Bicarbonate Hydrogenation Catalyzed by Iron: How the Choice of

Mar 25, 2016 - J. 2013, 19, 11869], with special emphasis herein on the effects of the choice of solvent. By using density functional theory we have l...
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Bicarbonate Hydrogenation Catalyzed by Iron – How the Choice of Solvent can Reverse the Reaction Rocío Marcos, Liqin Xue, Rocío Sánchez-de-Armas, and Mårten S.G. Ahlquist ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.6b00071 • Publication Date (Web): 25 Mar 2016 Downloaded from http://pubs.acs.org on March 30, 2016

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Bicarbonate Hydrogenation Catalyzed by Iron – How the Choice of Solvent can Reverse the Reaction Rocío Marcos, Liqin Xue, Rocío Sánchez-de-Armas, and Mårten S. G. Ahlquist* Division of Theoretical Chemistry & Biology, School of Biotechnology, KTH Royal Institute of Technology, SE-10691 Stockholm (Sweden) ABSTRACT: Here we report a mechanism study of the hydrogenation of bicarbonate by tetradentate phosphines ironcomplexes. It is an extension of our recent study on the reverse reaction by the same type of complexes [Chem. Eur. J. 2013, 19, 11869], with special emphasis herein on the effects of the choice of solvent. By using density functional theory we have located the most plausible mechanism, and have found remarkable effects of the solvent on the reversibility of this reaction. We predict that the solvent used in experiment, MeOH, for the hydrogenation of bicarbonate to formate could be replaced to enhance the activity of the system. There is a direct correlation of the solubility of the base to favor or disfavor the hydrogenation of bicarbonate to formate.

KEYWORDS: bicarbonate hydrogenation, formic acid synthesis, tetradentate phosphine ligand, iron, density functional theory, solvent influence.

INTRODUCTION Carbon dioxide is an attractive C1 source compared to other molecules, such as carbon monoxide, since it is a low cost reactant, abundant, relatively nontoxic, and renewable. However, utilizing carbon dioxide still poses a challenge owing to its thermodynamic. In the past decades, the homogeneous catalytic hydrogenation of carbon dioxide and its derivatives has been widely investigated, due to the fundamental and practical applicability of the products of CO2 hydrogenation as excellent fuels or as raw materials of interest for many chemical industries.1 Due to the importance of this reaction, the catalytic hydrogenation of bicarbonate to formate has been explored extensively; however the vast majority of the catalytic systems known to be active for this reaction are based on precious noble metals such as rhodium,2 ruthenium,3 iridium,4 and palladium.5 More attractive in this context is the use of iron−based catalysts due to the abundance of this non-noble metal.6 The tricarbonyl (cyclopentadienone)iron complex, a precursor of the Knölker´s iron complex, was used by Zhu et al. for the hydrogenation of bicarbonate at low hydrogen pressures.7 To date, this is the only successful report in which the catalytic system does not contain phosphine ligands. Iron complexes that do contain phosphine ligands, specifically multidentate phosphine ligands, have shown the best performance for this reaction. In 2003, Tai et al. used bidentate phosphines in combination with FeCl3 or Fe(OAc)2 for the hydrogenation of carbon dioxide to formic acid,8 however the hydrogenation of bicarbonate was not tested by the authors.

The most active additive-free Fe-based catalytic system for this reaction reported is based on tetradentate phosphines as ligands.9 A pioneering work by the group of Beller described an iron branched complex [FeH(PP3)] (PP3)= (P(CH2CH2PPh2)3) generated in situ from Fe(BF4) and a tetradentate phosphine ligand. HCO2H

[Fe] THF or PC 80 °C, 20 h [Fe] MeOH

HCO3 - + H 2

H 2 + CO2

HCO2 - + H2 O

80 °C, 20 h PPh 2 P

HCO2-

Fe

H2

H PPh2

PPh2

HCO2-

H2

PPh2 PPh2 H P Fe H P H Ph 2

PPh2 PPh2 P

O

Fe

P Ph 2 H

O H

HCO2 CO2

HCO2 H

HCO3 -

PPh2 PPh2 CO2

P

Fe

H

CO2 + H2O

P Ph2 H

Figure 1. Proposed mechanisms for the dehydrogenation of formic acid (in blue) and for the hydrogenation of bicarbonate (in red), both processes are catalyzed by the complex + [FeH(PP3)] .

