D2 Exchange in Pt

Article pubs.acs.org/IC

Strain Control: Reversible H2 Activation and H2/D2 Exchange in Pt Complexes Rameswar Bhattacharjee, Abdulrahiman Nijamudheen, Sharmistha Karmakar, and Ayan Datta* Department of Spectroscopy, Indian Association for the Cultivation of Science, Jadavpur 700032, West Bengal, India S Supporting Information *

ABSTRACT: Experiments have indicated that bulky ligands are required for efficient H2 activation by Pt−Sn complexes. Herein, we unravel the mechanisms for a Pt−Sn complex, Pt(SntBu3)2(CNtBu)2 (1a), catalyzed reversible H2 activation. Among a number of Pt−Sn catalysts used to model H2 activation and H2/D2 exchange reactions, only 1a with large strain was found to be suitable because the addition of H2 to 1a requires lowest distortion energy, minimal structural changes, and smallest entropy of activation. The activity of this Pt−Sn complex was compared vis-à-vis its Pt−Ge and Pt−Si analogues, and we predicted that strained Pt−Ge complex can efficiently activate H2 reversibly. Direct dynamics calculations for the rate of reductive elimination of H2, HD, and D2 from Pt(SntBu3)(CNtBu)2H3 (4a) and Pt(SntBu3)(CNtBu)2HD2 (4a[2D]) shows that H/D atom tunneling contributes significantly, which leads to an enhanced kinetic isotope effect. Strain control is suggested as a design concept in H2 activation.

■

INTRODUCTION The activation of molecular hydrogen by transition metals is an intermediate step for a number of chemical transformations.1 Dihydrogen activation forms an integral part of many petrochemical industries,2 the synthesis of ammonia,3 and reversible hydrogen storage.4 Environmentally benign nonmetal catalysts such as frustrated Lewis pairs5 and (alkyl) (amino) carbenes6 are of considerable interest for their high efficiency in catalyzing hydrogenation reactions. As a synthetic strategy, many catalysts based on metal and main-group elements7 require the presence of bulky ligands to achieve the optimal efficiency.8 Recently, Captain and co-workers have demonstrated that a mononuclear Pt complex Pt(SntBu3)2(CNtBu)2 (1a, Scheme 1) can reversibly add H2 even at low temperature (at −72 °C) to form Pt(SntBu3)2(CNtBu)2(H)2, 2a (Scheme 1).9 X-ray crystal structure revealed that both 1a and 2a were substantially distorted from ideal planar and octahedral geometries, respectively, due to the presence of steric crowding by bulky SntBu3 and CNtBu. Comparatively sterically less demanding stannyl and isocyanide ligands (1d and 1e, Scheme 1, and for 1 with R = R′ = SnMes3) are detrimental for the cycle thereby indicating that bulky ligands are essential for activating H2/D2. The presence of bulky ligands makes the initial H2 addition to 1a more favorable process compared to competing Sn−H addition.10 For rational design of catalysts for H2 activation utilizing molecular strain, a microscopic understanding of mechanisms of catalysis in such reactions warrants further studies. In this Article, we calculated the energetics for competitive H−H and Sn−H additions to a number of experimental (1a, 1d, 1e) and modeled (1b, 1c, 1f, and 1g) Pt−Sn bimetallic complexes using density functional theory (DFT). Our results provide strong © XXXX American Chemical Society

support to the mechanisms for reversible H2 activation and H2/ D2 exchange proposed from experimental observations. We quantified individual components that contribute to the activation energies for H−H additions to 1a−1g by performing distortion interaction analyses.11 Our results decipher the origin of higher activity of 1a to the strain induced by substitutions on isocyanide and stannyl ligands. The most favorable reaction requires the least distortion energy as a penalty while transforming from reactant to the transition state. Interestingly, among all H2 addition pathways (from 1a−1g), the reaction from 1a has the lowest entropy of activation (ΔS‡). Additional support for our hypothesis was also provided by calculations on model complexes where both Sn atoms of 1a were replaced by either Ge or Si. The rate of H2/D2 exchange over a temperature range of 200−400 K, kinetic isotope effects (KIE), and the role of quantum mechanical tunneling (QMT) in the reductive elimination step were investigated through direct dynamic calculation that incorporates multidimensional tunneling. It is shown that at 300 and 200 K, 70% and 99%, respectively, of the H2 reductive elimination from a Pt−Sn complex proceeds via QMT that may be subjected to experimental verification.

