Thermodynamic Hydricity across Solvents: Subtle Electronic Effects

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Thermodynamic Hydricity across Solvents: Subtle Electronic Effects and Striking Ligation Effects in Iridium Hydrides Kelsey R. Brereton, Caleb N. Jadrich, Bethany M. Stratakes, and Alexander J. M. Miller* Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290, United States

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S Supporting Information *

ABSTRACT: Insights into the influence of solvent on hydride transfer emerge from a study of iridium complexes with the formula [Cp*Ir(bpy-X)H]+ (bpy-X = 4,4′-X2-2,2′bipyridine, X = H, Me, tBu, OMe, CO2Me, and CF3). Hydricity (or hydride donor ability) is found to be equally sensitive to bipyridine ligand electronic structure in both CH3CN and H2O. In contrast, hydride transfer is found to be more strongly influenced by subsequent chloride ion binding to the metal center in CH3CN than in H2O. With thermochemical parameters for six iridium complexes available in both CH3CN and H2O, a general approach to comparing thermodynamic parameters across solvents was developed. The free energy to transfer the free hydride ion from water to eight organic solvents (acetonitrile, methanol, ethanol, ethylene glycol, dimethyl sulfoxide, N-methylpyrrolidin-2-one, ethylene carbonate, and tetrahydrofuran) was estimated. An equation based on the solvent transfer free energies of the hydride ion and the organometallic species involved in hydride transfer was developed, enabling accurate and quantitative predictions of the change in hydricity moving between solvents.

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INTRODUCTION Transition metal hydrides mediate catalytic transformations relevant to biology, commodity chemical and pharmaceutical synthesis, and energy conversion.1 These diverse transformations can be carried out in solvents ranging from aprotic organics to polar protic alcohols and water. Thermodynamic hydricity values (ΔG°H−) can help chemists understand and predict hydride transfer reactivity, but the vast majority of studies have been limited to a single solvent, acetonitrile (CH3CN).2 Starting with Creutz’s two Ru complexes,3,4 hydricities of a scant few complexes have been determined in two solvents.5−12 Almost all of these complexes exhibit a striking increase in hydride donor ability (decrease in ΔG°H− value) moving from CH3CN to H2O. However, ΔG°H− does not always change by the same amount, so the relative hydricity of hydrides is different in water and acetontrile. Furthermore, the hydricity values of small molecules are impacted by solvent differently than the metal hydride complexes. The formate ion (HCO2−), for example, has relatively similar hydricity in CH3CN and H2O.2 The different relative hydricities of metal complexes and small molecules in different solvents have been exploited for significant advances in the field of CO2 reduction, where a hydride transfer that is uphill in acetonitrile can become thermodynamically favorable in water.4,6,7,13,14 A quantitative understanding of how thermodynamic hydricity will change across solvents and the factors that determine the relative hydricity of hydride donors in each solvent has been lacking. Deeper conceptual and quantitative © XXXX American Chemical Society

insights into the influence of solvent on hydricity would provide more predictive power and new opportunities for applications in catalyst design. To better understand the impact of solvent on hydricity, we have sought to expand the number of metal hydrides with welldefined hydricity in multiple solvents. The first step in this direction was the development of a general method for the measurement of aqueous hydricity values.8 These prior studies by our group established trends in aqueous hydricity for [Cp*Ir(bpy-X)H]+ complexes that catalyze reactions, including CO2/HCO2− interconversion,2,15,16 in both organic and aqueous solvents. We now extend these studies to CH3CN, enabling unprecedented comparisons of hydricity across solvents. The influence of electronic tuning on a series of metal hydrides is examined in water and acetonitrile. In addition, the effect of chloride ion binding on hydricity is quantified in CH3CN, allowing comparisons of “effective hydricity”2 as a function of solvent. The new data were used to build a framework for understanding the influence of solvent on hydricity and the factors that influence the relative hydricity in each solvent. To facilitate comparisons across solvents, a unified set of solventspecific constants for use in each relative hydricity scale is provided. An equation based on the free energy of transfer between solvents is developed that enables accurate Received: April 28, 2019

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DOI: 10.1021/acs.organomet.9b00278 Organometallics XXXX, XXX, XXX−XXX

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the present studies in CH3CN (several days) than for the prior study in water (minutes).8 Creutz and co-workers also observed hydride transfer reactions to be slower in CH3CN than in H2O,4,18 which may be due to relatively slow dissociation of acetonitrile preceding hydride transfer. A Hammett plot in CH3CN (Figure 2A, red), where KX/KH is the ratio of hydricity equilibrium constants and Σσp− is the

predictions of changes in transition metal and small molecule hydricity between water and acetonitrile. This approach, which can in principle be extended to any of eight organic solvents based on the constants provided here, provides a new quantitative understanding of how the hydricity of individual species are differently influenced by changes in solvent.