The system showed remarkable capability and was used for the reactions in both directions: 1) the dehydrogenation of formic acid and 2) the hydrogenation of both car-

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bon dioxide and bicarbonates (Figure 1).9a Both [FeH(PP3)]+ and [FeH(η-H2)(PP3)]+ species were observed under catalytic conditions.9a The second generation of these complexes and the most active and productive ironbased catalyst system reported so far was obtained by combining Fe(BF4)2 and tris(2diphenylphosphine)phenyl)phosphine).9b In this case the mechanistic investigation established the [FeF(PPhP3)]+ species as the pre-catalyst of the reaction. Recently an iron-pincer complex, trans [(tBuPNP)FeH2], reported by Milstein and co-workers was successfully applied to reduce carbon dioxide and bicarbonate at remarkably low hydrogen pressures (initial H2 pressure 6.2 bar).10 In addition, Gonsalvi and co-workers reported the use of a linear tetradentate phosphine (Rac-tetraphos-1 (P4)) ligand in combination with Fe(BF4) as catalyst of sodium bicarbonate hydrogenation to sodium formate with good activities.11 Although significant progress has been made to try to explain the reaction mechanisms with these highly active catalyst systems, challenges still remain.12 In our previous work we performed a theoretical investigation of iron−catalyzed dehydrogenation of formic acid based on the experimental results of iron-tetradentate phosphine complexes.13 The results showed that the reaction cycle proceeds at only one coordination site on the iron center, despite two coordination sites being available at the metal. These results are in line with our studies on other metal catalysts which show that the reaction only needs one reactive site on the metal.14 Since we studied the dehydrogenation of formic acid in the first report on these iron catalysts, we were intrigued by the reversibility of the reaction and wanted to understand its driving forces (Figure 1). In this study we find a remarkable effect from the solvent. Moreover, we find completely opposite effects from hydrogen bond donor solvents compared to non-hydrogen bond donor solvents. Two recent studies by us15 and Yang and co-workers16 have showed that the properties of metal hydrides when the polarity of the solvent is changed. Combined, these studies the importance in understanding the solvation effects when designing catalytic systems for CO2 reduction.

RESULTS Based on our previous work on the theoretical investigation of dehydrogenation of formic acid catalyzed by an iron complex [generated in situ from Fe(BF4)2 and a tetradentate phosphine ligand P(CH2CH2PPh2)3 (PP3)] and the experimental results obtained from Beller, Laurenczy and co-workers, we investigated the mechanism of the hydrogenation of sodium bicarbonate using exactly the same in situ-generated catalyst. Initially, iron−hydride complex int_1 was selected as the active complex for this investigation. Experimental results showed that this species has a triplet electronic configuration,17 in agreement with our previous calculations showing that the complex with triplet ground state (m=3) is more stable than the com-

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plex with the singlet state (m=1). All other calculated species of this cycle have been found to be more stable in their single state, unless otherwise specified. The free-energy profile involves four major steps (Figure 2). The cycle starts with the addition of the dihydrogen molecule to the precursor complex [FeH(PP3)]+ (int_1) to form Bianchini’s complex [FeH(η-H2)(PP3)]+ (int_2)17,18 through TS1_2 with a free energy barrier of 23.6 kcal mol-1. The calculated intermediate, int_2, is consistent with experimental crystal structure19 in which the hydride and dihydrogen ligand are positioned cis to each other. Subsequent deprotonation of the dihydrogencoordinated iron complex int_2 from a bicarbonate molecule occurs then to give dihydride complex int_4 and carbonic acid (H2CO3), which disproportionates to CO2 and H2O. The transition state of the deprotonation step (TS3_4) is at 27.5 kcal mol-1, suggesting that this step could be limiting the reaction. Insertion of CO2 into the Fe−H bond of the int_4 intermediate yields formatecoordinated iron complex int_5 via TS4_5, via a barrier at 27.7 kcal mol-1. The elimination of the formate molecule from int_5 generates the iron-hydride complex int_1 and completes the cycle. The free energy barrier of TS5_1 is 29.1 kcal mol-1. However, the releasing of formate can happen through a transition state with a triplet state (m=3), which is just slightly lower than the singlet state (m=1) by 1.3 kcal mol-1 at 27.8 kcal mol-1. The overall reaction is exergonic with a reaction free energy of -4 kcal mol-1, which is in good agreement with the experimental results.9 Our calculated results show the overall freeenergy barrier to be 27.8 kcal mol-1, which corresponds to the energy of the elimination step (TS5_1). However, the energy barriers of TS3_4 (deprotonation step) and TS4_5 (insertion of CO2 into Fe-H bond) are close to identical to that of TS5_1. Therefore, these three steps could be competing and is not possible to state, based on our calculated free energy profile, which step is the rate-determining step. Other possible mechanisms were also investigated. We calculated barriers for addition of dihydrogen, deprotonation of dihydrogen-iron complex and insertion of CO2 into Fe−H when the hydride is trans to bridgehead phosphorus (Figure 3). The transition state with the highest energy on this reaction pathway (ATS4_5) is 10.6 kcal mol1 over the highest points in the first mechanism in Figure 2.The calculated results clearly show that the favored reaction cycle proceeds at only one coordination site on the iron metal center and a hydride should be trans to the central phosphorus.13 In one of the slow steps in Figure 2 CO2 is directly inserted into the Fe−H bond, leading to formation of the Obound formate complex int_5 directly. We also identified another possible path for the reaction of CO2 with the hydride complex, where the initial intermediate is an Hbound formate complex. This step is followed by an isomerization step, where the formate finally coordina-

