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Why is there a barrier in the coupling of two radicals in the water oxidation reaction? Ting Fan, Shaoqi Zhan, and Mårten S.G. Ahlquist ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.6b02697 • Publication Date (Web): 03 Nov 2016 Downloaded from http://pubs.acs.org on November 3, 2016
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Why is there a barrier in the coupling of two radicals in the water oxidation reaction? Ting Fan, Shaoqi Zhan, Mårten S. G. Ahlquist* Division of Theoretical Chemistry & Biology, School of Biotechnology, KTH Royal Institute of Technology, 10691 Stockholm, Sweden. ABSTRACT: Two radicals can form a bond without an energetic barrier. Yet, the radical coupling mechanism in ruthenium catalyzed water oxidation has been found to be associated with substantial activation energies. Here we have investigated the coupling reaction of [Ru=O(bda)L2]+ catalysts with different axial L ligands. The interaction between the two oxo radical moieties at the Ru(V) state was found to have a favorable interaction in the transition state compared to the prereactive complex. To further understand the existence of the activation energy, the activation energy has been decomposed into distortion energy and interaction energy. No correlation between the experimental rates and the calculated coupling barriers of different axial L was found, showing that more aspects such as solvation, supramolecular properties, and solvent dynamics likely play important roles in the equilibrium between the free RuV=O monomer and the [RuV=O•••O=RuV] dimer. Based on our findings, we give general guidelines for the design catalysts that operate by the radical coupling mechanism.
KEYWORDS: water oxidation, radical coupling, I2M, DFT, bda, catalysis, ruthenium
Introduction Why is there a barrier in the coupling of two radicals in the water oxidation reaction? Two radical typically form a bond in a potential well with no barrier. Take the simplest example where two hydrogen atoms form a bond. As the distance decreases between the nuclei the two unpaired electrons can occupy the same space and the kinetic energy decreases. 1 In the singlet state there is no Pauli repulsion between the two electrons with opposite spin, Therefore, other repulsive interactions are typically overcome by the attractive forces when the bond is formed between two radicals. In water oxidation catalyzed by transition metal complexes, a unifying feature of the mechanism is that high valent metal oxo species form, and that formation precedes the formation of the O-O bond. Typically the metal oxo has one or more unpaired electrons, with some exceptions including some iridium complexes.2 For the vast majority of synthetic catalysts, the O-O bond is formed via nucleophilic attack by a water molecule on the high valent oxo species.3 However, most of the active metal oxo species have unpaired electrons, where a significant fraction of the spin is located on the oxygen atom.4 Why do these catalysts not react via radical coupling? Figure 1 shows the water nucleophic attack (WNA) and radical coupling mechanisms (I2M = interaction of two metal centers) of the water oxidation reaction.5 If radicals couple without enthaplic barriers, the I2M reaction mechanism should have an intrinsic advantage. Of course diffusion and concentration effects are always counteracting the barrierless radical reactions, but these could be overcome by making the catalysts less hydrophilic and by increasing the catalyst concentration. There are a limited number of examples that actually react via radical coupling,6 where the most prominent example is the Ru(bda)L2 catalyst (bda = 2,2’-bipyridine-6,6’-dicarboxylate, L = typically nitrogen
heterocycle such as pyridine) by Sun and co-workers.7 According to the studies by Siegbahn on proposed structures at the S4 state of the oxygen evolving complex of photosystem II, the O-O bond forming step takes place at a species where the electrons on the two reacting metals are antiferromagnetically coupled.8 At the tested geometries the internal O-O bond formation was clearly favored over other mechanisms. These results indicate that also nature utilizes a form of radical interaction, although an intramolecular version, and thereby avoids unfavorable Pauli repulsion of a nucleophilic addition reaction. The activation energy present in the intramolecular version of radical couplings is due to the strain required for the two oxygen atoms to approach each other. Similar results were recently seen for a pentanuclear iron complex, which was found to react via internal O-O coupling. 9 Åkermark and Siegbahn recently presented experimental and theoretical investigations of the mechanism of water oxidation by a dinuclear ruthenium catalyst, which was found to proceed via internal O-O coupling.10 Despite low O-O bond formation activation energies, these catalysts were found to be limited by the release of O2. The single-site ruthenium catalysts by Sun show extreme reactivity even at pH 1, and their rates are approximately 1 000-1 000 000 times higher than similar catalysts reacting via water nucleophilic attack.11,12 The reaction of the Ru(bda)L2 complexes have been shown to follow second order kinetics,7 and was shown by Privalov and co-workers that the mechanism likely proceeds via a coupling of two RuV=O radicals.13 The most likely reaction was found to be on the open shell singlet surface. In Privalov’s report they found that most of the spin density was located on the oxo, and that the activation energy from the prereactive complex [(bda)L2RuV=O•••O=RuVL2(bda)]2+ was on the order of 5-11 kcal/mol.7b This clearly shows that this reaction can occur with relatively low activation energies. However, one thing still
