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Cite this: Chem. Commun., 2014, 50, 707 Received 24th September 2013, Accepted 6th November 2013
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Base-free hydrogen generation from methanol using a bi-catalytic system† Ange `le Monney, Enrico Barsch, Peter Sponholz, Henrik Junge, Ralf Ludwig and Matthias Beller*
DOI: 10.1039/c3cc47306f www.rsc.org/chemcomm
A bi-catalytic system, in which Ru-MACHO-BH and Ru(H)2(dppe)2 interact in a synergistic manner, was developed for the base-free dehydrogenation of methanol. A total TON > 4200 was obtained with only trace amounts of CO contamination (o8 ppm) in the produced gas.
Hydrogen is attracting increasing attention as an alternative energy carrier and is considered as a potential fuel of the future.1 It is therefore conceivable that a ‘‘hydrogen economy’’2–6 may solve the problems related to the use of fossil fuels whose reserves are depleting. Hydrogen can be used as an energy source for fuel cells or as a direct fuel for combustion with the release of water as the only side product. An important challenge towards the implementation of a hydrogen economy is the safe, practical and economical storage of the produced hydrogen.7 Unfortunately, the compression or liquefaction of H2 is costly and requires substantial energy. Alternatively, the chemical storage of hydrogen in liquid organic compounds and its release on demand hold great promise. In this respect, methanol is an ideal candidate due to its high hydrogen content (12.6% by weight) and its liquid state at room temperature. The so-called methanol-reforming process allows for the production of H2 from a mixture of methanol and water.8 However, the high temperatures needed (>200 1C) and the co-production of carbon monoxide render current methanol steam reforming processes unsuitable for applications in fuel cells which only tolerate very low CO concentrations. Recently, we described an aqueous phase methanol reforming process at temperatures below 100 1C using a homogeneous ruthenium catalyst with only traces of released carbon monoxide (o10 ppm).9 Hydrogen generation was observed with excellent catalyst turnover numbers (>350 000) and turnover frequencies
¨r Katalyse e.V. and der Universita ¨t Rostock, Leibniz-Institut fu Albert-Einstein-Strasse 29a, 18059 Rostock, Germany. E-mail: [email protected]; Fax: +49 381 1281 51113; Tel: +49 381 1281 113 † Electronic supplementary information (ESI) available. See DOI: 10.1039/c3cc47306f
This journal is © The Royal Society of Chemistry 2014
(4700 h 1) using ruthenium complexes bearing a cooperative PNP pincer ligand10–16 in the presence of base. From an environmental point of view and for practical applications, it would be advantageous to produce H2 from methanol under ¨tzmacher et al. neutral conditions. Parallel to our work, Gru described another ruthenium-based catalyst capable of dehydrogenating a methanol–water mixture without base with a TOF = 50 h 1.17 Here, we report an improved base-free methanol dehydrogenation using the ruthenium-based PNP pincer complex Ru-MACHO-BH18,19 (A1, see Table 1 for structure). This catalyst is comparable to Ru-MACHO (A0),20–22 used for the basemediated methanol reforming process, but lacks the chloride ligand and therefore does not require a base for its activation. Complex A1 showed comparable catalytic activity to A0 in the presence of a base (Table 1, entries 1 and 2). The addition of triglyme23 as solvent was found to increase the solubility of the catalyst (cf. ESI†) and the gas generation was improved, reaching a gas evolution rate of 149 mL h 1, corresponding to a Table 1
Base-free hydrogen generation from aqueous methanola
Catalyst loadingb Entry Cat. [mmol]/[ppm] Base
Solvent
1 2 3 4 5 6 7
None 113 None 133 Triglyme 150 Triglyme 1.7 Triglyme 17 Triglyme 37 Triglyme 61
A0 A1 A1 A1 A1 A1 A1
5/23 5/23 5/23 5/23 20/90 45/203 95/428
KOH KOH KOH None None None None
Gas evolution ratec [mL h 1] TON3h 2064 2392 2743 32 70 74 59
a
Reaction conditions: MeOH (9.0 mL), H2O (1.0 mL), solvent (4.0 mL) when indicated, base (80 mmol) when indicated, catalyst (5–95 mmol), Tset = 93.5 1C. The gas evolution was measured manually using burettes. b ppm relative to MeOH. c Average calculated over the first 3 hours.