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+

Figure 2. Favored free energy profile calculated starting with the precursor complex [FeH(PP3)] where the hydride occupies the trans position to the central phosporous at the M06 level. The relative solvation corrected Gibbs free energies are given in kcal -1 mol . The selected bond distances and angles are given in angstroms and degrees, respectively.

tes with one of the oxygen atoms to the metal instead of the hydrogen. The free energy barriers for this mechanism (TS4_6 and TS6_5) are calculated at 25.2 and 27.5 kcal mol-1, respectively (see Figure 4), indicating that both pathways are possible for this reaction. All the calculations showed in the Figures 2-4 were carried out using solvation field parameters to describe methanol, the solvent used in the experiments reported by the research group of Beller.9a Interestingly, the reverse reaction, dehydrogenation of formic acid, was originally performed in tetrahydrofuran (THF).20 Methanol and THF are solvents with very different characteristics including different polarity, miscibility, dielectric constant and capacity to form hydrogen bonds. Encouraged by the different effect of these two solvents in the presence of [FeH(PP3)] and formic acid, we decided to carry on a theoretical study on the effect of different solvents in this reaction. Solvent effect at the Fe-complexes was introduced via the self-consistent reaction filed method and using Poisson-Boltzman solver (PBF).21 For the small molecules, we used the SM8 model by Truhlar and Cramer with the standard parameters for the different solvents,22 which has been shown to give very high accuracy for solvation free energies in different solvents, and has been parametrized to capture some of the explicit effects from hydrogen bonding. This model can be applicable to any charged or uncharged solute in any solvent or liquid medium for which a few key descriptors are

known, in particular the dielectric constant, refractive index, macroscopic surface tension, and acidity and basicity parameters.

Figure 3. Alternative free energy profile calculated starting + with the precursor complex [FeH(PP3)] where the hydride occupies the trans position to the bridgehead phosporous at the M06 level. The relative solvation corrected Gibbs free -1 energies are given in kcal mol . The selected bond distances and angles are given in angstroms and degrees, respectively.

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H 1.67 O C O

1.50

Fe

P Ph 2 H

1.50

1.58

P

TS(4_5) TS(6_5) PPh2 PPh2 m=1 m =1 PPh2 PPh2 2.80 O O P Fe 2 P Fe 2.76 .20 O O 2.1 9 P P H H Ph2 H Ph2 H 1.50

TS(4_6) PPh 2 m=1 PPh2

TS(4_6) + H 2O TS(4_ 5) + H 2O 25.2

27.7

TS(6_5) + H2O 27.5

19.8 12.4

9.8 int_6 + H 2O

PPh2 PPh2 Fe

PPh2 PPh2

1.54

H

P Ph2 H int_4 m =1

P

Fe

1.68

1.50

P

int_5 + H2O

H

P Ph 2 H int_6 O m=1

PPh2 PPh2 O

P

Fe

2.01

1.50

int_4 + CO2 + H2O

1.50

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O

O

P Ph2 H int_5 H m =1

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to higher dielectric aprotic solvents the solvation of bicarbonate is improved, and in both DMSO and acetonitrile the thermodynamic sinks are removed. It turns out that the solvation of bicarbonate is optimal also in these solvent to minimize the free energy barriers of the reaction. The effect of the solvation of the bicarbonate vs the overall activation free energy is illustrated in Chart 1. An additional possible effect from the solvent is formation of coordination complexes between the solvent and the catalyst. If this is too favorable it could lead to new thermodynamic sinks. For acetonitrile we find such favorable coordination, whereas DMSO does not form a strong bond to the iron center(Figure S4).

Figure 4. Free energy profiles calculated starting with the + precursor complex [FeH(PP3)] showing other possible path (in blue) of reaction of CO2 with the hydride complex followed by an isomerization step at the M06 level. The relative solvation corrected Gibbs free energies are given in kcal -1 mol . The selected bond distances and angles are given in angstroms and degrees, respectively.