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puzzled us. If the reaction is indeed between two oxyl radicals, why does it at all have a barrier for the formation of the O-O bond? If these oxo groups are truly radicals they should couple in a barrierless process. To understand this in more detail we used density functional theory with empirical dispersion correction to study a series of [RuV=O(bda)L2]+ complexes. The activation energy has also been decomposed into different contributions, and based on this we show the origin of the activation energy. From the information of the different complexes, we find the key to the success of these complexes and we give a guideline for the design of new catalysts that utilize the repulsion free, barrierless radical coupling. We believe that this information is transferable to other metals than Ru, and we point out a direction that could assist in solving the challenge to design first row metals with reactivity similar to Ru(bda).
Figure 1. Illustration of WNA and radical coupling mechanisms for O-O bond formation. Computational Details All DFT calculations for free energy were carried out with the Jaguar 8.3 program package by Schrödinger LLC.14 Molecular geometries were optimized at the Becke’s threeparameter hybrid functional and the LYP correlation functional (B3LYP)15 with D3 correction of Grimme et al.16 with the LACVP** basis set17. We searched the potential energy surface by scanning the terminal O-O bond distance [RuV=O•••O=RuV] of the antiferromagnetic open shell singlet. On the basis of the gas phase optimized geometries, the solvation energies were estimated by the single-point calculations using the Poisson-Boltzmann reactive field implemented in Jaguar 8.3 (PBF) in water. We also investigated the reaction at the triplet state but as found by Privalov13 it is in every case associated with higher activation energies. Only electronic activation barrier ∆Eǂ and solvated electronic activation barrier ∆Eǂsolv are discussed in the main paper. Results and Discussions Under acidic conditions with Ce(IV) as the oxidant the mechanism of the majority of the Ru(bda)(L2) complexes follow second order kinetics with respect to the catalyst.7 A dimer of a seven coordinate RuIV complex was found, and this complex was found to be catalytically active.18 These results in combination with electrochemical data on the onset potential lead to the mechanistic proposal that two RuV=O species react to form the O-O bond leading to the RuIV-OO-RuIV intermediate.7,13,19 Based on the way of the two monomers interacted it was reasoned that larger and more hydrophobic axial ligands could favor the reaction. Indeed, isoquinoline,7b phtalazine7c and more recently halogenated versions of these ligands7d were found to have improved activity and turnover frequencies (TOF) exceeding 1000 s-1. In a study of a series of substituted pyridines as axial ligands, all catalysts showed second order kinetics, with the exception of the 3-pyridine-sulfonate. Since the sulfonylated pyridine is more water soluble the formation of the dimer is hindered, which again confirms the proposed mechanism. Also complexes where the pyridyl ligand was replaces by N-heterocyclic carbene ligands followed first
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order kinetics, but that was found to be due to ligand replacement by water which makes the complex too hydrophilic. 20 Among the other pyridines there appeared to be a correlation with the electron withdrawing/donating properties of the substituents on the pyridine, where electron withdrawing favored the reactivity.7a In the current study we have compared the activation electronic energies of eight complexes with different axial ligand with different experimental activity (Figure 2). A recent theoretical study probed the ligand effects on O-O bond formation of Ru-catalyzed water oxidation and it is found that bulky groups on the equatorial ligand can increase the barrier.21 In this report, we are only looking at the bond forming step, assuming that the dimer can form from the free solvated monomers. We are aware that the formation of the pre-reactive dimer is very important but that part is the topic of another study. In presence of an excess of Ce(IV) at pH 1 these catalysts show different activity as shown in Figure 2. One initial indicator of the reactivity could be the spin density at the oxo of the doublet RuV=O species. However, we find that all the studied complexes have more or less identical values, and the values only differ on the second decimal. This indicates that they should all be able to react via radical coupling given that the pre-reactive dimer forms. We then calculated the barriers of all the species. The transition states were located by scanning the O-O bond distance of the antiferromagnetic open shell singlet. All complexes have very similar reaction paths, although the interaction between the ligands differ somewhat. The isoquinoline based ligands seem to have a higher tendency to form π-stacking interactions in agreement with previous proposals.7b,22 Interestingly, we find no correlation between the observed rate and calculated barrier of the different axial ligands, meaning that the π-stacking does not seem to aid the OO bond forming step itself, but rather facilitates the formation of the pre-reactive dimer. The complexes have barriers in a range of 1.4 to 7.2 kcal/mol relative to the respective prereactive complexes (table 1). Not only are these barriers very low, but the lack of correlation between the activation energy and the experimental rates indicate that another step is causing the different reactivities, possibly the formation of the dimer from the free solvated monomer.
