Calorimetric Study of the Activation of Hydrogen by Tris

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Cite This: J. Phys. Chem. A XXXX, XXX, XXX-XXX

Calorimetric Study of the Activation of Hydrogen by Tris(pentafluorophenyl)borane and Trimesitylphosphine Adrian Y. Houghton and Tom Autrey* Pacific Northwest National Laboratory, 902 Battelle Boulevard, P.O. Box 999, Richland, Washington 99352 United States S Supporting Information *

ABSTRACT: The mechanism of H2 heterolysis by the frustrated Lewis pair of B(C6F5)3 and P(mes)3 was investigated by isothermal reaction calorimetry in the temperature range from 30 to 90 °C. The experimental heat curves were modeled in Berkeley Madonna to obtain both kinetic and thermodynamics data simultaneously. The H2 activation reaction is treated as a single, termolecular step with a rate constant of 0.61 M−2 s−1 at 303 K with an exothermic enthalpy of reaction, ΔHH2 = −141 kJ/mol. An Eyring analysis gave activation parameters of ΔH‡ = 13.6(9) kJ mol−1 and ΔS‡ = −204(85) J K−1 mol−1. Using D2 gas in place of H2 gas provided an opportunity to measure the relative rates of D2 versus H2 heterolysis to yield a the kinetic isotope effect, KIE = 1.1(1).

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INTRODUCTION Lewis acid−base pairs that are sterically or otherwise prevented from forming a normal dative adduct, “frustrated Lewis pairs” (FLPs),1−7 have received tremendous attention in the past decade due to their notable ability to activate small molecules. These unexpected findings have found applications in catalysis,8−14 organic synthesis,15−18 and CO2 capture19−23 and have introduced a new paradigm for conceptualizing chemical reactivity. The notion that three moieties, each of which appears inert to the other two, can so frequently be expected to react together has provided chemists with an extraordinary new toolset. Even though termolecular reactions involving small molecules have been known for some time,24−27 none has led to such a renaissance as frustrated Lewis pairs. One of the earliest observations in this regard was the activation of H2 by tris(pentafluorophenyl)borane and a PR3 (R = tBu, mesityl).28 To avoid proposing that all FLP reactions are termolecular, three possible “two-step” mechanisms (Scheme 1) are noted: (I) the reaction proceeds through a borane−H2 adduct, which is deprotonated by the phosphine, (II) the reaction proceeds through a phosphine−H2 adduct, followed by hydride abstraction by the borane, and (III) the reaction proceeds through a van der Waals complex between the acid and base, followed by a concerted, heterolytic cleavage of H2. The plausibility of the H2 interacting with the empty porbital of highly Lewic-acidic boranes has some justification in the literature: the spectroscopic detection of the H2−BH3 adduct in an argon matrix,29 the σ-bond metathesis reaction between H2 and HB(C6F5)2,30 and the activation of H2 by boroles.31,32 Furthermore, density functional theory (DFT) calculations indicate that mechanism I is operative where the acid/base pair is HBArF2 (ArF = 2,4,6-tris(trifluoromethyl)phenyl) and NEt3.33 Phosphine−H2 interactions have also been detected in an argon matrix,34 and the activation of hydrogen by a diazadiphosphapentalene has been reported recently.35 © XXXX American Chemical Society

Scheme 1

However, a seminal theoretical study by Pápai and co-workers indicated that both H2−B(C6F5)3 and H2−PR3 interactions are repulsive, and discovered that weakly bound B(C6F5)3−PR3 complexes exist at the minima of their potential energy surfaces in the gas phase.36 The nonlinked FLP encounter complexes have been the subject of intense theoretical investigation36−44 and are found to be stabilized in the gas phase. However, in solution the stabilizing interactions between the acid and base pair are significantly reduced due to competing solute−solvent interactions.45 For example, Dang43 used MD to compute the free energy surface of lutidine and B(C6F5)3 in the gas phase and in CH2Cl2 and found the solvent separated acid base pair to be more stable than the solvent cage pair. Using classical MD with explicit solvent, Pápai and co-workers found the formation Received: August 28, 2017 Revised: October 24, 2017 Published: October 27, 2017 A