Four protic solvents (hydrogen-donor solvents) were studied (see Figure 5a) with different dielectric constants (ε = 32.6, 24.8, 17.9 and 10.9). Comparing the rate determining steps of the calculated free energy profiles with the different solvents, we observed that in the presence of the solvent with the lowest dielectric constant (tBuOH) the barrier is 5 kcal mol-1 lower than the highest barrier for the highest dielectric constant solvent (MeOH). The barrier of the addition of the dihydrogen molecule to the iron complex does not change in energy with the application of different solvent (TS1_2 ≈ 23 kcal mol-1), since this reaction step does not involve any significant change in charge distribution and or solvation of small anions. However, the intermediate int_3 is 5 kcal mol-1 more stable in the lower dielectric constants solvents, and this difference is consistent with the difference observed in the barriers in the rate determining steps. Since this step involves transfer of the bicarbonate from full solvation to the inner-sphere of the iron complex the solvation free energy of the bicarbonate will have a very strong influence on this step. In methanol bicarbonate is simply too well solvated to react efficiently with the Fe dihydrogen complex. Since the reverse reaction was performed in THF originally we also studied the impact of three aprotic solvents (see Figure 5b) with different dielectric constants (ε = 47.2, 37.5 and 7.6). The free energy profile with the aprotic solvents shows that the largest barrier on the free-energy surface corresponds to the lowest dielectric constant solvent, THF (30.6 kcal mol-1) while the higher dielectric constant solvents (DMSO, CH3CN) have barriers of ≈ 23 kcal mol-1. As expected, bicarbonate is significantly less well solvated in THF compared to e.g. methanol. However, in THF the reactivity of bicarbonate is now so much higher that a new resting state is created, which also reflected in int_3, and more notably in the int_4 and int_5. These two latter intermediates are the resting states of the reaction in the presence of the THF. When changing

Figure 5. Free energy profiles calculated starting with the + precursor complex [FeH(PP3)] at the M06 level in different solvents: protic (a) and aprotic (b). The relative solvation -1 corrected Gibbs free energies are given in kcal mol .

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F_TS1_2) m=1 P

PPh2 PPh 2 3.15 Fe 3 H .43

P 0.74

P Ph2 F

H

P Ph2 F

PPh2 PPh2 F_TS(3_4) F_TS(4_5) m=1 m=1 1.56 0 H .89

Fe

H

1.60

O

P

O

P Ph2 F

H O

F_TS(1_2) + HCO3-

PPh2 PPh2 O Fe 2.76 1.50

According to the results presented we propose as ideal solvents for this reaction: a protic solvent, tBuOH, and aprotic solvents, DMSO (see Figure 6), with MeCN being very similar except that MeCN forms a stable metal solvent complex. Replacing the current methanol solvent with one of these two solvents (or one with similar properties) will make the reaction proceed under milder reaction conditions, meaning lower temperature and hydrogen pressure. Milder conditions should increase the lifetime of the catalyst, avoiding the decomposition of the sensitive ethyl-bridged ligand. The rate determining steps in the presence of either of these solvents is the dihydrogen addition to the active catalyst, [FeH(PP3)], with a free energy barrier of ≈ 23 kcal mol-1. There is a recent report by Zhu et al7 on the iron-catalyzed hydrogenation of sodium bicarbonate where the influence of several solvents was evaluated experimentally. Among the different solvents evaluated the highest activity was achieved using EtOH (in EtOH: TON of 447 and in MeOH: TON of 13) which agrees very well with our proposal of other solvent than MeOH for this reaction to improve the performance of the iron-complex. However, the reason for the effect of the solvent in the study by Zhu et al was not clarified.

hydridodihydrogen complex [FeH(η2-H2)(P(C6H4PPh2)3)]+ instead of the corresponding [FeF(η2-H2)(P(C6H4PPh2)3)]+ complex. From the calculated mechanistic studies of the simplified iron-complex F_int_1 with PP3 (see Figures 7), we can conclude that if the complex maintains the fluorine atom during all the catalytic cycle, the reaction should be very slow. The activation barriers for the dehydrogenation and the insertion of CO2 into Fe-H are very high (> 33 kcal/mol) while the activation barrier for the addition of H2 (22.2 kcal/mol) is easier than in the case of H_int_1 (23.6 kcal/mol). These results indicate that for the reaction to proceed the spectator ligand of the complex must be a hydride or transform into a hydride in the catalytic cycle. Indeed the computational studies agreed with the experimental mechanistic studies showing that in fact the complex [FeH(PP3)] will be formed from FeF(η2-H2)(PP3)]+ (see SI, Figures S1).

1.51

Chart 1. Plot of the activation free energy against the solvation free energy of the bicarbonate base.