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RESULTS AND DISCUSSION Solvent-Dependent Electronic Effects on Hydricity. A set of six chloride complexes [Cp*Ir(bpy-X)(Cl)][Cl] was synthesized by addition of various substituted bipyridine ligands to [Cp*Ir(Cl)2]2. Chloride abstraction in H2O was followed by dissolution in CH3CN to provide [Cp*Ir(bpyX)(NCCH3)][PF6]2. X-ray diffraction studies on [Cp*Ir(bpyMe)(NCCH3)][PF6]2 and [Cp*Ir(bpy-CF3)(NCCH3)][PF6]2 confirmed the expected “piano stool” structure (Figure S3 and S12).17 Hydride complexes were prepared from the chloride complexes following previously developed methods (Supporting Information section I). Hydricity values for the six Ir hydride complexes were determined in CH3CN by establishing hydride transfer equilibria (Figure 1). Reactions were monitored by 1H NMR

Figure 2. Hammett correlations for the hydricity of [Cp*Ir(bpyX)H]+ showing experimental values (A) and DFT values (B) in CH3CN (red, this work) and H2O (blue, ref 8).

sum of Hammett parameters,19−21 exhibits a linear correlation with a negative slope (ρ) consistent with the buildup of positive charge after hydride release. The parameter σp− exhibited the best correlation (Figure S38), consistent with significant resonance contributions between the bpy ligand and Ir d-orbitals.21 Combining the new data on electronic effects in CH3CN with data from our previous study in H2O8 provides the first direct comparison of electronic effects on thermodynamic hydricity in two different solvents. Comparison with the aqueous data (Figure 2A, blue) reveals identical ρ values, indicating no change in sensitivity to substituents between CH3CN and H2O. DFT computations provide further support for the lack of solvent dependence, with Figure 2B showing only slightly increased sensitivity in CH3CN relative to H2O. A different trend is observed for the acidity of organic acids (e.g., phenols and benzoic acids). These species, with localized

Figure 1. Equilibrium measurements and hydricities in CH3CN. Vertical arrows show established equilibria. Free energy relative to [Cp*Ir(bpy)H]+ is given next to arrows, and the effective hydricity (ΔG°H−(CH3CN)) is given in italics.

spectroscopy and concentrations at the end point provided an equilibrium constant (Keq) and the relative hydricity. The effective hydricities in CH3CN, ΔG°H−(CH3CN), were obtained by reference to the known hydricity of [Cp*Ir(bpy)H]+ (Figure 1).10 A similar method was used in the prior study of aqueous hydricity; approach to equilibrium was slower for B

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Organometallics charges and strong hydrogen bonding interactions with the solvent, exhibit more than 2-fold changes in sensitivity between water and organic solvent.22−24 We hypothesize that the minimal solvent dependence of electronic effects in hydricity is due to the delocalization of charge in the Ir complexes.9 NBO analysis (Figure S39) suggests that the charge on Ir is similar in CH3CN and H2O, in accord with this hypothesis. The lack of electronic sensitivity to solvent has important implications in catalyst design: When moving from one solvent to another, the magnitude of hydricity values change, but the relative hydricity for a series of electronically tuned complexes does not. This suggests that, at least for some systems, the influence of substituents on hydricity can be established in one solvent and extrapolated to other solvents. Solvent-Dependent Ligation Effects on Hydricity. The role of chloride additives in tuning hydricity was examined next, as an alternative way to tune hydricity without changing the bipyridine substituents. Previous studies in H2O showed that chloride binding promotes hydride transfer by enhancing the ef fective hydricity, which is a composite of the M−H heterolysis and the free energy of ligand association (Scheme 1).8

Figure 3. CVs of 2 mM [Cp*Ir(bpy)(NCCH3)]2+ in 100 mM [TBA][PF6] in CH3CN (black) and 2 mM [Cp*Ir(bpy)(Cl)]+ in 100 mM [TBA][Cl] in CH3CN with decamethylferrocene (Me10Fc) internal reference (red). Conditions: glassy carbon working electrode, Ag wire pseudoreference electrode, Pt wire counter electrode, and 100 mV/s scan rate.