reaction rate 8 > 7 > 1 ≈ 2 > 3 ≈ 4 ≈ 5 > 6 Figure 2. Molecular structures of complexes 1-8 Table 1. Activation electronic energy and solvated electronic energy barriers of coupling step.
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complexes
Electronic activation energy ∆E‡ (kcal/mol)
Solvated electronic activation energy ∆E‡solv(kcal/mol)
TOF (s )
1
2.2
3.1
119a
2
1.4
3.9
115a
3
2.4
5.4
25a
4
3.9
6.4
32b
5
3.5
5.0
25a
6.7
a
6
6.7
7
6.8
8 a
7.2 -4
IV
-1
14
b
5.8
300
5.4
1000c
catalyst (2.16 ×10 M) and Ce (0.327 M), . Duan, L.; Wang, L.; Inge, A. K.; Fischer, A.; Zou, X.; Sun, L. Inorg. Chem. 2013, 52, 7844-7852. b catalyst (2.16 ×10-4 M) and CeIV (0.48 M), Duan, L.; Bozoglian, F.; Mandal, S.; Stewart, B.; Privalov, T.; Llobet, A.; Sun, L. Nature. Chem. 2012, 4, 418-423. c catalyst (6.25 ×10-5 M) and CeIV (0.365 M), Wang, L.; Duan, L.; Wang, Y.; Ahlquist, M. S. G.; Sun, L. Chem. Commun. 2014, 50, 12947-12950. Still, as discussed in the introduction, why do any of these complexes show any activation energy at all? The two radicals should have a mostly attractive interaction. If we just look at the two oxyl radicals separately, and assume no Pauli repulsion, we can calculate the Coulomb repulsion from the partial charges and the dispersion attraction from assuming standard parameters for oxygen and thereby get a first measure on the interaction between the two radicals. Of course we discard many effects in this way, the interaction of the rest of the molecule and the stabilization from the interaction between the two unpaired electrons, yet it gives a measure of the direction of the interaction. For the dispersion we used the attractive Lennard-Jones (LJ) potential (eq 1) with parameters from the OPSL-AA force field of σ = 2.96 and ε = 0.21 for oxygen23. For the Coulomb interaction (eq 2) we used the ESP charge at the monomer of -0.113. If we combine the LJ attraction and the Coulomb repulsion we get an interaction that is altogether slightly repulsive at 2.43Å (+1.3 kcal/mol), due to neglect of all other interactions that keep the complexes together. The interaction then decreases and becomes slightly negative at the TS O-O distance of 1.9Å (-0.8 kcal/mol). This is an indication that the two oxygen atoms would prefer the transition state geometry over the pre-reactive complex by 2.1 kcal/mol. If the electron delocalization energy is included, the interaction could be even stronger. To further test the magnitude of the interaction, we exchanged the oxo groups, or oxyl radicals, for fluorines (Scheme 1). This gives singlet complexes, similar to the RuIV-OH complexes, where there are no unpaired electrons. The interaction when the distance between the two fluorine atoms decreases by the same amount as the oxos, from 2.43 to 1.9 Å, should therefore be significantly more repulsive due to Pauli repulsion. Indeed, the energy calculated with DFT increases by 16.0 kcal/mol as the F---F distance decreases. This value is 11.9 kcal/mol higher than for the original complex that contain the two oxo ligands. Again calculating the Lennard-Jones potential between the two atoms (with σ = 2.94 and ε = 0.061) combined with a Coulomb potential based on an ESP charge of -0.18 at fluorine, we get an estimate of the interaction of the two atoms. In this case we must include the repulsive LJ term since there are no unpaired
electrons. The interaction energy is then calculated to be 4.8 kcal/mol at 2.43 Å and to 16.5 kcal/mol at 1.9 Å, indicating an increased repulsion of 11.7 kcal/mol as the distance decreases, which gives further indication that the O---O interaction is not repulsive itself. This interaction is 13.8 kcal/mol (11.7 – (-2.1)) more repulsive than the interaction at the oxo-complex, in very good agreement with the full DFT difference 11.9 kcal/mol. Combined these results indicate a slightly favorable, nonrepulsive interaction between the two oxo radicals.