33.7

m=1 15.9

2.1 9

O

H

F_TS(4_5) + H 2O

F_TS(3_4)

32.4

28.7

24.7

17.2

15.5

13.4

0.0 -4.0

PPh2 PPh2 H 1.54 Fe F P Fe 0.86 H 118 PPh2 P F PPh2 Ph2 F_int_1 F_int_2 m=3 m=1

F_int_4 + CO2 + H2 O

PPh 2 PPh2

PPh2

P

F_int_3

P

Fe

P Ph 2 F

PPh 2 PPh2

H

P H

O

O H

F_int_3 m=1

F_int_5 + H 2O

O

Fe

1.54

H

P Ph2 F

F_int_4 m=1

F_int_1 + H 2 O + HCO2-

PPh2 PPh2 P

Fe

2.01

O

1.50

F_int_2 + HCO3 -

1.50

F_int_1 + H2 + HCO3-

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P Ph2 F

F_int_5 m=1

PPh2 O P

H

Fe

1.54

F

118 PPh2

PPh2 F_int_1 m=3

Figure 7. Free energy profile calculated starting with the + precursor complex [FeF(PP3)] at the M06 level. The relative solvation corrected Gibbs free energies are given in kcal -1 mol . The selected bond distances and angles are given in angstroms and degrees, respectively.

Figure 6. Free energy profiles calculated starting with the + precursor complex [FeH(PP3)] at the M06 level in two dift ferent solvents ( BuOH and DMSO). The relative solvation -1 corrected Gibbs free energies are given in kcal mol .

In addition to the use of tetraphos P(CH2CH2PPh2)3 (PP3) as the ligand for the reduction of bicarbonate, Beller and coworkers reported a catalytic system with tris(2(diarylphosphino)aryl)phosphines in combination with Fe(BF4)2.9b The use of an in situ-generated catalyst or the defined active complex in this reaction led to significantly higher activity, good yields, and high turnover numbers. Mechanistic studies performed by Beller and coworkers, proposed that [FeF(P(C6H4PPh2)3)]+ is an intermediate on the pathway to the active species of the catalytic cycle. This species under hydrogen pressure forms the iron

We were curious about the origin of the improved efficiency by the arylphosphine iron complex and we therefore investigated the reaction mechanism with [FeH(P(C6H4PPh2)3)]+ as the active species (see Figure 8 for the formation of active complex [FeH(P(C6H4PPh2)3)]+. The catalytic cycle begins with the formation of the iron hydrido dihydrogen complex PAr_int_2 followed with the deprotonation by bicarbonate, which led to the dihydride iron complex PAr_int_4 (see Figure 8). This complex behaves similarly to the one with the ethyl-bridge ligand, where both the direct insertion of CO2 into the Fe-H bond, as well as the reaction of the hydride with the CO2 and posterior isomerization, to form O-bound formateiron complex PAr_int_5 are feasible paths. The final elimination of the formate molecule closes the catalytic cycle.

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2.1 2

+

Figure 8. Free energy profiles calculated starting with the precursor complex [FeH(P(C6H4PPh2)3)] at the M06 level. The relative -1 solvation corrected Gibbs free energies are given in kcal mol . The selected bond distances and angles are given in angstroms and degrees, respectively.

According to the calculated data presented here the use of the arylphosphine as the ligand does not have a significant influence in the reaction mechanism. It is important to note that in this case it is not possible to discriminate which is the rate determining step of the mechanism. The higher free energy barrier is ~27 kcal mol-1, which is only ~1 kcal mol-1 lower than the one obtained with [FeH(PP3)]+. This difference is not enough to explain the better performance of the aryl phosphine ligands. Instead it indicates that the improved efficiency of this catalyst is mainly a result of higher stability of the multidentate ligand, where the ethyl-bridge phosphine ligands are not sufficiently stable at elevated temperature (100 °C), where the hydrogenation of bicarbonate normally proceeds.

lower free-energy barrier than the same catalyst in the presence of the experimental solvent (MeOH), 23.7 kcal mol-1 in tBuOH compared to of 27.0 kcal mol-1 in MeOH. The energy of PAr_TS5_1 (elimination of formate) relative to the reference point of the iron formate coordinate ligand PAr_int_5 + H2O (see Figure 9). This agrees very well with our conclusion that better results could be obtained in the presence of either of these two ideal proposed solvents, tBuOH or DMSO, which would allow for the reaction at iron multidentate phosphine complexes to proceed under milder reaction conditions.

CONCLUSIONS To conclude, the mechanism of the iron-catalyzed hydrogenation of bicarbonate has been theoretically investigated using tetraphos as a ligand. We confirmed that the reaction occurs at one coordination site on the iron center, in agreement with our previous studies on the dehydrogenation of formic acid by the same iron catalyst.13

Figure 9. Free energy profiles calculated starting with the + precursor complex [FeH(P(C6H4PPh2)3)] at the M06 level t with BuOH as solvent for the reaction. The relative solvation -1 corrected Gibbs free energies are given in kcal mol .