■

COMPUTATIONAL METHODS

All geometry optimizations and vibrational frequency calculations were performed in gas phase at M06-2X12/6-31+G(d),13 LANL2DZ14 (Pt, Sn, Ge) level using Gaussian 09.15 This method was represented as M06-2X/BS1 for simplicity. To reduce the computational cost, few large systems (see Table 1) were optimized using multilayer ONIOM16 method by employing M06-2X/BS1 level for high layer and PM617 level for low layer followed by which DFT single points Received: December 17, 2015

A

DOI: 10.1021/acs.inorgchem.5b02904 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

Scheme 1. Pt−Sn Bimetallic Complexes and Possible Pathways Modeled to Study the Mechanisms of Reversible H2 Activation and H2/D2 Exchange Reactions

Table 1. Gibbs Free Energies (at 298 K) for the H−H and Sn−H Addition on 1a−1g Are Provided reaction

H−H addition

Sn−H addition

entry

ΔG‡ (kcal/mol)

ΔH‡ (kcal/mol)

ΔS‡ (e.u.)a

ΔGr (kcal/mol)

dH−Hb in TS (Å)

1a 1b 1c 1d 1e 1f 1g

15.7 26.8 21.2 27.4 24.5 21.9 26.6

15.1 22.1 19.8 22.3 20.7 20.3 22.0

−2.2 −15.9 −4.5 −17.2 −12.7 −5.4 −15.4

0.5 11.4 7.6 12.7 11.5 9.3 13.5

0.96 1.01 0.99 0.98 1.03 1.02 1.03

ΔG‡ (kcal/mol)

ΔS‡ (e.u.)a

14.2

−52.3

10.2c

−60.6c

15.8 15.6

−58.5 −44.8

ΔGr (kcal/mol) 97.8c 6.1 89.7c 4.7c 45.8c 6.9 10.3

a

e.u. is the entropy unit. bdH−H is the H−H bond distance in the H−H addition TS. cRepresents that the geometry was optimized using QM/MM, ONIOM model followed by single point calculations at M06/BS2, level of theory.31

Figure 1. Computed Gibbs free energy profile for the reversible activation of H2 by 1a. Relative free energies (ΔGr and ΔG‡ in kcal/mol, at 298 K) and selected bond distance (in Å) are shown. energies were obtained. The specifications of the high and low layers are shown in the Supporting Information file. All the stationary points were confirmed by the absence of any imaginary frequency and transition states by single imaginary frequency. Intrinsic reaction coordinate calculations were performed to ensure that each transition state does connect correct reactants and products. We selected M062X functional over other recommended methods12 such as M06 or M06-L to perform geometry optimizations because some of the structures were not converged while using these latter methods. Nevertheless, reported thermochemical energies (at 298 K, 1 atm) were obtained from single-point calculations on these structures at M0618/TZVP,19 LANL2TZ(f)20 (Pt), LANL08(d)20 (Sn, Ge) level of

theory, which we represent as M06/BS2. Therefore, the error in energies due to the optimization at M06-2X level could be neglected, which was also verified by comparing the results from M06/BS2// M06/BS1 calculations for a few selected structures. We further performed single-point calculation on the full catalytic cycle shown in Figure 1 using other hybrid DFT methods such as M06L,21 ωB97XD,22 and B3LYP-D323,24 with the same basis set (BS1) to examine the reliability of the results. Solvent effects in the entropy of activation (ΔS‡) were explored by an implicit PCM solvation model (SMD)25 for n-hexane. The reaction rates were obtained based on canonical variation transition-state theory (CVT or CVTST)26 using GAUSSRATE27 as B


Article

Inorganic Chemistry an interface between POLYRATE28 and Gaussian 09. The effects of QMT on the reaction kinetics were taken into account through small curvature tunneling (SCT) approximation.29 We generated the VaG (vibrational adiabatic ground-state potential energy, which is the sum of the electronic energy and zero-point energy) surface by performing electronic structure calculation at each point along the minimum energy path using a very small and accurate step size (Δs = 0.02 Å amu−1/2). Then on the VaG surface, the reaction rate was calculated using the canonical variation transition-state theory, which maximizes the free energy of activation (thereby minimizing the reactive flux) along the reaction path and also takes into account the recrossing effect. Small curvature tunneling (SCT) is considered as a suitable model for describing quantum effect on the reaction dynamics for typical transition-metal based systems. These computationally expensive calculations were performed using M06/BS1 level of theory, which is in good agreement with M06/BS2 for this reductive elimination step.30 Optimized geometry of 1a is substantially distorted from square planar to sawhorselike structure and shows good agreement with its crystal structure, hence supporting the suitability of the theoretical methods utilized (see Supporting Information file for a detailed comparison).