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TOF > 900 h 1 after 3 hours (Table 1, entry 3). Although no activity was observed in the absence of base with 5 mmol of A1 (Table 1, entry 4), an increase in catalyst loading allowed the reaction to proceed. With 20 mmol of A1, the gas evolution was measured at an average rate of 17 mL h 1 (Table 1, entry 5). When the catalyst loading was increased to 45 mmol, a proportional increase in the gas evolution rate was observed reaching an average value of 37 mL h 1 (Table 1, entry 6). Finally, 95 mmol of A1 induced the highest gas evolution rate with 61 mL h 1 (Table 1, entry 7). The increase in catalyst loading is believed to significantly accelerate the decomposition of formic acid, an intermediate suspected of deactivating the catalyst. Even though the evolution of gas is increased by the use of higher catalyst loadings, the overall productivity of the catalytic system stayed low. In order to improve catalyst turnover numbers, the catalytic step that involves the dehydrogenation of formic acid needs to be accelerated in a different way. Therefore, we envisioned that the addition of a second catalyst capable of efficiently dehydrogenating formic acid would generate a bi-catalytic system that allows the reaction to proceed under neutral conditions at low catalyst loading. Although a similar approach was proposed by Cole-Hamilton in 1989 for the dehydrogenation of ethanol using a series of rhodium and ruthenium catalysts,24 only modest improvements in H2 production were achieved in certain cases. To our delight, the combination of Ru-MACHO-BH (A1) with Ru(H)2(dppm)2 (B1, Table 2) at a catalyst loading as low as 5 mmol per catalyst (i.e. 45 ppm total [Ru] relative to MeOH) was able to effectively dehydrogenate methanol under neutral conditions (Table 2, entry 1). Despite a rapid gas evolution at the beginning of the reaction with a turnover frequency of 239 h 1 after 1 hour, this catalytic system however slowly deactivated. Gratifyingly, when B1 was replaced with the less constrained Ru(H)2(dppe)2 (B2), a constant gas evolution was measured over 7 hours with an average TOF of 93 h 1 (Table 2, entry 2).25 Other variations of the bisphosphine ligand on the ruthenium catalyst B gave only a slower gas evolution rate (Table 2, entries 3–5). The combination of A1 with an iron catalyst decorated with the tetraphos26 ligand showed no activity in the dehydrogenation of methanol either as a defined (B6) or as an in situ prepared (B7) catalyst (Table 2, entries 6 and 7, respectively) even though these iron complexes were previously found to be highly active in the base-free dehydrogenation of formic acid.27 When catalyst A1 was replaced by A2 (Table 2, entry 8) or A3 (Table 2, entry 9) in combination with B2 similar activities were observed. Unlike previous observations,9 the substituents on the PNP ligand (phenyl in A1 and isopropyl in A2 and A3) did not influence considerably the catalytic activity of the system. The use of the iridium PNP complex A4 lowered the gas evolution rate (Table 2, entry 10) as previously observed in the catalytic dehydrogenation of alcohols.22,28 A Ru–arene complex bearing a triazolylidene ligand (A5), reported for its activity in base-free alcohol oxidation,29 was also tested here in combination with B2 but showed no activity (Table 2, entry 11). Interestingly, when catalyst A1 and catalyst B2 were used individually for the dehydrogenation of methanol, both of them
708 | Chem. Commun., 2014, 50, 707--709
ChemComm Table 2 Catalyst combinations for the base-free dehydrogenation of methanola
Entry
Cat. A
Cat. B
Gas evolution rateb [mL h 1]
TOF1hc [h 1]
TOF3h [h 1]
TOF7h [h 1]
1 2 3 4 5 6 7 8 9 10 11 12
A1 A1 A1 A1 A1 A1 A1 A2 A3 A4 A5 —
B1 B2 B3 B4 B5 B6 B7 B2 B2 B2 B2 B2
60 30 3.6 6.6 5.3 1.2 0.6 26 25 13 2.0 9.1
239 87 13 19 20 4 2 80 76 44 8
138 94 10 22 13 — 2 79 76 36 —
83 93 — — — — — 77 — — —
a Reaction conditions: MeOH (9.0 mL), H2O (1.0 mL), triglyme (4.0 mL), cat. A (5 mmol), cat. B (5 mmol), Tset = 93.5 1C. See ESI for graphical representation. b Average calculated over the first 3 hours. c Calculated using the theoretical value of 75% of H2 content of the gas phase. See ESI for details.
