Nickel Pincer Complexes with Frequent Aliphatic Alkoxo Ligands

Article pubs.acs.org/IC

Nickel Pincer Complexes with Frequent Aliphatic Alkoxo Ligands [(iPrPCP)Ni-OR] (R = Et, nBu, iPr, 2‑hydroxyethyl). An Assessment of the Hydrolytic Stability of Nickel and Palladium Alkoxides Luis M. Martínez-Prieto, Pilar Palma, Eleuterio Á lvarez, and Juan Cámpora* Instituto de Investigaciones Químicas. CSIC-Universidad de Sevilla. c/Américo Vespucio, 49. 41092, Sevilla, Spain S Supporting Information *

ABSTRACT: A series of nickel pincer complexes with terminal alkoxo ligands [(iPrPCP)Ni-OR] (R = Et, nBu, iPr, CH2CH2OH; iPrPCP is the 2,6-bis(diisopropylphosphinomethyl)phenyl pincer ligand) was synthesized and fully characterized. Together with the previously reported methoxo analogues of Ni and Pd, these complexes constitute a unique series of isostructural late transition-metal alkoxides. Spectroscopic and Xray diffraction data provide direct indications of the strong polarization of their covalent Ni-OR bonds. One of the most salient features of this class of compounds is their facile hydrolysis with traces of moisture, leading to equilibrium mixtures with the corresponding hydroxides [(iPrPCP)M− OH] (M = Ni or Pd) and alcohols, ROH. To compare the hydrolytic stability of nickel and palladium alkoxides, we performed NMR titrations of both hydroxides with several alcohols and determined the corresponding equilibrium constants. In general, these constants are ca. 1 order of magnitude smaller for M = Ni than Pd, indicating that Ni alkoxide complexes are more readily hydrolyzed than their Pd counterparts. For alkoxide complexes containing heteroatom-free R groups, the tendency to hydrolyze decreases as the parent alcohol ROH becomes more acidic, that is, R = Me > Et > iPr. This intuitive trend is broken for 2-methoxyethanol, the most acidic alcohol investigated. The hydroxo/2-methoxyethanol exchange equilibrium constants are comparable to those of ethanol (M = Ni) or methanol (M = Pd), showing that the corresponding 2-methoxyethoxide complexes are more prone to hydrolysis than anticipated. These experimental observations were rationalized in the light of density functional theory calculations.

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INTRODUCTION Well-defined late transition-metal complexes containing nondative σ metal−heteroatom bonded ligands, such as hydroxo, alkoxo, amido, or fluoro, compose a fascinating class of compounds.1 While these are commonly found as inert or spectator ligands in complexes with metals from groups 4−6, they become less common and more reactive in those of the second half of the transition series. Hydroxide and alkoxide complexes of Pd and other late transition metals have long been known to participate as key intermediates in important catalytic processes, such as the Wacker reaction,2 as well as in relevant biological processes.3 The origin of the distinct reactivity of this class of late transition-metal complexes has been largely debated. In general, the electronegativity difference between metals and heteroatoms E = N, O, or F leads to strongly polarized Mδ+−Eδ− bonds. In early transition-metal derivatives, the electron density excess on the heteroatom is alleviated by π donation from the nonbonding electron pairs of the heteroatom into empty metal d orbitals. The π interaction not only decreases the reactivity of E but contributes to the strength of the M−E bond. For late transition metals, there are fewer empty d orbitals, and these may not have suitable symmetry to accept π-donation from covalently bound heteroatoms. In consequence, late transition-metal hydroxide © XXXX American Chemical Society

or alkoxide complexes resemble those of electropositive maingroup elements or post-transition metals in their strongly basic and nucleophilic character.4 Moreover, it has been argued that the basicity of the heteroatom is further enhanced by nonbonding interactions of its lone electron pairs with the filled d orbitals of the metal.5,6 Although the general consensus is that the effect of repulsive p−d interactions is small,7 there are other factors that admittedly contribute to boost the reactivity of late transition-metal amides, alkoxides, fluorides, etc., such as the soft/hard mismatch between the metal and the heteroatom1a or kinetic effects due to the coexistence of contiguous Lewis-acidic metal and a basic heteroatom ligand.1f Over the last two decades, growing evidence has built showing that migratory insertion and other typical reactions of σorganometallic compounds is not limited to M−C bonds but can also involve reactive bonds between late transition metals and N, O, or even F atoms.8,9 The combination of basic/ nucleophilic and “organometallic-like” reactivity of late transition metal-heteroatom bonds has guided the design of new catalysts for many useful transformations, such as transfer hydrogenation,10 hydrogen-borrowing amination of alcohols,11 Received: July 26, 2017