reversible oxidative addition of H2 to 3a at the vacant site cis to SntBu3 would form the hexacoordinated octahedral complex 4a. This highly endergonic step (ΔGr = 12.2 kcal/mol) is identified as the process with the highest barrier (TS3, ΔG‡ = 24.0 kcal/ mol). Addition of H2 to 3a in an alternate fashion, at the vacant site trans to SntBu3, was discarded due to the requirement of a higher ΔG‡ (ΔG‡ = 36.9). The rate-determining step that involves the addition of H2 to RC2 would promote H2/D2 exchange in the presence of a gaseous mixture of H2 and D2. However, the reversible H2 activation is controlled mainly by the initial oxidative addition of H−H to RC1, and hence, a faster rate for competing SnR3− H addition to 1a (see path b, Scheme 1) can inhibit the entire cycle in Figure 1. We found that the addition of SntBu3−H to 1a is highly energy demanding with a calculated reaction free energy (ΔGr) of +97.8 kcal/mol. Such strong thermodynamic instability for the product for Sn(tBu)3−H addition arises primarily due to the steric crowding of tBu3 groups. Nevertheless, this result would readily exclude any possible role of SntBu3−H addition in the mechanism. In general, the catalytic cycle for the H2 activation and H2/D2 exchange starting from 1 with different substituents could be represented as shown in Scheme 2.

■

RESULTS AND DISCUSSION Calculated Gibbs free energy profile for H2 activation by Pt−Sn complex 1a is shown in Figure 1. The envisaged mechanism would be initiated by the formation of a reactant complex (RC1) through the weak association of a H2 molecule with 1a (ΔGr = 5.2 kcal/mol). In the successive step, RC1 would be converted to the hexacoordinated complex 2a via crossing of the barrier TS1 (see Figure 2 for the optimized geometries of

Scheme 2. Catalytic Cycle Depicting the H2 Activation and H2/D2 Exchange by 1

We compared thermodynamic and kinetic preferences for competing pathways of H−H and Sn−H additions to a number of Pt−Sn complexes (1a−1g) consisting of different substitutions on stannyl and isocyanide ligands. As shown in Figure 1, the H−H addition to 1a to form 2a is a kinetically favorable process due to the requirements of a small ΔG‡ (15.7 kcal/ mol), and the reaction is readily reversible because ΔGr = 0.5 kcal/mol. These results are in very good agreement with the experimental observations that the conversion from 1a to 2a is reversible in nature and can occur even at low temperature. Comparing these results with other systems (1b−1g) shows that the presence of bulky tBu groups on both stannyl and isocyanide ligands (1a) is an essential condition for the kinetically favorable H−H addition. The complexes with tBu substitution on either the isocyanide (1b, Table 1) or stannyl ligand (1c, 1d, and 1e, Table 1) were found to be insufficient to

Figure 2. Optimized geometries for the transition states involved in the reaction coordinate for the H2 activation by Pt−Sn complex (1a). Selected bond distances (in Å) and ΔG‡ values (in kcal/mol) are provided. Only relevant H atoms are shown for clarity.

TSs involved in the mechanism). This oxidative addition process is slightly endergonic (ΔGr = 0.5 kcal/mol) and requires an activation barrier (ΔG‡) of 15.7 kcal/mol. A facile Sn−H reductive elimination from 2a forming product complex PC1 (ΔGr = 3.1 kcal/mol, ΔG‡ = 12.1 kcal/mol) and subsequent loss of Sn−H ligand from PC1 would furnish the four-coordinated intermediate 3a. Unlike 1a and 2a, optimized geometry of 3a is almost planar (dihedral Φ(Sn−C−C−H) = 0.0°), which could be attributed to the release of strain in 3a upon the reductive elimination of Sn−H from PC1. Finally, a C


Article

Inorganic Chemistry

regioselectivity of Au/Ga catalyzed Nakamura reactions.36 In the present case, our results in Figure 3 elucidate that the