produced little gas evolution (Fig. 1). More specifically, complex B2 generated only about 30 mL of gas in 3 hours (Table 2, entry 12) while A1 was essentially inactive (Table 1, entry 4). However, when these two complexes are combined together they interact in a synergistic manner and are able to dehydrogenate methanol with an increased rate. A stepwise addition of the catalysts showed that the effective dehydrogenation of methanol is triggered by the addition of B2 to A1 or vice versa by the addition of A1 to B2 (Fig. 1). Remarkably, the volume of gas generated by the bi-catalytic system is significantly greater than the sum of the gas volumes produced by each catalyst separately. The possibility
Fig. 1 Gas evolution from methanol catalysed by A1 (blue line) and B2 (red line), individually and together (green line). Rapid gas evolution starts when the two catalysts are mixed together (green dotted lines).
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of a ligand exchange process between the catalysts was ruled out since neither the addition of dppe (Ph2PCH2CH2PPh2) to complex A1 nor the addition of the PNP ligand (HN(CH2CH2PPh2)2) to B2 altered the activity of the individual catalysts. This positive interaction between the catalysts however decreases at higher catalyst loading (cf. ESI†). This is tentatively explained by the low solubility of B2 at higher catalyst loading, preventing an optimal interaction between the catalysts. In order to assess the role of formic acid in this methanol reforming process, HCOOH was added to the reaction mixture. The bi-catalytic system A1/B2 (5 mmol of each catalyst) dehydrogenated quickly and completely the extra formic acid (0.5 mL) with a gas evolution rate of 464 mL h 1, corresponding to a TOF = 940 h 1 (cf. ESI†).30–32 A positive synergistic effect between the two catalysts was observed here as well, although to a lower extent. Catalyst B2 (5 mmol) on its own was able to decompose the added HCOOH with a gas evolution rate of 352 mL h 1. While formic acid decomposition by A1 at a catalyst loading of 5 mmol was slow (31 mL h 1), an increased loading (45 mmol) allowed for the efficient and complete HCOOH decomposition with a gas evolution rate of 306 mL h 1. When the added formic acid was fully converted, methanol dehydrogenation was carried out at the same rate as before the addition of HCOOH, suggesting no significant deactivation of the catalysts. These results support our initial hypothesis that either an increase in catalyst loading or the addition of a second particular catalyst can improve the rate of the methanol dehydrogenation by accelerating the decomposition of in situ formed formic acid. Further insights into the mechanism of the methanol dehydrogenation were provided by in situ infrared spectroscopy. During the reaction, a vibration band was detected at 1730 cm 1 and was attributed to formic acid since its shape and position were very similar to free formic acid under the same conditions (cf. ESI†). This signal appeared as soon as the two catalysts were added and its intensity continuously increased until it reached a steady-state after about 20 hours. Our catalytic system therefore seems to reach an equilibrium in which a certain amount of formic acid is present. The appearance of this band was accompanied by a negative contribution at 1650 cm 1 which we attributed to the consumption of water since this region of the spectrum hosts the frequency of the deformation vibration of water. The stepwise addition of A1 to B2 and of B2 to A1 was also followed by IR spectroscopy and led to comparable conclusions to those obtained from gas evolution measurements (cf. ESI†). Further experiments involving the addition of formic acid and formaldehyde supported the proposed reaction pathway via these two intermediates9 and suggested the release of the first molecule of hydrogen as a rate limiting step (cf. ESI†). Finally, a long term experiment was performed resulting in the production of a total gas volume of about 1400 mL. This corresponds to a hydrogen yield of 26% (relative to H2O) and a TON > 4200 (cf. ESI†). Importantly, less than 8 ppm of carbon monoxide was observed by gas phase GC measurements throughout the reaction course. In conclusion, we developed a base-free bi-catalytic system for the production of H2 from aqueous methanol at low temperature.