DOI: 10.1021/acs.jpca.7b08582 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A of the t(Bu)3P/B(C6F5)3 encounter complex, previously found to be stable in the gas phase (ΔE ca. −54 kJ/mol)36 and to be unstable by ca. 5 kJ/mol in toluene.36 Consistent with the MD simulations the HOESY NMR study by Rocchigiani et al. determined that the formation of the B(C6F5)3/P(mes)3 encounter complex has a ΔG°(298) of ca. +1.7 kJ/mol.46 It was also found that these encounter complexes did not exhibit any preferential orientation,38 although theoretical studies indicate that the phosphine needs to be oriented toward the borane in the solvent cage for the H2 activation reaction to occur. The nature of hydrogen activation has been probed experimentally by several groups.47−50 Notably, the Piers group discovered that P(tBu)3 and B(C6F5)3 do indeed react with each other to form isobutene, [(tBu)3PH][FB(C6F5)3], and an intramolecular phosphinoborane, adding to the kinetic complexity of hydrogen activation.47 The thermodynamics have been shown to be very tunable using different combinations of acids50,51 and bases,52 in line with theoretical predictions.42 However, detailed kinetic information about basic three-bodied FLP H2 activations is still absent in the literature. We therefore chose to study the reaction between B(C6F5)3, P(mes)3, and H2, using isothermal reaction calorimetry, and report here the activation parameters, kinetic isotope effect, and orders in each reactant.

solutions of B(C6F5)3 (1.0 mL) and P(mes)3 (1.0 mL), and a reference cell was charged with CH2Cl2 (2.0 mL). The cells were pressurized with H2 (100−400 psig) and placed in the calorimeter at the desired reaction temperature (30−90 °C). Once heatflow had stabilized near zero, the reaction was initiated by reversal mixing, and data points were collected every 1.2 s until heat flow returned to the baseline. The degree of reaction was confirmed by 19F and 31P NMR spectroscopy.

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RESULTS AND DISCUSSION Calorimetry has been successfully applied to the study of several FLP systems48,49,53 and offers the advantages of a simple experimental setup and the ability to collect thousands of data points per experiment. The relationship between heat flow and reaction rate is given by eq 1: ∂Q R ∂t

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∂Q (t ) e =A ∂t

∂C ∂t

(1)

where ∂QR/∂t is the heat flow of the reaction (in mW), V is the volume of the reaction, ΔH is the enthalpy, and ∂C/∂t is the rate of the reaction. However, the output of the calorimeter does not respond instantaneously to the heat generated by the process under study and must be corrected using the Tian equation:54 ∂Q R

METHODS General Considerations. All manipulations were carried out under an inert nitrogen (N2) atmosphere using standard Schlenk or glovebox techniques unless otherwise stated. Dichloromethane was passed through a neutral alumina column under argon prior to use. Tris(pentafluorophenyl)borane was obtained from Boulder Scientific and purified by sublimation under reduced pressure at 90 °C. All NMR spectra were recorded on 500 MHz Varian INOVA spectrometers. 19F NMR spectra were referenced to fluorobenzene as an external standard (δ = −113.15 ppm). Calorimetric measurements were performed on a Setaram C80 Calvet calorimeter, and the instrument was operated in isothermal mode. Measurements were conducted in modified Hastealloy reversal mixing cells (6.7 mL). The commercial mixing vessels were modified to include an inlet with 0.030 in. i.d. PEEK tubing that allowed gases to be introduced. Instrument Time Constant Determination. The two compartments of a C80 cell were charged with CH2Cl2 solutions of B(C6F5)3 (51.2 mg, 0.100 mmol, 1.0 mL) and pyrazine (80.1 mg, 1.00 mmol, 1.0 mL), and a reference cell was charged with CH2Cl2 (2.0 mL). The cells were pressurized with N2 (100 psig) and placed in the calorimeter. Once heat flow had stabilized near zero, the reaction was initiated by reversal mixing. This experiment was carried out at 29.1 and 88.5 °C, and the heat curves were fit to the following formula: −(t − to)/ τ1

= V ΔH

∂t

=

∂Q C ∂t

+ tc

∂ 2Q C ∂t 2

(2)

where ∂QC/∂t is the observed heat flow and tc is the instrument time constant, which is a measure of how long the signal takes to return to the baseline after an instantaneous thermal event takes place (Supporting Information). In a typical experiment a cylindrical steel cell was charged with dichloromethane solutions of B(C6F5)3 and P(mes)3 that are kept separate by a deep inner well, and a reference cell was charged with dichloromethane only. The cells were then pressurized with H2 (typically 6.89−27.6 bar, in at least 13-fold excess) via a modified cell cap equipped with PEEK tubing (0.03 in. i.d.) and allowed to equilibrate in a Setaram C80 Calvet calorimeter until the heat flow stabilized. Reversal mixing was used to initiate the reaction, and to keep the solution saturated with H2 with approximately six inversions per minute. Once the heat flow had returned to the baseline, the completion of the reaction was verified by 19F and 31P NMR spectroscopy. We first sought to determine the orders of reaction in each of phosphine (P), borane (B), and H2, because this could potentially distinguish among the three mechanistic possibilities discussed above. Applying the steady state approximation for each intermediate in the three mechanisms, one obtains the following rate laws:

−(t − to)/ τ2

−e 1/τ2 − 1/τ1

where ∂Q(t)/∂t is the heat flow at time t, τ1 is the mixing time, τ2 is the instrument time constant, and A is a fitting parameter. τ1 = 15.1 s; τ2 = 310 s (29.1 °C), 278 s (88.5 °C). It was assumed that τ2 varied linearly over the experimental temperature range. C80 Experiments. In a typical experiment, the two compartments of a C80 cell were charged with CH2Cl2

k k [B][P][H 2] ∂C = 1 2 ∂t k −1 + k 2[H 2]

(3)

k k [B][P][H 2] ∂C = 1 2 ∂t k −1 + k 2[B]

(4)

k k [B][P][H 2] ∂C = 1 2 ∂t k −1 + k 2[P]

(5)

It bears mentioning that the steady state approximation implies that the first step (intermediate formation) is rate limiting, B


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The Journal of Physical Chemistry A

the product mixtures indicated the presence of small amounts (ca. 1−5 mol %) of the water activation product [HP(mes)3][(μ-OH)(B(C6F5)3)2].56 A separate calorimetry experiment with stoichiometric quantities of P(mes)3, B(C6F5)3, and H2O revealed that the reaction is effectively instantaneous, with a ΔH = −176 kJ/mol H2O. Second, the nonlinearity of the red curve in Figure 1b indicates that the reaction is not adequately captured by the above mechanistic descriptions. 19F NMR spectrum of the product mixture showed significant broadening of the signals, which has previously been attributed to the rapid, reversible formation of [(μ-(H)(B(C6F5)3)2]−.57 Addition of a few drops of MeCN-d3 to the sample resulted in the sharpening of all signals in the sample, consistent with the formation of the MeCN−B(C6F5)3 adduct and release of [HB(C6F5)3]− anion. Given the large excess of B(C6F5)3 used to create pseudo-firstorder conditions in phosphine, this equilibrium likely played a significant role in this experiment. It also bears mentioning that the mixing of solutions of B(C6F5)3 and P(mes)3 produces no appreciable heat in the absence of hydrogen. To gauge the impact of the formation of [(μ-(H)(B(C6F5)3)2]−, we measured the heat output upon mixing equimolar B(C6F5)3 amounts of and [HP(mes)3][HB(C6F5)3], which amounted to only −1.9 kJ/mol B(C6F5)3. The equilibrium appears to lie far to the left in any case, based on the work done by Collins et al.57 Furthermore, comparing fractional conversion of the pseudo-first-order experiments to a pseudo-second-order experiment in which [B] ≈ [P] reveals approximately equal order dependence on both phosphine and borane (Figure 2). Again, the slight differences in the two

whereas DFT calculations predict that the second step (hydrogen activation) is the rate-limiting step.55 In this case, the above three rate laws would collapse into a single, indistinct rate law (i.e., k2[X] ≪ k−1): kk ∂C = 1 2 [B][P][H 2] ∂t k −1

(6)

Two experiments in which pseudo-first-order conditions in [P] and [B] were carried out; the excess reagent (phosphine or borane) and hydrogen were held in approximately 10-fold and 66-fold excess, respectively, over the limiting reagent. Assuming that the only species in appreciable, changing amounts were the product and limiting reagent, the concentration of product as a function of time C(t) was estimated from the ratio of heat evolved Q(t) over the total heat from the reaction Qtotal times the initial concentration of limiting reagent [L]o: C(t ) = [L]o

Q (t ) Q total

(7)

At first glance, graphical analysis of the heat flow data (Figure 1) indicates that the reaction is first order in borane and ca. 0.65 order in phosphine. This would appear to be consistent with mechanism 2 (Scheme 1); however, there are several factors that bear consideration. First, the rapid deceleration in the early stages of the reaction (blue curves) indicates the occurrence of a rapid side reaction. Indeed, 19F NMR spectra of

Figure 2. Fractional conversions (Q(t)/Qtotal) for the pseudo-first- and second-order experiments in which H2 was held in excess at ca. 6.9 bar at 30 °C in dichloromethane.

pseudo-first-order experiments are attributable to the side reaction with water and to the reversible formation of [(μ(H)(B(C6F5)3)2]−. Pressure dependence experiments were carried out with equimolar B(C6F5)3 and P(mes)3 and 6.92, 13.7, 20.7, and 27.6 bar H2, and the heat traces are overlaid in Figure 3. It had been hoped that by increasing the pressure, the rate-limiting step would change and we would be able to distinguish between rate laws 1, 2, and 3. However, the absence of a change in rate limiting step at high hydrogen pressures leads us to treat the reaction as a single termolecular step (rate law 4). There is

Figure 1. Logarithmic plots for pseudo-first-order reactions (a) in B(C6F5)3 and (b) in P(mes)3 at 30 °C and 6.9 bar H2 in dichloromethane. The limiting reagent was kept at 10 mM and the excess reagent at 100 mM. Data points included in the linear best fit are in red. C


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Figure 4. Heat curves and model fits for variable temperature experiments at p(H2)o ≈ 6.9 bar and [B(C6F5)3]o ≈ [P(mes)3]o = 50 mM in CH2Cl2. Inset: Eyring plot of variable temperature experiments.