The influence of chloride can now be compared quantitatively between CH3CN and H2O solvents. Added chloride changes the effective hydricity by 10.6 kcal/mol in CH3CN, whereas in H2O, the effective hydricity only changes by 4.5 kcal/mol (Scheme 1).8 Thus, while ligand electronics influence hydricity similarly in CH3CN and H2O, the impact of chloride ions on the effective hydricity is twice as great in CH3CN as it is in H2O. It is also noteworthy that chloride binding tunes hydricity much more than bpy electronic tuning in both solvents; the change from chloride binding is 3.7 times larger than the full span of electronic effects in CH3CN. The relative hydricity with chloride binding was obtained for the full series of Ir complexes through hydride transfer equilibria (Figure S54). A Hammett plot again shows a linear correlation and is qualitatively reproduced by DFT computations (Figure S56). The slopes are similar for Ir-NCCH3 and Ir−Cl complexes, suggesting that the 1+/2+ charge difference between complexes does not significantly impact the solvent influence on electronic effects and that the chloride association energy is not greatly influenced by bpy substituents. Chloride binding tunes the effective hydricity in a solvent-dependent manner without perturbing the bpy electronic influence. Thermochemical Analysis of Solvent-Dependent Hydricity. Having greatly increased the number of complexes with well-defined hydricity values in two different solvents, we sought to better understand the general solvent dependence of hydricity and develop a quantitative equation to predict changes in hydricity between solvents. A general thermochemical cycle for assessing hydricity in different solvents is shown in Scheme 2. As first noted qualitatively by Creutz,4 and recognized elsewhere,28 hydricity values in two different solvents are connected by transfer free energies (ΔG°tr). Solvent binds after hydride transfer in many cases, so the term ΔG°tr(AL(n+1)+) also includes the free energy of solvent association (cases where solvent does not bind are also accommodated, see Supporting Information section VIII). All thermochemical data take standard states of 1 M solutes, 1 atm gases, and constant solvent activity.2

Scheme 1. Thermochemical Cycle Relating Effective Hydricity with Chloride Binding (ΔG°H−(Cl)), Hydricity with Solvent Binding (ΔG°H−(solv)), and the Free Energy of Chloride Ion Binding (ΔG°solv→Cl)

We set out to establish the effective hydricity with chloride binding, ΔG°H−(Cl) in CH3CN, by summing ΔG°H−(NCCH3) and the free energy of CH3CN substitution by chloride (ΔG°NCCH3→Cl). Electrochemical methods proved useful for the experimental determination of ΔG°NCCH3→Cl. Because the 2e− reduction of [Cp*Ir(bpy-X)L]2+ is coupled to the dissociation of L,8,25 a thermochemical cycle relates the difference in reduction potentials and ΔG°NCCH3→Cl (Supporting Information section IV). Cyclic voltammograms (CVs) of [Cp*Ir(bpy)(Cl)]+ in CH3CN with 0.1 M [TBA][Cl] electrolyte (TBA is tetra-nbutylammonium) exhibit a quasi-reversible 2e− reduction with a half-wave potential E1/2= −1.30 V vs Fc+/0 (ΔEp = 173 mV). The reduction event was previously shown to involve an electrochemical−chemical−electrochemical (ECE) mechanism.25,27 The chloride ions in the electrolyte facilitate rapid halide binding upon reoxidation,26 thus producing a reversible wave in the CV (in contrast to the irreversible wave with [TBA][PF6] electrolyte). The chloride complex is more difficult to reduce than the nitrile complex (E1/2= −1.07 V vs Fc+/0),10 as in Figure 3. This potential difference gives ΔG°NCCH3→Cl = −10.6 kcal/mol (Supporting Information section IV). Adding this value to ΔG°H−(NCCH3) gives the effective hydricity with chloride binding ΔG°H−(Cl) = 51.4 kcal/mol for [Cp*Ir(bpy)(H)]+. C

