hz (lit. [12,13,26,37]: 87.5, 82.0, 81.0 Hz). The JBH value for NaBH4 was 80.6 Hz (lit. [12,13,26]: 82.0, 81.0 Hz). The presence of the hydroxyborates BH2(OH) 2 and BH(OH)3 was especially scrutinized by exploring the shifts within the range defined by the signals of the borates BH3(OH) and B(OH) 4 , a shielding being expected. As shown in Fig. 8, no BH2(OH) 2 or BH(OH)3 were detected. In other words, when NaBH3(OH) was synthesized by reacting 1:1H2O with NaBH4 or by quenching the NaBH4 hydrolysis at a conversion of 25%, mainly NaB(OH)4 was formed with traces of NaBH3(OH).
4.
Conclusion
The spontaneous hydrolysis of NaBH4 was studied under experimental conditions that are much more severe than those
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reported to date (i.e. T ¼ 30e80 C, [NaBH4]0 ¼ 0.63e6.19 mol L1, solutions not buffered). Under these conditions, an apparent activation energy of 98 10 kJ mol1 and a reaction order versus the initial NaBH4 concentration of 0 were found. Besides, a parallelity in the evolution of the solution pH with that of the NaBH4 conversion was observed. The hydrolysis intermediates at conversions up to 25% were recovered by dipping the hydrolysis slurry in liquid air (<140 C) and then dried under vacuum. The as-obtained solids were analyzed by XRD, IR and 11B NMR. The three characterization techniques gave the same results; they evidenced the formation of 25 mol% of B(OH) 4 and 75 mol% of BH 4 at a conversion of 25%. This was consistent with the kinetic data. It was besides confirmed that BH3(OH), which is the first hydrolysis intermediate, is a very short-lived intermediate. Traces of it were detected. No other hydroxyborate was detected by 11B NMR. As a result, one concludes that (i) the direct reaction preferentially occurred under our experiBH 4 / BH3(OH) mental conditions and (ii) the B(OH) 4 anions controlled the solution pH at conversions 22%. Fig. 9 depicts schematically our observations. Notably, our study has confirmed that at basic pH, the H2 release through spontaneous hydrolysis is harshly hindered.
Acknowledgements The present study was mainly funded by the ‘Cluster Energies’ of ‘Re´gion Rhoˆne-Alpes, France’. It was also funded by the ANR project BoraHCx. The authors are grateful to Mrs. Caroline TOPPAN (Univ. Lyon 1, CCRMN) and Mr. Olivier Majoulet (Univ. Lyon 1, LMI) for some of the 11B NMR spectra and the subsequent valuable discussions.
references
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