promote the H−H addition. For the system in which R = Ph and R′ = tBu (1d), the largest activation free energy (27.4 kcal/ mol) was obtained. Incidentally, all four coordinated complexes other than 1a were found to be stable in nearly planar geometries (see Supporting Information file). The transitionstate geometry corresponding to H2 addition to 1a has a reactant-like structure with small H−H (0.96 Å) and long Pt− H (1.74 Å) bonds (see Figure 2). A comparison of all TS geometries indicates that the activation free energy increases sharply for systems in which the H−H bond is almost cleaved and strong Pt−H bonds already formed. A reactant-like TS is expected to be formed via minimal changes in enthalpy as well as entropy.32 To gain direct evidence for this concept, we calculated the enthalpy and entropy of activation for all systems. Intriguingly, H2 addition to 1a has the lowest entropy of activation (ΔS‡ = −2.2 e.u.) compared to any other system. Similar trend is also reflected in the calculated enthalpy of activation (ΔH‡) for 1a that shows similar value as ΔG‡ (ΔH‡ = 15.1 and ΔG‡ = 15.7 kcal/mol) compared to 1d (ΔH‡ = 22.3 and ΔG‡ = 27.4 kcal/mol) and 1b (ΔH‡ = 22.1 and ΔG‡ = 26.8 kcal/mol, respectively). Hence, a lower ΔH‡ as well as lower ΔS‡ value contributes to a low free-energy barrier for H2 addition to 1a. As seen from Table 1, the enthalpy of activation is decisive in controlling the overall free energy of activation for 1a, 1c, and 1f, while the entropy of activation is important in 1b, 1d, 1e, and 1g. We had recently shown that unusually fast reductive eliminations of C−C bonds from a number of Au(III) complexes proceed via concerted pathways those involving very low entropy contribution since their transition states have minimal structural changes with respect to the reactant.33 While accurate calculation of ΔS‡ requires high-quality calculations incorporating explicit solvent molecules, gas-phase values reported do provide useful qualitative understanding. We also calculated the ΔS‡ using implicit solvation model and obtained similar trends reported here (see Supporting Information Table S7 for details). The reliability of ΔS‡ obtained from both gas phase and implicit solvent model DFT calculations for rationalizing organic reactions has been reported in literature.34 Opposite to the trends obtained for H−H addition, Sn−H addition was found to be highly unfavorable for sterically crowded systems 1a, 1c, and 1e and relatively favorable for smaller systems 1f and 1g. However, Sn−H addition to the relatively bulky system 1d proceeds through a low activation barrier (10.2 kcal/mol) due to the presence of a number of attractive CH···π interactions, which were visualized using the NCI plot program 35 (see Figure S16, of Supporting Information file). The above calculations clearly show that only the highly strained complex 1a can efficiently activate H−H bond. One fruitful approach to quantify the strain induced by ligands is to calculate the distortion energy along the reaction coordinate. In this regard, activation strain analysis and closely related distortion interaction models have been successful for a number of reactions.11 In the distortion interaction method, the activation energy along the reaction coordinate (ΔEact(ζ)) is separated into the contributions due to distortion and interaction energies ΔEdist(ζ) and ΔEint(ζ), respectively, as ΔEact(ζ) = ΔEdist(ζ) + ΔEint(ζ). Here, ΔEdist(ζ) is the sum of energies required for individual reactants to distort from their reactant geometries to the transition-state geometries, whereas ΔEint(ζ) is the interaction energy between individual reactants in the transition-state geometry. Recently, we applied distortion interaction analysis to unravel the role of ligands in tuning

Figure 3. Distortion interaction analyses along the reaction coordinates (ζ) for H2 addition to Pt−Sn complexes (1a−1g) are shown. Distortion energy (ΔEdist(ζ), black line), activation energy (ΔEact(ζ), blue line), and interaction energy (ΔEint(ζ), red line) are plotted as a function of the change in H−H bond distance (dH−H(ζ)).