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Organometallics Scheme 2. Thermochemical Cycle Relating Hydricity across Solventsa

Table 1. Hydride Transfer Free Energy and H2 Acidity/ Hydricity as a Function of Solvent solvent CH3CN methanol ethanol ethylene glycol DMSO DMF N-methyl-2-pyrrolidone ethylene carbonate tetrahydrofuran H2O

a

In H2O (blue) and organic solvent (red). ΔG°H −(org) − ΔG°H −(H 2O) transfer free energies

solvent corr. factor

13.8 7.0 8.6 5.9 18.6 24.6b 22.8b 18.7b 16.4 0

76.0 (62.3)a 43.3 45.4 41.6 60.7 (48.1)a 55.6b

68.7 34.2

a

Parenthetical value, based on the recently developed aqueous scale, recommended for cross-solvent comparisons. bIncreased uncertainty due to unavailable F− transfer free energy.

    = ´Δ G°trÖ≠ÖÖÖÖÖÖÖÖÖÖÖÖ (H−)ÖÆ + ΔG°tr (AL(n + 1) +) − ΔG°tr (AHn +) +  GÖ≠°ÖÖÖÖÖÖÖÖ SCFÖÆ ÖÖÖÖÖÖÖÖÖÖÖÖ ´Δ ÖÖÖÖÖÖÖÖ constant

ΔG°tr,aq→org(H−) (kcal/mol) ΔG°H2 (kcal/mol)

constant

(1)

(ΔG°H2) was calculated as a point of reference (Table 1). These parameters can act as anchors of hydricity scales for the organic solvents for which experimental thermodynamic hydricity values have not yet been explored. The parameters for tetrahydrofuran are particularly noteworthy because hydricity measurements have been carried out in this solvent,37 and the ΔG°H2 value agrees nicely with the H2 acidity in THF reported by Morris and coworkers.38 While most of the organic solvents in Table 1 have not been the subject of extensive studies, Parker did report free energies for H2 heterolysis in DMF and DMSO,33 and the DMSO value was used in an initial study of that solvent.5 We recommend using ΔG°H2 = 48.1 kcal/mol for future studies, because this value is based on the accepted aqueous parameters and will enable seamless cross-solvent comparisons. If the previous value is used, a solvent correction factor analogous to the one for CH3CN would be required for comparisons across solvents. The difference in hydricity of [Cp*Ir(bpy-X)H]+ between CH3CN and H2O was predicted for each Ir complex using eq 1. While experimental methods could, in principle, be used to determine the transfer free energies needed for eq 1, we elected to compute the transfer free energies using DFT methods. Recent computational advances in solvation methods,40,41 including our own contributions in the area of hydricity,42 suggested that ion transfer free energies could be estimated with reasonably accuracy. Computed free energies were converted to the standard state conventions adopted for experimental measurements.43,44 For full computational details, see Supporting Information section II. Each Ir species is computed to be more stable in CH3CN than in H2O, but the ΔG°tr values are similar and thus constitute a relatively minor contribution to the change in hydricity. Thus, the main contribution to the change in hydricity for these complexes is the transfer free energy of the free hydride ion, H−. As shown in Table 2, eq 1 provides remarkably accurate predictions for the Ir complexes, within the estimated 1.4 kcal/mol uncertainty in experimental hydricity differences. To test the generality of the method, the same approach was applied to the small collection of hydride complexes with experimental hydricity values reported in both CH3CN and H2O. Excellent agreement was again observed across a series of hydride complexes varying the metal ion, overall complex charge, supporting ligands, and propensity to bind solvent