This journal is © The Royal Society of Chemistry 2014
Communication
The synergistic interaction between Ru-MACHO-BH and Ru(H)2(dppe)2 allows the reaction to proceed at low catalyst loading. The unprecedented mild reaction conditions and the low CO contamination of the produced gas render this system interesting for a combination with PEM fuel cells. Additional experiments using in situ IR spectroscopy supported the proposed reaction pathway through formaldehyde and formic acid.
Notes and references 1 J. O. M. Bockris, Science, 1972, 176, 1323. 2 J. O. M. Bockris, Int. J. Hydrogen Energy, 2013, 38, 2579–2588. ´, D. M. Minic ´, D. G. Minic ´ and J. G. Novakovic ´, 3 V. A. Blagojevic Hydrogen Economy: Modern Concepts, Challenges and Perspectives, 2012. 4 J. Andrews and B. Shabani, Procedia Eng., 2012, 49, 15–25. 5 N. Armaroli and V. Balzani, ChemSusChem, 2011, 4, 21–36. 6 G. W. Crabtree, M. S. Dresselhaus and M. V. Buchanan, Phys. Today, 2004, 57, 39–44. 7 P. Jena, J. Phys. Chem. Lett., 2011, 2, 206–211. 8 D. R. Palo, R. A. Dagle and J. D. Holladay, Chem. Rev., 2007, 107, 3992–4021. 9 M. Nielsen, E. Alberico, W. Baumann, H.-J. Drexler, H. Junge, S. Gladiali and M. Beller, Nature, 2013, 495, 85–89. 10 R. Langer, I. Fuchs, M. Vogt, E. Balaraman, Y. Diskin-Posner, L. J. W. Shimon, Y. Ben-David and D. Milstein, Chem.–Eur. J., 2013, 19, 3407–3414. 11 S. Musa, S. Fronton, L. Vaccaro and D. Gelman, Organometallics, 2013, 32, 3069–3073. 12 K.-N. T. Tseng, J. W. Kampf and N. K. Szymczak, Organometallics, 2013, 32, 2046–2049. 13 K. E. Allen, D. M. Heinekey, A. S. Goldman and K. I. Goldberg, Organometallics, 2013, 32, 1579–1582. 14 S. Smith, D. Spasyuk and D. G. Gusev, Angew. Chem., Int. Ed., 2012, 51, 2772–2775. 15 S. Schneider, J. Meiners and B. Askevold, Eur. J. Inorg. Chem., 2012, 412–429. 16 J. I. van der Vlugt and J. N. H. Reek, Angew. Chem., Int. Ed., 2009, 48, 8832–8846. 17 R. E. Rodrı´guez-Lugo, M. Trincado, M. Vogt, F. Tewes, G. Santiso¨tzmacher, Nat. Chem., 2013, 5, 342–347. Quinones and H. Gru 18 Takasago International Corp., PCT Int. Appl., WO2011048727A, 2011. 