difference in ΔEact(ζ) arises primarily due to the contribution from ΔEdist(ζ). Therefore, while all H2 addition TSs have similar ΔEint(ζ) along the reaction coordinate, ΔEdist(ζ) values differ significantly. This difference in ΔEdist(ζ) is directly reflected in their ΔEact(ζ). For example, ΔEdist(ζ) calculated at TS geometry for 1a and 1g are of 37.4 and 51.9 kcal/mol, respectively, while their ΔEint(ζ) are −23.0 and −30.1 kcal/ mol, respectively, and their ΔEact(ζ) are 14.4 and 21.8 kcal/ mol, respectively. Therefore, Pt complexes with planar or nearly planar systems must distort substantially to reach the TS geometry, leading to higher ΔEdist(ζ) and ΔS‡. Hence, such reactions are kinetically not favored. Further calculations were performed using modeled germyl and silyl analogues of 1a, namely, 6a and 10a, respectively, to unravel the role of cooperativity between Sn and Pt in the reaction. Figure 4 depicts the Gibbs free energy profile for H2 activation by Pt−Ge (6a) and Pt−Si (10a). It has been found that activation free energies for H2 addition to 6a is 2.5 kcal/ mol lower in energy than to 1a (ΔG‡(1a) = 15.7 kcal/mol and ΔG‡(6a) = 13.9 kcal/mol). Additionally, all intermediate steps involved are reversible, unveiling its potential for H2 activation and H2/D2 exchange reactions. The comparable reactivities predicted for 1a and 6a could be due to the combined effects from their similar distortion energy and entropy of activation values (ΔEdist(ζ) = 34.4 kcal/mol at TS geometry and ΔS‡ = −3.2 e.u. for 6a). However, for the reaction using Pt−Si complex, all steps involved require higher ΔG‡, and reversibility for individual steps is reduced (see Figure 5). However, The enthalpy of activation of hydrogen addition to 6a and 10a are comparable with 1a (ΔH‡(1a) = 15.1 kcal/mol; ΔH‡(6a) = 12.9 kcal/mol, and ΔH‡(10a) = 13.4 kcal/mol). Even though optimized geometry of Pt−Si complex 10a was found to be highly strained (ΔEdist(ζ) = 33.3 kcal/mol at TS geometry), H2 addition requires large negative ΔS‡ (−22.5 e.u). This large negative value of entropy of activation for Pt−Si complexes (10a) may be attributed to the fact that 10a is highly crowded at Pt (because the Pt−Si bond length is short; Pt−Si = 2.51 Å, D


Article

Inorganic Chemistry

Figure 4. Computed Gibbs free energy (G) profile for the reversible activation of H2 by Pt−Ge complex (6a). Relative free energies (ΔGr and ΔG‡ in kcal/mol) and selected bond distance (in Å) are shown.

Figure 5. Computed Gibbs free energy (G) profile for the reversible activation of H2 by Pt−Si (10a). Relative free energies (ΔGr and ΔG‡ in kcal/ mol) and selected bond distance (in Å) are shown.

Figure 6. (A) Calculated reactions rates (in s−1) for the reductive elimination of D1D2 and HD1 from 4a′[2D] at room temperature to form 3a′[H] and 3a′[D2], respectively. (B) Gibbs free energy (ΔGr and ΔG‡ in kcal/mol) profiles for alternative routes of H2 elimination from 4a′. (C) Arrhenius plots for HD1 and D1D2 reductive elimination from 4a′[2D] in the temperature range of 200−400 K.

and hence increase its ΔS‡. Therefore, while the strain energy decreases from Pt−Sn to Pt−Si, the change in ΔS‡ become a crucial factor in controlling the reaction. In the intermediate

Pt−Ge = 2.59 Å, and Pt−Sn = 2.69 Å) and has less space to accommodate H2 at the metal center. This forces the system to undergo significant rearrangements to reach the TS geometry E