Figure 3. Heat curves and model fits for variable pressure experiments at 30 °C and [B(C6F5)3]o ≈ [P(mes)3]o = 50 mM in CH2Cl2.

It is instructive to compare the activation parameters for heterolysis of H2 in this bimolecular FLP to the activation parameters in a unimolecular FLP. Repo and colleagues found a near identical activation free energy for heterolysis of hydrogen by o-TMP−C6H4−BH2, ΔG‡ ca. 77 kJ/mol.60 The entropy ΔS‡ = −90 kJ/mol is typical for a bimolecular reaction, however, this comes at an apparent cost of an increase in the enthalpy, ΔH‡ ca. 50 kJ/mol, compared to 14 kJ/mol determined in this work. The similarity in ΔG‡ is a little surprising given that the heterolysis by o-TMP−C6H4−BH2 is near thermoneutral compared to the exothermic heterolysis by B(C6F5)3 and P(mes)3. This may be due to an inherent difference in the reactivity of a nitrogen vs a phosphorus center Lewis base. Certainly more research to understand the surprisingly low enthalpic nature of the transition state in FLP activation of H2 is desirable. Finally, we carried out the reaction using D2 to measure the kinetic isotope effect (KIE) and found that kH2/kD2 = 1.1(1). Given that H2 activation appears to be rate-limiting, the nearunity of the KIE might indicate that the true rate-limiting step is the diffusion of H2(D2) into the “encounter complex pocket”, and that hydrogen cleavage follows in a second step. However, the nonzero enthalpy of activation indicates that there must be some degree of bond breaking and making in the transition state. Examples of KIE measurements of the termolacular activation of H2 are rare, but there is an instructive example that comes from the work of Wayland et al.:25 the homolytic cleavage of H2 by RhII(TMP)•, which has activation parameters similar to those of the reaction under study (vide supra), was shown to have a KIE of only 1.6. Indeed, two RhII(TMP)• molecules can be regarded as a “frustrated radical pair,” and the activation of H2 almost certainly proceeds through a fourcentered, linear transition state (as with B(C6F5)3/P(mes)3). Such a transition state can generally be expected to result in a low KIE because it contains many more isotopically sensitive vibrational modes relative to its ground state, as compared to a three-centered system. Consider the following expressio,n which estimates a KIE as a ratio of the ZPE terms for isotopically sensitive vibrations in the reactants over those in the transition state (eq 8). The ∏ symbol indicates the product of terms i through n, and μi = hυi/ kBT, where υi is the frequency of vibration i.61,62

experimental precedence for molecular hydrogen activation by organometallic radical pairs in solution (vide inf ra). Notably, a recent report from Stephan et al. indicates that the B(C6F5)3/ P(mes)3 pair exhibits radical character in solution;58 thus there may be some further parallel reactivity with organometallic radical pairs. With the side reactions in mind, we opted to explicitly model the whole system using Berkeley Madonna. Inclusion of the reversible formation of [(μ-H)(B(C6F5)3)2]− in the model had little impact on the fit and was therefore left out. The side reaction with water was included as an “instantaneous” reaction, and virtually no heat was evolved upon mixing solutions of B(C6F5)3 and P(mes)3 in the absence of H2. The explicit model can be seen in the Supporting Information, and fits to the data were achieved by allowing “k3”, “DH3” and Co terms to vary. We measured the rate constant to be 0.613(9) M−2 s−1 and the ΔH = −141(2) kJ mol−1. Previously, our group had measured an enthalpy of hydrogen activation by P(tBu)3, B(C6F5)3, as being ΔH = −131(1) kJ mol−1; in these experiments we correct for heat associated with H2 dissolving dichloromethane, ΔH = 10.4 kJ mol−1, which was calculated from variable temperature solubility data.59 Variable temperature experiments at 30, 50, 70, and 90 °C allowed us to obtain the activation parameters from an Eyring analysis of the rate constants, which ranged from 0.613(9) to 1.77(2) M−2 s−1. The fitted heat traces and Eyring plot can be observed in Figure 4, with activation parameters of ΔH‡ = 13.6(9) kJ mol−1 and ΔS‡ = −204(85) J K−1 mol−1, corresponding to ΔG‡(298) = 74.4 kJ mol−1. The ΔS‡ compares well to other termolecular H2 activation reactions, namely homolytic cleavage by metalloradicals like •Co(CN)53− (ΔS‡ = −230(20) J K−1 mol−1),27 RhII(TMP)• (TMP = tetramesitylporphyrin) (ΔS‡ = −170(20) J K−1 mol−1),25 and Cp*Cr(CO)3• (ΔS‡ = −200(20) J K−1 mol−1),26 and ΔH‡ is reasonably close to that for RhII(TMP)• (ΔH‡ = 20(4) kJ mol−1).25 Prior DFT studies predicted that ΔH‡ = 104 kJ mol−1, ΔS‡ = 47.6 J K−1 mol−1, and ΔG‡(298) = 137 kJ mol−1 in toluene,55 though it is unclear why these calculations are so vastly different from measurements; one would expect the reaction to be several orders of magnitude slower with an activation free energy as high as predicted by DFT. D