19 For another complex with a [Ru(H)(H–BH3)] motif, see: T. Ohkuma, ˜iz, G. Hilt, C. Kabuto and R. Noyori, J. Am. Chem. M. Koizumi, K. Mun Soc., 2002, 124, 6508–6509. 20 W. Kuriyama, T. Matsumoto, O. Ogata, Y. Ino, K. Aoki, S. Tanaka, K. Ishida, T. Kobayashi, N. Sayo and T. Saito, Org. Process Res. Dev., 2012, 16, 166–171. 21 Z. Han, L. Rong, J. Wu, L. Zhang, Z. Wang and K. Ding, Angew. Chem., Int. Ed., 2012, 51, 13041–13045. 22 M. Nielsen, H. Junge, A. Kammer and M. Beller, Angew. Chem., Int. Ed., 2012, 51, 5711–5713. 23 Triethylene glycol dimethyl ether. 24 D. Morton, D. J. Cole-Hamilton, I. D. Utuk, M. Paneque-Sosa and M. Lopez-Poveda, J. Chem. Soc., Dalton Trans., 1989, 489–495. 25 These optimal conditions were reproduced several times with a maximum error of 6%. 26 Tris-[(2-diphenylphosphino)ethyl]phosphine. ¨rtner, R. Jackstell, H. Junge, 27 A. Boddien, D. Mellmann, F. Ga P. J. Dyson, G. Laurenczy, R. Ludwig and M. Beller, Science, 2011, 333, 1733–1735. 28 M. Nielsen, A. Kammer, D. Cozzula, H. Junge, S. Gladiali and M. Beller, Angew. Chem., Int. Ed., 2011, 50, 9593–9597. 29 D. Canseco-Gonzalez and M. Albrecht, Dalton Trans., 2013, 42, 7424–7432. 30 I. Mellone, M. Peruzzini, L. Rosi, D. Mellmann, H. Junge, M. Beller and L. Gonsalvi, Dalton Trans., 2013, 42, 2495–2501. 31 M. Grasemann and G. Laurenczy, Energy Environ. Sci., 2012, 5, 8171–8181. ¨rtner, H. Junge and M. Beller, Top. Catal., 32 B. Loges, A. Boddien, F. Ga 2010, 53, 902–914.
Chem. Commun., 2014, 50, 707--709 | 709
Published on 07 November 2013. Downloaded by York University on 21/06/2014 08:50:20.
COMMUNICATION
Cite this: Chem. Commun., 2014, 50, 707 Received 24th September 2013, Accepted 6th November 2013
View Journal | View Issue
Base-free hydrogen generation from methanol using a bi-catalytic system† Ange `le Monney, Enrico Barsch, Peter Sponholz, Henrik Junge, Ralf Ludwig and Matthias Beller*
DOI: 10.1039/c3cc47306f www.rsc.org/chemcomm
A bi-catalytic system, in which Ru-MACHO-BH and Ru(H)2(dppe)2 interact in a synergistic manner, was developed for the base-free dehydrogenation of methanol. A total TON > 4200 was obtained with only trace amounts of CO contamination (o8 ppm) in the produced gas.