The thermochemical cycle of Scheme 2 was used to develop eq 1, a generally applicable expression for predicting the difference in hydricity between solvents. Equation 1 is based on the transfer free energy of the hydride donor (ΔG°tr(AHn+)) and the hydride acceptor (ΔG°tr AL(n+1)+)) as well as the transfer free energy of the free hydride ion, H−, (ΔG°tr(H−)) and a solvent correction factor (ΔG°SCF). The transfer free energies of the hydride donor and acceptor could be measured experimentally or estimated computationally. These parameters will be variables, leading to scales that do not correlate linearly with each other as a function of solvent, as observed for pKa scales in different solvents.29 The values ΔG°tr(H−) and ΔG°SCF, however, are constants that are specific to each solvent. The first value sought was the transfer free energy of the hydride ion (H−), ΔG°tr(H−), which was estimated using Kelly and Rosseinsky’s approach with the most accurate experimental data available.30 Correlations between halide ionic radii, hydration energies, and solvation energies provided ΔG°tr,aq→NCCH3(H−) = 13.8 kcal/mol (Supporting Information section VII), similar to prior estimates.4,30,31 The second constant is a solvent correction factor (ΔG°SCF) required because of longstanding thermochemical conventions in CH3CN. The experimental values for the aqueous H+/H• and H•/H− reduction potentials underpin both the H2O and CH3CN hydricity scales,31−34 but different values for these potentials were utilized when the respective scales were developed decades apart. The correction factor, ΔG°SCF = 13.75 kcal/mol, accounts for these inherent differences without the disruption of changing the widely adopted CH3CN scale. Although the present focus is on CH3CN, eq 1 can be extended to other organic solvents if the appropriate ΔG°tr(H−) values are available. This is significant in light of the broad range of solvents used in catalysis involving hydride transfer. To enable broader future application of eq 1, ΔG°tr(H−) values for eight other organic solvents were calculated (Table 1, see Supporting Information section VII). The ΔG°tr(H−) values of Table 1 correlate nicely with the acceptor number and ET solvent parameters, but do not show any correlation with relative permittivity (Figures S66−S68). These correlations align with trends observed for the kinetics of hydride transfer to CO2 in various solvents.35,36 For solvents where both ΔG°tr(H−) and the H+ transfer free energy, ΔG°tr(H+), are available, the H2 acidity/hydricity D

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thermochemical grounding complements previously observed empirical correlations between CH3CN and H2O hydricity.14

Table 2. Predicting the Difference in Hydricity between CH3CN and H2O for Several Classes of Hydrides

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CONCLUSIONS The hydricity of a structurally homologous set of transition metal hydride complexes has been studied in two solvents, providing a unique platform for understanding hydride transfer in different media. A summary of the thermochemical parameters related to this series of iridium complexes is given in Table S1. Figure 4A gathers all of the thermodynamic hydricity values discussed here on a unified scale that enables direct

ΔG°H−(NCCH3) − ΔG°H−(OH2) (kcal/mol) complex

expc

calcd (eq 1)

exp − calcd

+

30.5 30.5 30.3 30.2 33.49 31.14 25.54 33.87 33.56 19.92 41.82

30.1 29.7 29.9 30.2 31.5 28.9 27.8 33.0 26.1 16.6b 41.7

0.4 0.8 0.4 0.1 2.0 2.3 −2.3 0.8a 7.4 3.3 0.1

[Cp*Ir(bpy)H] [Cp*Ir(bpy-Me)H]+ [Cp*Ir(bpy-OMe)H]+ [Cp*Ir(bpy-CO2Me)H]+ [RuIrH]3+ [(C6Me6)Ru(bpy)H]+ [Ru(tpy)(bpy)H]+ [HNi(dmpe)2]+ [HFe4(CO)12N]− HCO2− H2 a

When solvent binding is not considered, the difference is 7.25 kcal/ mol. b Assuming HCO 2 − and acetate have the same value ΔG°tr,aq→MeCN = +11 kcal/mol (see Supporting Information section VII for details).28,39 cExperimental differences in hydricity have estimated uncertainty of ±1.4 kcal/mol.