= 4[
=
− ]
(1)
(2)
Scheme 1 From the results above, we see that the interaction between the oxo groups is not disfavoring the transition state geometry, but since there is a barrier, something else must be disfavoring it. Since the radical is part of a large unit, it is likely that the sum of the weak interactions, mainly electrostatic and dispersion interactions, will favor a geometry, which is not optimal for O-O bond formation. To understand the forces leading to the barrier we have decomposed the activation energy in two terms, distortion energy (Edist) and interaction energy (Eint) (Figure 3). Edist is the energy required to distort the two reactants from the starting pre-encounter complex geometry to the transition state geometry. Eint is the interaction energy difference of the distorted fragments between the pre-reactive complex and transition state. As is seen in table 2 the Edist accounts for 1.7 to 6.7 kcal/mol of the strain. The Eint unfortunately also includes the favorable O-O interaction, but since this is of similar magnitude in all cases we still get a measure of the intermolecular interaction. The magnitude of Eint ranges from -3.1 to 4.2 kcal/mol. There appears to be a slight reverse relationship between the magnitude of the two terms. If Eint is high, then Edist is low, and vice versa. In that way the complex can either adopt a geometry that has very small effect on the intermolecular interactions, however it pays the price in the Edist, or it could make minimal adjustments to the monomer but instead disrupt the favorable intermolecular interaction Eint. But summed up, the activation energies all fall within a quite small range. From these results we conclude that the origin of the gas phase activation energy is the combination of the distortion and interaction of the pre-encounter complex.
∆Eaǂ TSA
A
TSA΄
A΄
Sin A
A΄
TSA
TSA΄
∆Edist= E(TSA) + E(TSA΄) – E(A)-E(A΄) ∆Eint = ∆Eaǂ - ∆Edist Figure 3. Illustration of distortion and interaction energy.
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face-to-face 0.0 kcal/mol
face-to-end 3.1 kcal/mol
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end-to-end 28.8 kcal/mol
Figure 4. Face-to-face, face-to-end and end-to-end conformations of pre-reactive complexes 3 with relative gas phase free energies. Table 2. Distortion/interaction energy analysis for TSs. complexes
∆Edist (kcal/mol)
∆Eint (kcal/mol)
TOF (s-1)
1
1.7
0.5
119
2
4.5
-3.1
115
3
5.0
-2.0
25
4
3.0
0.9
32
5
1.7
1.8
25
6
2.5
4.2
14
7
3.1
3.7
300
8
6.7 0.5 1000 The effect of solvation on the reaction could in a simple form be estimated with a continuum model. Here we used the PBF model in Jaguar. 24This model cannot account for any explicit solvation effects or give insight to the effects of the dynamical behavior of the solvent, which is a limitation. However, it is still interesting to see the effect of merely applying a solvation field. The explicit solvent effect will be the topic of another study. The solvated activation energy does not change much compared to the gas phase barrier. The effect of implicit solvation was relatively small in all cases and seems to even out the relative values (table 1). In vacuum the difference in activation energy ranged from 1.4 to 7.2 kcal/mol, while with implicit solvation this range is reduced to 3.9 to 6.7 kcal/mol. Interestingly, we find no clear correlation between the observed rate and the calculated activation energy for the O-O bond forming step, neither with or without implicit solvation. Ever since the crystallization of the RuIV dimer18 we have assumed that the complex prefers a structure with the oxo groups facing each other in a face-to-face complex. However, if a face-to-end or end-to-end complex is more stable this would cause additional distortion energy to reach the transi-