Hydrogen is attracting increasing attention as an alternative energy carrier and is considered as a potential fuel of the future.1 It is therefore conceivable that a ‘‘hydrogen economy’’2–6 may solve the problems related to the use of fossil fuels whose reserves are depleting. Hydrogen can be used as an energy source for fuel cells or as a direct fuel for combustion with the release of water as the only side product. An important challenge towards the implementation of a hydrogen economy is the safe, practical and economical storage of the produced hydrogen.7 Unfortunately, the compression or liquefaction of H2 is costly and requires substantial energy. Alternatively, the chemical storage of hydrogen in liquid organic compounds and its release on demand hold great promise. In this respect, methanol is an ideal candidate due to its high hydrogen content (12.6% by weight) and its liquid state at room temperature. The so-called methanol-reforming process allows for the production of H2 from a mixture of methanol and water.8 However, the high temperatures needed (>200 1C) and the co-production of carbon monoxide render current methanol steam reforming processes unsuitable for applications in fuel cells which only tolerate very low CO concentrations. Recently, we described an aqueous phase methanol reforming process at temperatures below 100 1C using a homogeneous ruthenium catalyst with only traces of released carbon monoxide (o10 ppm).9 Hydrogen generation was observed with excellent catalyst turnover numbers (>350 000) and turnover frequencies
¨r Katalyse e.V. and der Universita ¨t Rostock, Leibniz-Institut fu Albert-Einstein-Strasse 29a, 18059 Rostock, Germany. E-mail: [email protected]; Fax: +49 381 1281 51113; Tel: +49 381 1281 113 † Electronic supplementary information (ESI) available. See DOI: 10.1039/c3cc47306f
This journal is © The Royal Society of Chemistry 2014
(4700 h 1) using ruthenium complexes bearing a cooperative PNP pincer ligand10–16 in the presence of base. From an environmental point of view and for practical applications, it would be advantageous to produce H2 from methanol under ¨tzmacher et al. neutral conditions. Parallel to our work, Gru described another ruthenium-based catalyst capable of dehydrogenating a methanol–water mixture without base with a TOF = 50 h 1.17 Here, we report an improved base-free methanol dehydrogenation using the ruthenium-based PNP pincer complex Ru-MACHO-BH18,19 (A1, see Table 1 for structure). This catalyst is comparable to Ru-MACHO (A0),20–22 used for the basemediated methanol reforming process, but lacks the chloride ligand and therefore does not require a base for its activation. Complex A1 showed comparable catalytic activity to A0 in the presence of a base (Table 1, entries 1 and 2). The addition of triglyme23 as solvent was found to increase the solubility of the catalyst (cf. ESI†) and the gas generation was improved, reaching a gas evolution rate of 149 mL h 1, corresponding to a Table 1
Base-free hydrogen generation from aqueous methanola
Catalyst loadingb Entry Cat. [mmol]/[ppm] Base
Solvent
1 2 3 4 5 6 7
None 113 None 133 Triglyme 150 Triglyme 1.7 Triglyme 17 Triglyme 37 Triglyme 61
A0 A1 A1 A1 A1 A1 A1
5/23 5/23 5/23 5/23 20/90 45/203 95/428
KOH KOH KOH None None None None
Gas evolution ratec [mL h 1] TON3h 2064 2392 2743 32 70 74 59
a
Reaction conditions: MeOH (9.0 mL), H2O (1.0 mL), solvent (4.0 mL) when indicated, base (80 mmol) when indicated, catalyst (5–95 mmol), Tset = 93.5 1C. The gas evolution was measured manually using burettes. b ppm relative to MeOH. c Average calculated over the first 3 hours.