Subsequent application of one of our computed ideal solvent (tBuOH) to the iron arylphosphine complex for the reduction of bicarbonate resulted in a significantly

In the experimental reports aprotic solvents were typically used for dehydrogenation of formic acid, while protic solvents were used for the reverse reaction, hydrogenation of bicarbonate or CO2. Our theoretical results show that the reaction is very dependent on the choice of the solvent used. In the presence of a hydrogen bond donor solvent (protic solvent), the bicarbonate substrate will be very well solvated and the further reaction with the [FeH(η2-H2)(P(C6H4PPh2)3)]+ complex results in a high free energy intermediate. According to our results the reactivity of the bicarbonate can be increased by simply using a protic solvent with slightly lower ability of solvating the bicarbonate anion. Such solvent should have a lower dielectric constant, which will improve reactivity of

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the substrate with the iron-hydride and thus the free energy barrier of the entire reaction is reduced. In the case of a non-hydrogen bonding solvent (non protic solvent), the situation is reverse. The bicarbonate is too poorly solvated in a medium with a low dielectric, such as THF, and the reaction with the metal complex results in a too stable intermediate, which creates new resting states in the catalytic cycle and overall a higher free energy barrier. For non-protic solvents it is therefore preferable to use a solvent with a high dielectric constant that reduces the possibility of forming stable and unreactive intermediates. Interestingly, the low dielectric protic solvents (tBuOH) and the high dielectric aprotic solvents (DMSO) have very similar free energy barriers. The possibility of the iron complex containing a fluoride ligand and maintaining it during all the catalytic cycle was also investigated. Based on our results the fluoride is most likely dissociated and replaced by a hydride [FeH(PP3)]+. The reaction mechanism was studied also for the tris(2-diphenylphosphine)phenyl)phosphine) ligand, which experimentally was described as the best experimental ligand for this reaction; the mechanistic investigation established that the active complex of the catalytic cycle is [FeH(PPhP3)]+. The higher performance of this catalyst [FeH(PP3)]+ versus [FeH(PPhP3)]+ cannot be explained by the current study of the reaction mechanism, in which there are not big differences between these two iron complexes. A plausible explanation of the higher efficiency is that the aryl bridges have higher stability than the ethyl bridges elevated temperatures.

Supporting Information. This material is available free of charge via the Internet at http://pubs.acs.org Table of energies of the optimized complexes. Cartesian coordinates of the computed structures.

AUTHOR INFORMATION Corresponding Author *E-mail: [email protected]

Present Addresses †L. X.: Department of Chemistry-Ångström Laboratory, Ångströmlaboratoriet, Lägerhyddsvägen 1, 75120 Uppsala, Sweden. ‡ R. S.-de A.: Materials Theory Division, Department of Physics and Astronomy, Uppsala University, P.O Box 516, S75120, Uppsala, Sweden.

Author Contributions The manuscript was written through contributions of all authors. / All authors have given approval to the final version of the manuscript.

Notes The authors declare no competing financial interest.

ACKNOWLEDGMENT Computational resources have been provided by the National Supercomputer Centre in Linköping, Sweden. Financial support has been received from Vetenskapsrådet.

REFERENCES COMPUTATIONAL DETAILS All the calculations were carried out with the Jaguar 7.6 program package.23 The Gibbs free energies were defined as GM06 = E(M06/LACV3P**++2f on Fe) + Gsolv + ZPE +H298-TS298 +1.9 (concentration correction to the free energy of solvation from M(g)- M(aq) to atm (g) – M(aq)). Molecular geometries were optimized at the B3PW9124 level of densityfunctional theory with the LACVP** basis set.25 Frequency calculations at the level were performed on the optimized geometries to verify that the geometries correspond to minima (no imaginary frequency) or transition states (one imaginary frequency) and to provide the thermochemical data at 298K, which include entropy contributions. Based on the gas-phase-optimized structures, the single-point solvation free energies were also calculated using the Poisson-Boltzmann reactive field implemented in Jaguar 7.6 (PBF)21 with the standard parameters for the different solvents using the calculations, except for the carbonate, formate and water which were calculated by using Solvation Model 8 (SM8)22 with the standard parameters of the solvents used. The single-point energy corrections were performed with M0626 functional using LACV3P**++ basis set27 augmented with two f-polarization functions on iron as suggested by Martin and Sundermann.28