tion state. We therefore optimized face-to-end and end-to-end geometries (Figure 4). We find that these configurations are less favored compared to the face-to-face geometry. Clearly this feature is important for any complex that react via radical coupling, that no low energy conformations exist that could increase the Eint required to reach the TS. It could also have implications to design of covalently linked complexes where the link could disrupt the favorable formation of the prereactive dimer type structure. Conclusions To conclude we have found that the intermolecular I2M, or radical coupling reaction, does not have an intrinsic barrier between the oxo fragments of RuV=O complexes. The activation energy is very low, and is composed of distortion and interaction energy that is required to break the favorable prereactive dimer geometry. The effect of the axial ligand appears to be minimal on the barrier of this particular step. From that we conclude that the effect of the axial ligand is likely mostly on the equilibrium between the free RuV=O monomer and the [RuV=O•••O=RuV] dimer. The radical coupling between two M=O fragments is typically intermolecular, and effects such as solvation, dimerization, diffusion and supramolecular interactions play major roles. Therefore, the synthesis of a complex with the appropriate properties could be more challenging than complexes reacting via water nucleophilic attack. However, the lack of intrinsic barrier between the oxo groups means that the reaction rate could be arbitrarily high, up to the limit of diffusion control. Based on the findings here we propose a few guidelines for synthesizing catalysts that operate by the radical coupling mechanism. 1) The oxo must have significant spin density at the oxo for the reaction to be possible. Closed shell singlet species are unlikely to function well. Species with very low spin density at the oxo could require excessive distortion of the electron density that could cause a significant activation energy.
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3) 4) 5)
The complexes studied herein have charge of 1 per unit. More highly charged species could be less likely to form prereactive dimers. The complexes must prefer a face-to-face geometry. Any other preference will lead to higher Eint. The steric bulk must be limited around the reactive site to allow for two oxos to couple. The hydrophobic and hydrophilic properties need to be balanced.
AUTHOR INFORMATION Corresponding Author Email for M. S. G. Ahlquist:
[email protected] Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript.
Funding Sources This research has been funded by Vetenskapsrådet, Wenner-Gren Foundation and China Scholarship Council (CSC). Computational resources have been provided by the National Supercomputer Centre in Linkӧping, Sweden.
ASSOCIATED CONTENT Supporting Information. Cartesian coordinates of Optimized structures. This material is available free of charge via the Internet at http://pubs.acs.org.
ACKNOWLEDGMENT MSGA Prof. Yi Luo and Prof Licheng Sun are acknowledged for insightful discussions.
ABBREVIATIONS WNA water nucleophilic attack, LJ Lennard-Jones, bda bipyridine dicarboxylate
REFERENCES
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( 20 ) Staehle, R.; Tong, L.; Wang, L.; Duan, L.; Fischer, A.; Ahlquist, M. S. G.; Sun, L.; Rau, S. Inorg. Chem. 2014, 53, 13071319. (21) Kang, R.; Chen, K.; Yao, J.; Shaik, S.; Chen, H. Inorg. Chem. 2014. 53, 7130-7136. (22) Richmond, C. J.; Matheu, R.; Poater, A.; Falivene, L.; BenetBuchholz, J.; Sala, X.; Cavallo, L.; Llobet, A. Chem. Eur. J. 2014, 20, 17282-17286. (23) The parameters supplied with TINKER are from "OPLS AllAtom Parameters for Organic Molecules, Ions, Peptides & Nucleic Acids, July 2008" as provided by W. L. Jorgensen, Yale University during June 2009. These parameters are taken from those distributed with BOSS Version 4.8. (24) Marten, B.; Kim, K.; Cortis, C.; Friesner, R. A.; Murphy, R. B.; Ringalda, M. N.; Sitkoff, D.; Honig, B. J. Phys. Chem. 1996, 100, 11775-11788.
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