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TOF > 900 h 1 after 3 hours (Table 1, entry 3). Although no activity was observed in the absence of base with 5 mmol of A1 (Table 1, entry 4), an increase in catalyst loading allowed the reaction to proceed. With 20 mmol of A1, the gas evolution was measured at an average rate of 17 mL h 1 (Table 1, entry 5). When the catalyst loading was increased to 45 mmol, a proportional increase in the gas evolution rate was observed reaching an average value of 37 mL h 1 (Table 1, entry 6). Finally, 95 mmol of A1 induced the highest gas evolution rate with 61 mL h 1 (Table 1, entry 7). The increase in catalyst loading is believed to significantly accelerate the decomposition of formic acid, an intermediate suspected of deactivating the catalyst. Even though the evolution of gas is increased by the use of higher catalyst loadings, the overall productivity of the catalytic system stayed low. In order to improve catalyst turnover numbers, the catalytic step that involves the dehydrogenation of formic acid needs to be accelerated in a different way. Therefore, we envisioned that the addition of a second catalyst capable of efficiently dehydrogenating formic acid would generate a bi-catalytic system that allows the reaction to proceed under neutral conditions at low catalyst loading. Although a similar approach was proposed by Cole-Hamilton in 1989 for the dehydrogenation of ethanol using a series of rhodium and ruthenium catalysts,24 only modest improvements in H2 production were achieved in certain cases. To our delight, the combination of Ru-MACHO-BH (A1) with Ru(H)2(dppm)2 (B1, Table 2) at a catalyst loading as low as 5 mmol per catalyst (i.e. 45 ppm total [Ru] relative to MeOH) was able to effectively dehydrogenate methanol under neutral conditions (Table 2, entry 1). Despite a rapid gas evolution at the beginning of the reaction with a turnover frequency of 239 h 1 after 1 hour, this catalytic system however slowly deactivated. Gratifyingly, when B1 was replaced with the less constrained Ru(H)2(dppe)2 (B2), a constant gas evolution was measured over 7 hours with an average TOF of 93 h 1 (Table 2, entry 2).25 Other variations of the bisphosphine ligand on the ruthenium catalyst B gave only a slower gas evolution rate (Table 2, entries 3–5). The combination of A1 with an iron catalyst decorated with the tetraphos26 ligand showed no activity in the dehydrogenation of methanol either as a defined (B6) or as an in situ prepared (B7) catalyst (Table 2, entries 6 and 7, respectively) even though these iron complexes were previously found to be highly active in the base-free dehydrogenation of formic acid.27 When catalyst A1 was replaced by A2 (Table 2, entry 8) or A3 (Table 2, entry 9) in combination with B2 similar activities were observed. Unlike previous observations,9 the substituents on the PNP ligand (phenyl in A1 and isopropyl in A2 and A3) did not influence considerably the catalytic activity of the system. The use of the iridium PNP complex A4 lowered the gas evolution rate (Table 2, entry 10) as previously observed in the catalytic dehydrogenation of alcohols.22,28 A Ru–arene complex bearing a triazolylidene ligand (A5), reported for its activity in base-free alcohol oxidation,29 was also tested here in combination with B2 but showed no activity (Table 2, entry 11). Interestingly, when catalyst A1 and catalyst B2 were used individually for the dehydrogenation of methanol, both of them
708 | Chem. Commun., 2014, 50, 707--709
ChemComm Table 2 Catalyst combinations for the base-free dehydrogenation of methanola
Entry
Cat. A
Cat. B
Gas evolution rateb [mL h 1]
TOF1hc [h 1]
TOF3h [h 1]
TOF7h [h 1]
1 2 3 4 5 6 7 8 9 10 11 12
A1 A1 A1 A1 A1 A1 A1 A2 A3 A4 A5 —
B1 B2 B3 B4 B5 B6 B7 B2 B2 B2 B2 B2
60 30 3.6 6.6 5.3 1.2 0.6 26 25 13 2.0 9.1
239 87 13 19 20 4 2 80 76 44 8
138 94 10 22 13 — 2 79 76 36 —
83 93 — — — — — 77 — — —
a Reaction conditions: MeOH (9.0 mL), H2O (1.0 mL), triglyme (4.0 mL), cat. A (5 mmol), cat. B (5 mmol), Tset = 93.5 1C. See ESI for graphical representation. b Average calculated over the first 3 hours. c Calculated using the theoretical value of 75% of H2 content of the gas phase. See ESI for details.
produced little gas evolution (Fig. 1). More specifically, complex B2 generated only about 30 mL of gas in 3 hours (Table 2, entry 12) while A1 was essentially inactive (Table 1, entry 4). However, when these two complexes are combined together they interact in a synergistic manner and are able to dehydrogenate methanol with an increased rate. A stepwise addition of the catalysts showed that the effective dehydrogenation of methanol is triggered by the addition of B2 to A1 or vice versa by the addition of A1 to B2 (Fig. 1). Remarkably, the volume of gas generated by the bi-catalytic system is significantly greater than the sum of the gas volumes produced by each catalyst separately. The possibility
Fig. 1 Gas evolution from methanol catalysed by A1 (blue line) and B2 (red line), individually and together (green line). Rapid gas evolution starts when the two catalysts are mixed together (green dotted lines).