Article

Inorganic Chemistry steps, the barriers for the elimination of HEtBu3 (E = Sn, Ge, and Si) from Pt(EtBu3)2(CNtBu)2H2 was found to follow the order Si < Ge < Sn, whereas the addition of HEtBu3 (E = Sn, Ge, and Si) to Pt(EtBu3)2(CNtBu)2 follow exactly opposite trends. These results can be explained from the orders of bond strengths Pt−E and H−E bonds; Si−H > Ge−H > Sn−H and Pt−Sn > Pt−Ge > Pt−Si.37 Notably, intermediates 8a and 12a are more stable than 6a and 10a, respectively, while for the reaction of Pt−Sn system, 3a is less stable than 1a, which indicates more strain being released during the formation of 8a and 12a. Further H2 addition to 8a and 12a requires relatively more energy compared to 3a (ΔG‡(8a) = 26.9 kcal/mol and ΔG‡(12a) = 29.1 kcal/mol). Hence, the H2 addition to platinum complexes is shown to depend on steric effects and nature of the ligand. As reported,10 in the presence of a gaseous mixture of H2 and D2, 3a (see Figure 1) can activate H2−D2 exchange and hence produce HD. By means of direct dynamics calculations, we calculated the kinetics for this process for a temperature range of 200−400 K. It was found that, unlike the oxidative addition or reductive elimination on or from 2a, the activation barrier for these processes for the dyads 3a and 4a do not depend on the substitution on isocyanide and stannyl ligands. Replacement of SntBu3 and CNtBu by SnH3 and CNH, respectively, led to negligible change in reaction energetics.38 Therefore, we used small model systems 3a′ (3a′[H], 3a′[H1], 3a′[H2], and 3a′[D2]), 4a′, and 4a′[2D] for 3a, 4a, and their deuterated structures, respectively (see Figure 6), because direct dynamics calculations on the actual systems 3a and 4a are computationally unrealistic. The hexacoordinated complex 4a′ can undergo reductive elimination via three unique pathways leading to the elimination of H−H 1 , H−H 2 , and H 1 −H 2 . Reductive elimination of H−H1 and H1−H2 from 4a′ were found to be iso-energetic (ΔG‡ = 10.1 kcal/mol, Figure 6B) and more favorable compared to the reductive elimination of H−H2 (ΔG‡ = 12.0 kcal/mol). Arrhenius plots calculated for the reductive elimination of HD1 (or HD) and D1D2 (or DD) from 4a′[2D] with (CVT + SCT) and without (CVT) inclusion of tunneling for a temperature range of 200−400 K are shown in Figure 6C. At room temperature (298 K), classical reaction rates for both HD and DD eliminations were predicted to be similar; kCVT(HD1) = 1.73 × 105 s−1, and kCVT(D1D2) = 1.15 × 105 s−1. However, the computed CVT + SCT rates were significantly faster for HD elimination; kCVT+SCT(HD1) = 5.18 × 105 s−1, and kCVT+SCT(D1D2) = 2.48 × 105 s−1. Similar calculations performed for H2 (H−H1) elimination from 4a′ lad to CVT + SCT rate of (kCVT+SCT(HH1)) of 1.96 × 106. Hence, inclusion of tunneling enhances the rate of HH and HD elimination compared to DD by 7.9 and 2.2 times, respectively, at 298 K. Clearly, the enhanced KIE is not captured by zeropoint energy effects only, and one needs to explicitly consider QMT. At lower temperatures, the contributions from QMT increase substantially. At 200 K, ∼99% of the reaction occurs via thermally assisted tunneling. The representative tunneling energy is found to be 0.4 kcal/mol lower than the top of barrier (maxima of the vibrationally adiabatic ground-state potential energy, VaG(s = 0)). Therefore, at 200 K, kCVT+SCT(HH1) = 2.99 × 10 3 s −1 , k CVT+SCT (HD 1 ) = 4.69 × 10 2 s −1 , and kCVT+SCT(D1D2) = 6.36 × 101 s−1. Hence, QMT increases the rates for the eliminations of HH and HD over DD by a factor of 47.0 and 7.4, respectively. The curvature in the CVT + SCT Arrhenius plot is evident even for an ambient temperature of ∼250 K for HD elimination. However, for DD elimination a

curvature can be observed only below 200 K. The importance of QMT in the preference of HD elimination over DD may be experimentally verified.

■

CONCLUSIONS In conclusion, we have provided an understanding of the mechanisms of reversible H2 activation by a highly strained Pt− Sn complex, Pt(SntBu3)2(CNtBu)2 (1a). All Pt−Sn systems other than 1a (1b−1g) were found to be inefficient for H2 activation due to the requirement of higher distortion energies. Minimal structural rearrangements required for 1a for H2 addition leads to small entropy of activation obtained for this reaction. It has been predicted that the Pt−Ge analogue of 1a can also exist as a highly strained complex, which can catalyze the reversible H2 activation and H2/D2 exchange reactions. Calculations reveal nontrivial contribution of QMT for the reductive elimination of H2 and HD at experimentally relevant temperatures. Contrary to the general perception that bulky ligands hamper oxidative addition on transition metal, the present study depicts strain induced by bulky ligands as a positive factor toward small-molecule activation.

■

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.5b02904. Complete reference for Gaussian 09, crystal structure comparison with DFT optimized geometry, spatial orientation of the transition states, ONIOM specification, suitability of ONIOM method, NCI plot, distortion/ interaction data, entropy of activation in n-hexane, CVT and CVT+SCT rate constants, Cartesian coordinates, electronic energies, zero point energy correction to electronic energy (EZPE), zero point energy correction to Gibbs free energies (GZPE), and zero point energy corrected enthalpies (HZPE). (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■ ■

ACKNOWLEDGMENTS R.B. and S.K. thank CSIR India for SRF. A.D. thanks INSA, DST, and BRNS for partial funding. REFERENCES