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r

∏ e−(1/2)(uiH − uiD)r kH = ‡i −(1/2)(u − u ) iH iD ‡ kD ∏i e

(8)

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CONCLUSIONS In summary, we used isothermal reaction calorimetry to perform a detailed kinetic study of the activation of H2 by B(C6F5)3 and P(mes)3 in dichloromethane. We found the reaction to be well modeled as a single, termolecular reaction step. The activation parameters indicated that the reaction was entropy controlled, had a surprisingly small enthalpic barrier, and had virtually no kinetic isotope effect. These results are consistent with a rate-determining step involving the assembling the reactants into the solvent cage with right configuration rather than breaking the H−H bond/forming the B−H/N−H bonds. ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b08582. General information for the time constant determination and C80 model, the Berkeley Madonna Model, and figures of heat traces (PDF)

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REFERENCES

(1) Frustrated Lewis Pairs I, 1st ed.; Springer-Verlag: Berlin, 2013. (2) Frustrated Lewis Pairs II, 1st ed.; Springer-Verlag: Berlin, 2013. (3) Stephan, D. W.; Erker, G. Frustrated Lewis Pair Chemistry: Development and Perspectives. Angew. Chem., Int. Ed. 2015, 54, 6400−6441. (4) Stephan, D. W.; Erker, G. Frustrated Lewis Pairs: Metal-free Hydrogen Activation and More. Angew. Chem., Int. Ed. 2010, 49, 46− 76. (5) Stephan, D. W. Frustrated Lewis Pairs: From Concept to Catalysis. Acc. Chem. Res. 2015, 48, 306−316. (6) Rocchigiani, L. Experimental Insights into the Structure and Reactivity of Frustrated Lewis Pairs. Isr. J. Chem. 2015, 55, 134−149. (7) Stephan, D. W. The broadening reach of frustrated Lewis pair chemistry. Science 2016, 354, aaf7229. (8) Hounjet, L. J.; Stephan, D. W. Hydrogenation by Frustrated Lewis Pairs: Main Group Alternatives to Transition Metal Catalysts? Org. Process Res. Dev. 2014, 18, 385−391. (9) Chase, P. A.; Welch, G. C.; Jurca, T.; Stephan, D. W. Metal-Free Catalytic Hydrogenation. Angew. Chem., Int. Ed. 2007, 46, 8050−8053. (10) Erő s, G.; Nagy, K.; Mehdi, H.; Pápai, I.; Nagy, P.; Király, P.; Tárkányi, G.; Soós, T. Catalytic Hydrogenation with Frustrated Lewis Pairs: Selectivity Achieved by Size-Exclusion Design of Lewis Acids. Chem. - Eur. J. 2012, 18, 574−585. (11) Mohr, J.; Oestreich, M. B(C6F5)3-Catalyzed Hydrogenation of Oxime Ethers without Cleavage of the NO Bond. Angew. Chem., Int. Ed. 2014, 53, 13278−13281. (12) Paradies, J. Metal-Free Hydrogenation of Unsaturated Hydrocarbons Employing Molecular Hydrogen. Angew. Chem., Int. Ed. 2014, 53, 3552−3557. (13) Segawa, Y.; Stephan, D. W. Metal-free hydrogenation catalysis of polycyclic aromatic hydrocarbons. Chem. Commun. 2012, 48, 11963− 11965. (14) Hounjet, L. J.; Bannwarth, C.; Garon, C. N.; Caputo, C. B.; Grimme, S.; Stephan, D. W. Combinations of Ethers and B(C6F5)3 Function as Hydrogenation Catalysts. Angew. Chem., Int. Ed. 2013, 52, 7492−7495. (15) McCahill, J. S. J.; Welch, G. C.; Stephan, D. W. Reactivity of “Frustrated Lewis Pairs”: Three-Component Reactions of Phosphines, a Borane, and Olefins. Angew. Chem., Int. Ed. 2007, 46, 4968−4971. (16) Dureen, M. A.; Stephan, D. W. Terminal Alkyne Activation by Frustrated and Classical Lewis Acid/Phosphine Pairs. J. Am. Chem. Soc. 2009, 131, 8396−8397. (17) Mahdi, T.; Stephan, D. W. Frustrated Lewis Pair Catalyzed Hydroamination of Terminal Alkynes. Angew. Chem., Int. Ed. 2013, 52, 12418−12421. (18) Melen, R. L.; Hansmann, M. M.; Lough, A. J.; Hashmi, A. S. K.; Stephan, D. W. Cyclisation versus 1,1-Carboboration: Reactions of B(C6F5)3 with Propargyl Amides. Chem. - Eur. J. 2013, 19, 11928− 11938. (19) Ashley, A. E.; Thompson, A. L.; O’Hare, D. Non-MetalMediated Homogeneous Hydrogenation of CO2 to CH3OH. Angew. Chem., Int. Ed. 2009, 48, 9839−9843. (20) Courtemanche, M.-A.; Légaré, M.-A.; Maron, L.; Fontaine, F.-G. Reducing CO2 to Methanol Using Frustrated Lewis Pairs: On the Mechanism of Phosphine−Borane-Mediated Hydroboration of CO2. J. Am. Chem. Soc. 2014, 136, 10708−10717. (21) Stephan, D. W.; Erker, G. Frustrated Lewis pair chemistry of carbon, nitrogen and sulfur oxides. Chem. Sci. 2014, 5, 2625−2641. (22) Voicu, D.; Abolhasani, M.; Choueiri, R.; Lestari, G.; Seiler, C.; Menard, G.; Greener, J.; Guenther, A.; Stephan, D. W.; Kumacheva, E. Microfluidic studies of CO2 sequestration by frustrated Lewis pairs. J. Am. Chem. Soc. 2014, 136, 3875−80. (23) Chi, J. J.; Johnstone, T. C.; Voicu, D.; Mehlmann, P.; Dielmann, F.; Kumacheva, E.; Stephan, D. W. Quantifying the efficiency of CO2 capture by Lewis pairs. Chem. Sci. 2017, 8, 3270. (24) Norton, J. R.; Spataru, T.; Camaioni, D. M.; Lee, S.-J.; Li, G.; Choi, J.; Franz, J. A. Kinetics and Mechanism of the Hydrogenation of