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of a ligand exchange process between the catalysts was ruled out since neither the addition of dppe (Ph2PCH2CH2PPh2) to complex A1 nor the addition of the PNP ligand (HN(CH2CH2PPh2)2) to B2 altered the activity of the individual catalysts. This positive interaction between the catalysts however decreases at higher catalyst loading (cf. ESI†). This is tentatively explained by the low solubility of B2 at higher catalyst loading, preventing an optimal interaction between the catalysts. In order to assess the role of formic acid in this methanol reforming process, HCOOH was added to the reaction mixture. The bi-catalytic system A1/B2 (5 mmol of each catalyst) dehydrogenated quickly and completely the extra formic acid (0.5 mL) with a gas evolution rate of 464 mL h 1, corresponding to a TOF = 940 h 1 (cf. ESI†).30–32 A positive synergistic effect between the two catalysts was observed here as well, although to a lower extent. Catalyst B2 (5 mmol) on its own was able to decompose the added HCOOH with a gas evolution rate of 352 mL h 1. While formic acid decomposition by A1 at a catalyst loading of 5 mmol was slow (31 mL h 1), an increased loading (45 mmol) allowed for the efficient and complete HCOOH decomposition with a gas evolution rate of 306 mL h 1. When the added formic acid was fully converted, methanol dehydrogenation was carried out at the same rate as before the addition of HCOOH, suggesting no significant deactivation of the catalysts. These results support our initial hypothesis that either an increase in catalyst loading or the addition of a second particular catalyst can improve the rate of the methanol dehydrogenation by accelerating the decomposition of in situ formed formic acid. Further insights into the mechanism of the methanol dehydrogenation were provided by in situ infrared spectroscopy. During the reaction, a vibration band was detected at 1730 cm 1 and was attributed to formic acid since its shape and position were very similar to free formic acid under the same conditions (cf. ESI†). This signal appeared as soon as the two catalysts were added and its intensity continuously increased until it reached a steady-state after about 20 hours. Our catalytic system therefore seems to reach an equilibrium in which a certain amount of formic acid is present. The appearance of this band was accompanied by a negative contribution at 1650 cm 1 which we attributed to the consumption of water since this region of the spectrum hosts the frequency of the deformation vibration of water. The stepwise addition of A1 to B2 and of B2 to A1 was also followed by IR spectroscopy and led to comparable conclusions to those obtained from gas evolution measurements (cf. ESI†). Further experiments involving the addition of formic acid and formaldehyde supported the proposed reaction pathway via these two intermediates9 and suggested the release of the first molecule of hydrogen as a rate limiting step (cf. ESI†). Finally, a long term experiment was performed resulting in the production of a total gas volume of about 1400 mL. This corresponds to a hydrogen yield of 26% (relative to H2O) and a TON > 4200 (cf. ESI†). Importantly, less than 8 ppm of carbon monoxide was observed by gas phase GC measurements throughout the reaction course. In conclusion, we developed a base-free bi-catalytic system for the production of H2 from aqueous methanol at low temperature.
This journal is © The Royal Society of Chemistry 2014
Communication
The synergistic interaction between Ru-MACHO-BH and Ru(H)2(dppe)2 allows the reaction to proceed at low catalyst loading. The unprecedented mild reaction conditions and the low CO contamination of the produced gas render this system interesting for a combination with PEM fuel cells. Additional experiments using in situ IR spectroscopy supported the proposed reaction pathway through formaldehyde and formic acid.
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