(1) (a) Vaska, L. Acc. Chem. Res. 1968, 1, 335. (b) Saillard, J.; Hoffmann, R. J. Am. Chem. Soc. 1984, 106, 2006. (c) Cui, Q.; Musaev, D. G.; Morokuma, K. J. Chem. Phys. 1998, 108, 8418. (2) (a) Burch, R.; Garla, L. C. J. Catal. 1981, 71, 360. (b) Fujikawa, T.; Ribeiro, F. H.; Somorjai, G. A. J. Catal. 1998, 178, 58. (3) (a) Pool, J. A.; Lobkovsky, E.; Chirik, P. J. Nature 2004, 427, 527. (b) Kubas, G. J. Science 2006, 314, 1096. (4) (a) Hull, J. F.; Himeda, Y.; Wang, W.-H.; Hashiguchi, B.; Periana, R.; Szalda, D. J.; Muckerman, J. T.; Fujita, E. Nat. Chem. 2012, 4, 383. (c) Jena, P. J. Phys. Chem. Lett. 2011, 2, 206. (5) (a) Grimme, S.; Kruse, H.; Goerigk, L.; Erker, G. Angew. Chem., Int. Ed. 2010, 49, 1402. (b) Kronig, S.; Theuergarten, E.; Holschumacher, D.; Bannenberg, T.; Daniliuc, C. G.; Jones, P. G.; Tamm, M. Inorg. Chem. 2011, 50, 7344. (c) Fan, C.; Mercier, L. G.; F


Article

Inorganic Chemistry Piers, W. E.; Tuononen, H. M.; Parvez, M. J. Am. Chem. Soc. 2010, 132, 9604. (6) (a) Frey, G. D.; Lavallo, V.; Donnadieu, B.; Schoeller, W. W.; Bertrand, G. Science 2007, 316, 439. (b) Henkel, S.; Sander, W. Angew. Chem., Int. Ed. 2015, 54, 4603. (7) (a) Power, P. P. Nature 2010, 463, 171. (b) Li, B.; Xu, Z. J. Am. Chem. Soc. 2009, 131, 16380. (8) (a) Gunanathan, C.; Milstein, D. Acc. Chem. Res. 2011, 44, 588. (b) Hauger, B. E.; Gusev, D.; Caulton, K. G. J. Am. Chem. Soc. 1994, 116, 208. (9) Yempally, V.; Zhu, L.; Captain, B. Inorg. Chem. 2010, 49, 7238. (10) Koppaka, A.; Zhu, L.; Yempally, V.; Isrow, D.; Pellechia, P. J.; Captain, B. J. Am. Chem. Soc. 2015, 137, 445. (11) (a) Kitaura, K.; Morokuma, K. Int. J. Quantum Chem. 1976, 10, 325. (b) Nagase, S.; Morokuma, K. J. Am. Chem. Soc. 1978, 100, 1666. (c) Bickelhaupt, F. M. J. Comput. Chem. 1999, 20, 114. (d) Diefenbach, A.; de Jong, G. T.; Bickelhaupt, F. M. Mol. Phys. 2005, 103, 995. (e) van Zeist, W.-J.; Bickelhaupt, F. M. Org. Biomol. Chem. 2010, 8, 3118. (f) Schoenebeck, F.; Houk, K. N. J. Am. Chem. Soc. 2010, 132, 2496. (g) Green, A. G.; Liu, P.; Merlic, C. A.; Houk, K. N. J. Am. Chem. Soc. 2014, 136, 4575. (12) (a) Zhao, Y.; Truhlar, D. G. J. Chem. Theory Comput. 2008, 4, 1849. (b) Zhao, Y.; Truhlar, D. J. Theor. Chem. Acc. 2008, 120, 215. (c) Nijamudheen, A.; Jose, D.; Shine, A.; Datta, A. J. Phys. Chem. Lett. 2012, 3, 1493. (d) Sun, Y. Y.; Chen, H. J. Chem. Theory Comput. 2013, 9, 4735. (13) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213− 222. (14) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 299. (15) Frisch, M. J.; et al. Gaussian 09; Gaussian, Inc: Pittsburgh, PA, 2009. Full citations are available in the Supporting Information. (16) (a) Dapprich, S.; Komáromi, I.; Byun, K. S.; Morokuma, K.; Frisch, M. J. J. Mol. Struct.: THEOCHEM 1999, 461-462, 1. (b) Svensson, M.; Humbel, S.; Froese, R. D. J.; Matsubara, T.; Sieber, S.; Morokuma, K. J. Phys. Chem. 1996, 100, 19357. (17) Stewart, J. J. P. J. Mol. Model. 2007, 13 (12), 1173. (18) (a) Faza, O. N.; Rodríguez, R. Á .; López, S. C. Theor. Chem. Acc. 2011, 128, 647. (b) Laury, M. L.; Wilson, A. K. J. Chem. Theory Comput. 2013, 9, 3939. (c) Zhao, Y.; Truhlar, D. Theor. Chem. Acc. 2008, 120, 215. (19) Dunning, T. H., Jr. J. Chem. Phys. 1971, 55, 716. (20) Wadt, W. R.; Hay, P. J. J. Chem. Phys. 1985, 82, 284. (21) Zhao, Y.; Truhlar, D. G. J. Chem. Phys. 2006, 125, 194101. (22) Chai, J. D.; Head-Gordon, M. J. Chem. Phys. 2009, 131, 174105. (23) Becke, A. D. Phys. Rev. A: At., Mol., Opt. Phys. 1988, 38, 3098. (b) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev. B: Condens. Matter Mater. Phys. 1988, 37, 785. (24) Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. J. Chem. Phys. 2010, 132, 154104. (25) Marenich, A. V.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. B 2009, 113, 6378. (26) Truhlar, D. G.; Garrett, B. C. Annu. Rev. Phys. Chem. 1984, 35, 159. (27) Zheng, J.; Zhang, S.; Corchado, J. C.; Chuang, Y. Y.; Coitino, E. L.; Ellingson, B. A.; Truhlar, D. G. GAUSSRATE, version 2010; University of Minnesota: Minneapolis, MN, 2010. (28) Zheng, J.; POLYRATE, version 2010; University of Minnesota: Minneapolis, MN, 2010. (29) (a) Hu, W.-P.; Liu, Y.-P.; Truhlar, D. G. J. Chem. Soc., Faraday Trans. 1994, 90, 1715. (b) Fernandez-Ramos, A.; Ellingson, B. A.; Garrett, B. C.; Truhlar, D. G. Variational Transition State Theory with Multidimensional Tunneling. In Reviews in Computational Chemistry; Lipkowitz, K. B., Cundari, T. R., Eds.; Wiley-VCH: Hoboken, NJ, 2007; Vol. 23, pp 125−132. (30) The free energy barrier for HH1 and H1H2 (Figure 6B) elimination from 4a′ is computed to be 10.1 kcal/mol with in M06/ BS1 level of theory, which is in good agreement with the M06/BS2 method, where the barrier height found to be 10.3 kcal/mol.