Qualitatively, this expression indicates that as the number of isotopically sensitive vibrations in the transition state (i.e., the number of terms in the denominator) increase relative to those in the reactants (i.e., the number of terms in the numerator), the KIE will decrease. In the case of a three-centered system of the type X H−Y ⇔ [X--H--Y]‡, the reactants would exhibit one stretch and two bending modes, while the transition state would have two stretching and two bending modes. When this system is idealized as having a perfectly linear, symmetric transition state, the theoretical KIE is about 6.5. The situation is somewhat different for a four-centered transition state of the type X H−H Y ⇔ [X--H--H--Y]‡. The reactants exhibit only one stretching mode, and in the case of Wayland’s system where X = Y = Rh, the transition state has three stretching and four bending modes. In the system under study the transition state is less symmetric, resulting in four stretching and eight bending modes, which is consistent with the smaller KIE than for RhII(TMP)• (1.1 vs 1.6).

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AUTHOR INFORMATION

Corresponding Author

*T.A. E-mail: [email protected]. ORCID

Tom Autrey: 0000-0002-7983-3667 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences is gratefully acknowledged for financial support. The authors thank Dr. Don Camaioni for critical discussions and useful suggestions for fitting strategies in Berkeley Madonna. Pacific Northwest National Laboratory is a multiprogram national laboratory operated by Battelle for the U.S. Department of Energy under Contract DE-AC0576RL01830. . E