(Table 1). In all cases, the DFT-aided predictions using eq 1 were more accurate (often by several kcal/mol) than those predictions based only on the constants ΔG°tr(H−) and ΔG°SCF (Table S22). Hydricity changes for the small molecules HCO2− and H2 were also calculated using eq 1.28,39 The hydricity of HCO2− differs by only 20 kcal/mol in water (ΔG°H− = 24.1 kcal/ mol)34 and acetonitrile (ΔG°H− = 44 kcal/mol).45,46 The small difference in hydricity relative to transition metal hydride complexes has been exploited for solvent-dependent CO2 reduction reactions.4,6,7,13,14 The small solvent dependent hydricity of formate has previously been attributed to Hbonding stabilization of HCO2−.7 Equation 1 reveals that the CO2 being a gas phase hydride acceptor is also an important, previously unrecognized factor. The independent hydricity scales actually prove somewhat misleading in this case: one would conclude that formate is a 20 kcal/mol stronger hydride donor in H2O than CH3CN, but ∼14 kcal/mol of the apparent hydricity difference is due to distinct scale conventions in the two solvents (ΔG°SCF). Thus, the hydride donor ability of HCO2−is almost the same in H2O and CH3CN (∼6 kcal/mol difference in hydricity). The fact that H2 is a gas-phase hydride donor, in contrast to CO2 which is a gas-phase hydride acceptor, helps to explain the large solvent-dependent hydricity for this other quintessential small molecule. Even with corrections for scale conventions, the hydricity of H2 changes by ∼30 kcal/mol between CH3CN and H2O. Equation 1 is thus able to accurately predict the hydricity changes for both transition metal and small-molecule hydride donor/acceptor pairs. This general approach with a strong

Figure 4. Effective hydricity (ΔG°H−) of hydride complexes [Cp*Ir(bpy-X)(H)]+ plotted on a unified scale against Σσp− (A) provides an overview of how hydricity can be tuned. Plotting effective hydricity in acetonitrile vs water (ΔG°H−(NCCH3) vs ΔG°H−(H2O)). (B) Series of related Ir complexes described here are related by a constant value, whereas the more diverse collection of hydride donors with hydricity known in each solvent are not linearilty related.

comparisons across solvents. Hydricity can be tuned by changing the solvent, changing the ligand electronic structure, or introducing a chloride ligand to bind after hydride transfer. The Ir hydrides can be made more hydridic by incorporating electron-donating groups in the backbone of the bipyridine ligand, although this effect is relatively small for these Ir complexes. In contrast to established trends in organic acidity,22−24 the iridium complexes have the same electronic E

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ACKNOWLEDGMENTS This work was supported by the NSF Center for Enabling New Technologies through Catalysis (CENTC) through grant CHE-1205189 (synthesis and NMR studies) and the Division of Chemical Sciences, Geosciences & Biosciences, Office of Basic Energy Sciences of the U.S. Department of Energy through grant DE-SC0014255 (electrochemistry and thermodynamic analysis). K.R.B. acknowledges support from a University of North Carolina Dissertation Completion Fellowship.

sensitivity in both CH3CN and H2O. The hydricity also increased when chloride binds after hydride release. Both of these trends hold with chemical intuition, and the quantitative analysis reveals that the chloride association energy is much higher in CH3CN than in H2O, which leads to a dramatic change in the ability of chloride to tune effective hydricity in the different solvents. The change in hydricity moving from acetonitrile to water is quantitatively larger than the other tuning methods discussed here, although it is important to emphasize the importance of relative hydricity within a given solvent scale. Figure 4B plots the hydricity in acetonitrile against the hydricity in water for H2, HCO2−, and other metal hydrides with hydricities reported in each solvent. There is not a strong correlation in the combined data, as expected from eq 1. There is substantial compression in the aqueous hydricity scale, as observed with solvent-dependent pKa scales.29 Considering this nonlinear cross-solvent relationship, it is impressive that a quantitative analysis of solvent dependence is provided by eq 1. Hydricity measurements made in one solvent can be extrapolated to others. The most accurate predictions require experimental or computational estimates of transfer free energies. However, as can be seen from the linear relationship for the series of Ir complexes studied here, it is expected that the same transfer free energy values can be assigned to each complex in a structurally analogous series if needed. The ability to predict how hydricity will change from one solvent to another provides an important tool for catalyst design, enabling predictions of whether relative hydricity of two species will switch as a function of solvent and become thermodynamically favorable.

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REFERENCES

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EXPERIMENTAL SECTION

Experimental details are presented in the Supporting Information.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.organomet.9b00278. Experimental, crystallographic, and computational details (PDF) Coordinates for DFT optimized geometries (XYZ) Accession Codes

CCDC 1885711 and 1885712 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Alexander J. M. Miller: 0000-0001-9390-3951 Notes

The authors declare no competing financial interest. F

DOI: 10.1021/acs.organomet.9b00278 Organometallics XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.organomet.9b00278 Organometallics XXXX, XXX, XXX−XXX