(31) Comparative studies for H2 addition on 1d using both ONIOM and DFT-optimized geometries give ΔG‡ to be 27.1 and 27.4 kcal/ mol, respectively, and ΔE‡ to be 23.6 and 23.7 kcal/mol, respectively. See also Figure S15 (Supporting Information) for a comparison of geometries from both methods. (32) Hauger, B. E.; Gusev, D.; Caulton, K. G. J. Am. Chem. Soc. 1994, 116, 208. (33) Nijamudheen, A.; Karmakar, S.; Datta, A. Chem. - Eur. J. 2014, 20, 14650. (34) (a) Dogan, B.; Catak, S.; Van Speybroeck, V.; Waroquier, M.; Aviyente, V. Polymer 2012, 53, 3211. (b) Ruff, F.; Farkas, Ö . J. Org. Chem. 2006, 71, 3409. (c) Feliz, M.; Freixa, Z.; van Leeuwen, P. W. N. M.; Bo, C. Organometallics 2005, 24, 5718. (d) Domingo, L. R.; Aurell, M. J.; Arnó, M.; Sáez, J. A. J. Org. Chem. 2007, 72, 4220. (35) (a) Contreras-Garcia, J.; Johnson, E. R.; Keinan, S.; Chaudret, R.; Piquemal, J.-P.; Beratan, D. N.; Yang, W. J. Chem. Theory Comput. 2011, 7, 625. (b) Johnson, E. R.; Keinan, S.; Mori-Sanchez, P.; Contreras-Garcia, J.; Cohen, A. J.; Yang, W. J. Am. Chem. Soc. 2010, 132, 6498. (36) Bhattacharjee, R.; Nijamudheen, A.; Datta, A. Org. Biomol. Chem. 2015, 13, 7412. (37) (a) Ozawa, F.; Hikida, T.; Hasebe, K.; Mori, T. Organometallics 1998, 17, 1018. (b) Ozawa, F. J. J. Organomet. Chem. 2000, 611, 332. (c) Sakaki, S.; Ieki, M. J. Am. Chem. Soc. 1993, 115, 2373. (d) Basch, H. Inorg. Chim. Acta 1996, 252, 265. (e) Liao, W. H.; Ho, P. Y.; Su, M. D. Inorg. Chem. 2013, 52, 1338. (38) ΔG‡ for H2 (H1H2 and H1H3) elimination from 4a′, where tBu has been replaced by H, is 10.1 kcal/mol, and for the experimental system (4a) ΔG‡ = 11.8 kcal/mol. So, the energy barriers for both the modeled and experimental systems are comparable.

G


D2 Exchange in Pt

Recommend Documents