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The Journal of Physical Chemistry A the CpCr(CO)3•/[CpCr(CO)3]2 Equilibrium to CpCr(CO)3H. Organometallics 2014, 33, 2496−2502. (25) Wayland, B. B.; Ba, S.; Sherry, A. E. Reactions of hydrogen or deuterium molecule with a rhodium(II) metalloradical: kinetic evidence for a four-centered transition state. Inorg. Chem. 1992, 31, 148−150. (26) Capps, K. B.; Bauer, A.; Kiss, G.; Hoff, C. D. The rate and mechanism of oxidative addition of H2 to the Cr(CO)3C5Me5 radicalgeneration of a model for reaction of H2 with the Co(CO)4 radical. J. Organomet. Chem. 1999, 586, 23−30. (27) Halpern, J.; Pribanic, M. Hydrogenation of pentacyanocobaltate(II) at high pressures. Inorg. Chem. 1970, 9, 2616−2618. (28) Welch, G. C.; Stephan, D. W. Facile Heterolytic Cleavage of Dihydrogen by Phosphines and Boranes. J. Am. Chem. Soc. 2007, 129, 1880−1881. (29) Tague, T. J.; Andrews, L. Reactions of Pulsed-Laser Evaporated Boron Atoms with Hydrogen. Infrared Spectra of Boron Hydride Intermediate Species in Solid Argon. J. Am. Chem. Soc. 1994, 116, 4970−4976. (30) Nikonov, G. I.; Vyboishchikov, S. F.; Shirobokov, O. G. Facile Activation of H−H and Si−H Bonds by Boranes. J. Am. Chem. Soc. 2012, 134, 5488−5491. (31) Houghton, A. Y.; Karttunen, V. A.; Fan, C.; Piers, W. E.; Tuononen, H. M. Mechanistic Studies on the Metal-Free Activation of Dihydrogen by Antiaromatic Pentarylboroles. J. Am. Chem. Soc. 2013, 135, 941−947. (32) Fan, C.; Mercier, L. G.; Piers, W. E.; Tuononen, H. M.; Parvez, M. Dihydrogen Activation by Antiaromatic Pentaarylboroles. J. Am. Chem. Soc. 2010, 132, 9604−9606. (33) Lu, Z.; Cheng, Z.; Chen, Z.; Weng, L.; Li, Z. H.; Wang, H. Heterolytic Cleavage of Dihydrogen by “Frustrated Lewis Pairs” Comprising Bis(2,4,6-tris(trifluoromethyl)phenyl)borane and Amines: Stepwise versus Concerted Mechanism. Angew. Chem., Int. Ed. 2011, 50, 12227−12231. (34) Moroz, A.; Sweany, R. L.; Whittenburg, S. L. Van der Waals complexes between Lewis bases and molecular hydrogen in argon matrixes. J. Phys. Chem. 1990, 94, 1352−1357. (35) Cui, J.; Li, Y.; Ganguly, R.; Kinjo, R. Reactivity Studies on a Diazadiphosphapentalene. Chem. - Eur. J. 2016, 22, 9976−9985. (36) Rokob, T. A.; Hamza, A.; Stirling, A.; Soós, T.; Pápai, I. Turning Frustration into Bond Activation: A Theoretical Mechanistic Study on Heterolytic Hydrogen Splitting by Frustrated Lewis Pairs. Angew. Chem., Int. Ed. 2008, 47, 2435−2438. (37) Grimme, S.; Kruse, H.; Goerigk, L.; Erker, G. The Mechanism of Dihydrogen Activation by Frustrated Lewis Pairs Revisited. Angew. Chem., Int. Ed. 2010, 49, 1402−1405. (38) Bannwarth, C.; Hansen, A.; Grimme, S. The Association of Two “Frustrated” Lewis Pairs by State-of-the-Art Quantum Chemical Methods. Isr. J. Chem. 2015, 55, 235−242. (39) Bako, I.; Stirling, A.; Balint, S.; Papai, I. Association of frustrated phosphine-borane pairs in toluene: molecular dynamics simulations. Dalton Trans. 2012, 41, 9023−9025. (40) Stirling, A.; Hamza, A.; Rokob, T. A.; Papai, I. Concerted attack of frustrated Lewis acid-base pairs on olefinic double bonds: a theoretical study. Chem. Commun. 2008, 3148−3150. (41) Hamza, A.; Stirling, A.; András Rokob, T.; Pápai, I. Mechanism of hydrogen activation by frustrated Lewis pairs: A molecular orbital approach. Int. J. Quantum Chem. 2009, 109, 2416−2425. (42) Rokob, T. A.; Hamza, A.; Pápai, I. Rationalizing the Reactivity of Frustrated Lewis Pairs: Thermodynamics of H2 Activation and the Role of Acid−Base Properties. J. Am. Chem. Soc. 2009, 131, 10701− 10710. (43) Dang, L. X.; Schenter, G. K.; Chang, T.-M.; Kathmann, S. M.; Autrey, T. Role of Solvents on the Thermodynamics and Kinetics of Forming Frustrated Lewis Pairs. J. Phys. Chem. Lett. 2012, 3, 3312− 3319. (44) Camaioni, D. M.; Ginovska-Pangovska, B.; Schenter, G. K.; Kathmann, S. M.; Autrey, T. Analysis of the Activation and Heterolytic