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(1) (a) Jessop, P. G. In The Handbook of Homogeneous Hydrogenation, De Vries, J. G., Elsevier C. J., Eds.; Wiley-VCH: Weinheim, 2007; p 489–511; (b) Jessop, P. G.; Ikariya, T.; Noyori, R. Chem. Rev. 1995, 95, 259-272; (c) Sakakura, T.; Choi, J.-C.; Yasuda, H. Chem. Rev. 2007, 107, 2365–2387; (d) Enthaler, S. ChemSusChem 2008, 1, 801–804; (e) Wang, W.; Wang, S.; Ma, X.; Gong, J. Chem. Soc. Rev. 2011, 40, 3703–3727. (2) (a) Tsai, J. C.; Nicolas, K. M. J. Am. Chem. Soc. 1992, 114, 5117–5124; (b) Graf, E.; Leitner, W. J. Chem. Soc. Chem. Commun. 1992, 1465–1466; (c) Burgemeister, T.; Kastner, F.; Leitner, W. Angew. Chem. Int. Ed. Engl. 1993, 32, 739–741; Angew. Chem. 1993, 105, 781–783; (d) Fornika, R.; Görls, H.; Seemann, B.; Leitner, W. J. Chem. Soc. Chem. Commun. 1995, 1479–1491. (3) (a) Jessop, P. G.; Ikariya, T.; Noyori, R. Nature 1994, 368, 231–233; (b) Lau, C. P.; Chen, Y. Z. J. Mol. Catal. A 1995, 101, 33– 36; (c) Jessop, P. G.; Hsiao, Y.; Ikariya, T.; Noyori, R. J. Am. Chem. Soc. 1996, 118, 344–355; (d) Gao, J. C.; Kuncheria, J. K.; Jenkins, H. A.; Puddephatt, R. J.; Yap, G. P. A. J. Chem. Soc. Dalton Trans. 2000, 3212–3217; (f) Yin, C.; Xu, Z.; Yang, S. Y.; Ng, S. M.; Wong, K. Y.; Lin, Z.; Lau, C. P. Organometallics 2001, 20, 1216–1222; (g) Elek, J.; Nádashi, L.; Papp, G.; Laurenczy, G.; Joó, F. Appl. Catal. A 2003, 255, 59–67; (h) Federsel, C.; Jackste, R.; Boddien, A.; Laurenczy, G.; Beller, M. ChemSusChem 2010, 3, 1048–1050; (i) Balaraman, E.; Gunanthan, C.; Zhang, J.; Shimon, L. J. W.; Milstein, D. Nat. Chem. 2011, 3, 609–614; (j) Huff, C. A.; Sanford, M. S. J. Am. Chem. Soc. 2011, 133, 18122–18125; (k) Himeda, Y.; Miyazawa, S.; Hirose, T. ChemSusChem 2011, 4, 487–493; (l) Baffert, M.; Maishal, T. K.; Mathery, L.; Copéret, C.; Thieuleux, C. ChemSusChem 2011, 4, 1762–1765; (m) Wesselbaum, S.; Vom Stein, T.; Klankermayer, J.; Leitner, W. Angew. Chem. Int. Ed. 2012, 51, 7499–7502; Angew. Chem. 2012, 124, 7617–7620; (n ) Han,

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Z.; Rong, L.; Wu, J.; Zhang, L.; Wang, Z.; Ding, K. Angew. Chem. Int. Ed. 2012, 51, 13041–13045; Angew. Chem. 2012, 124, 13218–13222; (o) Gowrisankar, S.; Federsel, C.; Neumann, H.; Ziebart, D.; Jackstell, R.; Spannenberg, A.; Beller, M. ChemSusChem 2013, 6, 85–91. (4) (a) Himeda, Y.; Onozawa-Komatsuzaki, N.; Sugihara, H.; Arakawa, H.; Kasuga, K. Organometallics 2004, 23, 1480–1483; (b) Himeda, Y.; Onozawa-Komatsuzaki, N.; Sugihara, H.; Arakawa, H.; Kasuga, K. Organometallics 2007, 26, 702–712; (c) Tanaka, R.; Yamashita, M.; Nozaki, K. J. Am. Chem. Soc. 2009, 131, 14168– 14169; (d) Schmeier, T. J.; Dobereiner, G. E.; Crabtree, R. H.; Hazari, N. J. Am. Chem. Soc. 2011, 133, 9274–9277. (5) (a) Wiener, H.; Blum, J.; Feilchenfeld, H.; Sasson, Y.; Zalmanov, N. J. Catal. 1988, 110, 184–190; (b) Stalder, C. J.; Chao, S.; Summers, D. P.; Wrighton, M. S. J. Am. Chem. Soc. 1983, 105, 6318–6320. (6) For recent reviews see: (a) Bolm, C.; Legros, J.; Paith, J. L.; Zani, L. Chem. Rev. 2004, 104, 6217–6254; (b) Enthaler, S.; Junge, K.; Beller, M. Angew. Chem. Int. Ed. 2008, 47, 3317–3317; Angew. Chem. 2008, 120, 3363–3367; (c) Gaillard, S.; Renaud, J.-L. ChemSusChem 2008, 1, 505–509; (d) Morris, R. H. Chem. Soc. Rev. 2009, 38, 2282–2291; (e) Junge, K.; Schröder, K.; Beller, M. Chem. Commun. 2011, 47, 4849–4859; (f) Carwish, M.; Wills, M. Catal. Sci. Technol. 2012, 2, 243–255. (7) Zhu, F.; Zhu-Ge, L.; Yang, G.; Zhou, S. ChemSusChem 2015, 8, 609–612. (8) Tai, C.-C.; Chang, T.; Roller, B.; Jessop, P. G. Inorg. Chem. 2003, 42, 7340–7341. (9) (a) Federsel, C.; Boddien, A.; Jackstell, R.; Jennerjahn, R.; Dyson, P. J.; Scopelliti, R.; Laurenczy, G.; Beller, M. Angew. Chem. Int. Ed. 2010, 49, 9777–9780; (b) Ziebart, C.; Federsel, C.; Anbarasan, P.; Jackstell, R.; Baumann, W.; Spannenberg, A.; Beller, M. J. Am. Chem. Soc. 2012, 134, 20701–20704. (10) Langer, R.; Diskin-Posner, Y.; Leitus, G.; Shimon, L. J. W.; Ben-David, Y.; Milstein, D. Angew. Chem. Int. Ed. 2011, 50, 9948– 9952; Angew. Chem. 2011, 123, 10122–10126. (11) Bertini, F.; Mellone, I.; Ienco, A.; Peruzzini, M.; Gonsalvi, L. ACS Catal. 2015, 5, 1254–1265. (12) Yang, X. Dalton Trans. 2013, 42, 11987–11997. (13) Sánchez-de-Armas, R.; Xue, L.; Ahlquist, M. S. G. Chem. Eur. J. 2013, 19, 11869–11873. (14) (a) Osadchuk, I.; Tamm, T.; Ahlquist, M. S. G. Organometallics 2015, 34, 4932–4940; (b) Xue, L.; Ahlquist, M. S. G.; Inorg.

Page 8 of 9

Chem. 2014, 53, 3281–3289; (c) Ahlquist, M. S. G. J. Mol. Cat. A 2010, 324, 3–8. (15) Polukeev, A. V.; Marcos, R.; Ahlquist, M. S. G.; Wendt, O. F. Chem. Eur. J. 2016, 22, 4078-4086. (16) Tsay, C.; Livesay, B. N.; Ruelas, S.; Yang, J. Y. J. Am. Chem. Soc., 2015, 137, 14114–14121. (17) Bianchini, C.; Laschi, F.; Peruzzini, M.; Ottaviani, F. M.; Vacca, A.; Zanello, P. Inorg. Chem. 1990, 29, 3394–3402. (18) Bianchini, C.; Peruzzini, M.; Zanobini, F. J. Organomet. Chem. 1988, 354, C19–C22. (19) Bianchini, C.; Masi, D.; Peruzzini, M. Inorg. Chem. 1997, 36, 1061–1069. (20) Boddien, A.; Mellmann, D.; Gärtner, F.; Jackstell, R.; Junge, H.; Dyson, P. J.; Laurenczy, G.; Ludwig, R.; Beller, M. Science 2011, 333, 1733–1736. (21) (a) Tannor, D. J.; Marten, B.; Murphy, R.; Friesner, R. A.; Sitkoff, D.; Nicholls, A.; Ringnalda, M.; Goddard, W. A.; Honig, B. J. Am. Chem. Soc. 1994, 116, 11875–11882; (b) Marten, B.; Kim, K.; Cortis, C.; Friesner, R. A.; Murphy, R. B.; Ringnalda, M. N.; Sitkoff, D.; Honig, B. J. Phys. Chem. 1996, 100, 11775–11788. (22) (a) Marenich, A. V.; Olson, R. M.; Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Chem. Theory Comput. 2007, 3, 2011–2033; (b) Marenich, A. V.; Cramer, C. J.; Truhlar. D. G. J. Phys. Chem. B 2009, 113, 4538–4543, (23) Jaguar, version 7.6, Schrödinger, LLC, New York, NY, 2009; www.schrodinger.com (24) (a) Becke, A. D. Phys. Rev. A. 1998, 38, 3098–3100; (b) Perdew, J. P. In Electronic Structure Theory of Solids; Ziesche, P., Eschrig, H., Eds.; Akademic Verlag: Berlin, 1991; p 11–20; (c) Perdew, J. P.; Chevary, J. A.; Vosko, S. H.; Jackson, K. A.; Pedersen, M. R.; Singh, D. J.; Fiolhais, C. Phys. Rev. B. 1992, 46, 6671– 6687. (25) LACVP basis sets use 6-31G for main group elements and the Hay-Wadt ECP for Fe: Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 299–310. (26) Zhao, Y.; Truhlar, D. G. Theor. Chem. Acc. 2008, 120, 215– 241. (27) The LACV3P basis set is a triple-zeta contraction of the LACVP basis set developed and tested at Schrödinger, Inc. LACV3P basis sets use 6-311G for main-group elements. (28) Martin, J. M. L.; Sundermann, A. J. Chem. Phys. 2001, 114, 3408–3420.

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