Dissociation of H2 by Frustrated Lewis Pairs: NH3/BX3 (X = H, F, and Cl). J. Phys. Chem. A 2012, 116, 7228−7237. (45) Pu, M. P.; Privalov, T. Chemistry of Intermolecular Frustrated Lewis Pairs in Motion: Emerging Perspectives and Prospects. Isr. J. Chem. 2015, 55, 179−195. (46) Rocchigiani, L.; Ciancaleoni, G.; Zuccaccia, C.; Macchioni, A. Probing the Association of Frustrated Phosphine−Borane Lewis Pairs in Solution by NMR Spectroscopy. J. Am. Chem. Soc. 2014, 136, 112− 115. (47) Marwitz, A. J. V.; Dutton, J. L.; Mercier, L. G.; Piers, W. E. Dihydrogen Activation with tBu3P/B(C6F5)3: A Chemically Competent Indirect Mechanism via in Situ-Generated p-tBu2P− C6F4−B(C6F5)2. J. Am. Chem. Soc. 2011, 133, 10026−10029. (48) Karkamkar, A.; Parab, K.; Camaioni, D. M.; Neiner, D.; Cho, H.; Nielsen, T. K.; Autrey, T. A thermodynamic and kinetic study of the heterolytic activation of hydrogen by frustrated borane-amine Lewis pairs. Dalton Trans. 2013, 42, 615−619. (49) Whittemore, S. M.; Edvenson, G.; Camaioni, D. M.; Karkamkar, A.; Neiner, D.; Parab, K.; Autrey, T. Catalytic reduction of polar substrates without metals: A thermodynamic and kinetic study of heterolytic activation of hydrogen by vacancies in frustrated Lewis pairs. Catal. Today 2015, 251, 28−33. (50) Welch, G. C.; Juan, R. R. S.; Masuda, J. D.; Stephan, D. W. Reversible, Metal-Free Hydrogen Activation. Science 2006, 314, 1124− 1126. (51) Jiang, C.; Blacque, O.; Fox, T.; Berke, H. Reversible, metal-free hydrogen activation by frustrated Lewis pairs. Dalton Trans. 2011, 40, 1091−1097. (52) Greb, L.; Tussing, S.; Schirmer, B.; Ona-Burgos, P.; Kaupmees, K.; Lokov, M.; Leito, I.; Grimme, S.; Paradies, J. Electronic effects of triarylphosphines in metal-free hydrogen activation: a kinetic and computational study. Chem. Sci. 2013, 4, 2788−2796. (53) Whittemore, S. M.; Autrey, T. Kinetic and Thermodynamic Study of the Reduction of 1,1-Diphenylethylene by a Thermally Frustrated Diethyl Ether-BCF Lewis Pair. Isr. J. Chem. 2015, 55, 196− 201. (54) Hansen, C. W.; Hansen, L. D.; Nicholson, A. D.; Chilton, M. C.; Thomas, N.; Clark, J.; Hansen, J. C. Correction for instrument time constant and baseline in determination of reaction kinetics. Int. J. Chem. Kinet. 2011, 43, 53−61. (55) Zeonjuk, L. L.; Vankova, N.; Mavrandonakis, A.; Heine, T.; Röschenthaler, G.-V.; Eicher, J. On the Mechanism of Hydrogen Activation by Frustrated Lewis Pairs. Chem. - Eur. J. 2013, 19, 17413− 17424. (56) Xu, T.; Chen, E. Y. X. Probing Site Cooperativity of Frustrated Phosphine/Borane Lewis Pairs by a Polymerization Study. J. Am. Chem. Soc. 2014, 136, 1774−1777. (57) González-Hernández, R.; Chai, J.; Charles, R.; Pérez-Camacho, O.; Kniajanski, S.; Collins, S. Catalytic System for Homogeneous Ethylene Polymerization Based on Aluminohydride−Zirconocene Complexes. Organometallics 2006, 25, 5366−5373. (58) Liu, L.; Cao, L. L.; Shao, Y.; Ménard, G.; Stephan, D. W. A Radical Mechanism for Frustrated Lewis Pair Reactivity. Chem. 2017, 3, 259−267. (59) Shirono, K.; Morimatsu, T.; Takemura, F. Gas Solubilities (CO2, O2, Ar, N2, H2, and He) in Liquid Chlorinated Methanes. J. Chem. Eng. Data 2008, 53, 1867−1871. (60) Zhivonitko, V. V.; Sorochkina, K.; Chernichenko, K.; Kotai, B.; Foldes, T.; Papai, I.; Telkki, V.-V.; Repo, T.; Koptyug, I. Nuclear spin hyperpolarization with ansa-aminoboranes: a metal-free perspective for parahydrogen-induced polarization. Phys. Chem. Chem. Phys. 2016, 18, 27784−27795. (61) Lowry, T. H.; Richardson, K. S. Mechanism and Theory in Organic Chemistry, 2nd ed.; Harper & Row: New York, 1981. (62) Melander, L.; Saunders, W. H. Reaction Rates of Isotopic Molecules; Wiley-Interscience: New York, 1987.

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