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DOI: 10.1021/acs.inorgchem.7b01868 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry alcohol dehydrogenation12 and oxidation,13 or carbon-heteroatom cross-coupling reactions,14 that have become essential parts in the contemporary toolbox of organic syntheses. The high reactivity of late transition-metal complexes with σheteroatom bound ligands, although essential for their role as catalysts, can be also detrimental in terms of catalyst lifetime and selectivity. It also poses many practical challenges for fundamental reactivity studies. Late transition-metal alkoxides can be thermally unstable15 and are in general very sensitive to hydrolysis by residual moisture. Thus, despite their importance in catalysis, few late transition-metal alkoxides other than methoxides have been fully characterized. For group 10 metals, most of them are derived from considerably acidic fluoroalcohols16 or have metallacyclic structures.17 The Cambridge Structural Database contains no crystal structures of Ni, Pd, or Pt monomeric complexes containing such common alkoxo ligands as ethoxide, n-butoxide, or isopropoxide. In recent years, strongly chelating pincer ligands have been widely used to stabilize many types of otherwise difficult to handle metal species, decisively contributing to the advancement of homogeneous catalysis.18 Nickel19 and platinum20 tbutoxide complexes containing different tridentate pincer-type ligands have been reported, although the M−OR linkages in these complexes are stabilized by intramolecular hydrogen bonding or interactions with an alkaline cation. Hitherto, the only structurally authenticated example of a monomeric Pd complex bearing a primary alkoxo ligand contains a PCP scaffold.21 Our group has contributed to the exploration of pincer ligands for the stabilization of nickel and palladium complexes containing reactive σ-metal-heteroatom bonds. We have shown that the 2,5-bis(isopropylphosphino)methylphenyl ligand (iPrPCP) is particularly well-suited for this purpose and reported some of the first examples of pincer-type complexes of nickel and palladium containing terminal hydroxo,22 amido,23 and fluoro6b ligands. Very recently, we investigated β-hydrogen elimination in methoxides of the type [M(iPrPCP)OMe] (M = Ni, Pd), which initially leads to the corresponding hydrides [(iPrPCP)M-H] and, subsequently, to M(0) species.9b While the Pd-OMe complex decomposes readily in alcohol solvents, its nickel analogue is considerably more stable. In view of their potential utility in catalytic processes, we decided to investigate iPr PCP-stabilized Ni and Pd complexes containing higher alkoxo ligands arising from common alcohols. We are particularly interested in assessing the stability and hydrolytic sensitivity of this type of compound, as this is an important factor for catalytic and biological processes in which water is a reaction product, like alcohol carboxylation24 or the oxidation of alcohols in the liver by the alcohol dehydrogenase enzyme.25 Herein we report the synthesis and full characterization of a series of isostructural nickel alkoxides (ethoxide, n-butoxide, ipropoxide, and glycoxide). To assess the hydrolytic stability of nickel alkoxides and compare with those of their palladium analogues, we measured the equilibrium constants from the exchange reactions of hydroxide complexes and different alcohols and used density functional theory (DFT) calculations to inquire on their dependency on the nature of the alkoxo ligand and the metal.

or Pd, 1b26). Both compounds can be prepared either by fluoride exchange from the corresponding fluorocomplexes [(iPrPCP)M-F] with LiOMe or by proton exchange between the parent amides [(iPrPCP)M-NH2] and methanol.6b In practice, we found the latter method more convenient for routine synthetic procedures, since both Ni an Pd amides can be generated in situ from the readily available halide precursors [(iPrPCP)M-X] (M = Ni, X = Br, or M = Pd, X = Cl). Now, we extended this procedure to prepare complexes with higher alkoxo ligands (Scheme 1). The 31P NMR spectra of the crude Scheme 1

reaction mixtures confirm that these are very clean reactions but usually show a residual peak of the starting halide complex that apparently cannot be converted ( MeOH > EtOH > iPrOH (see literature pKa data in dimethyl sulfoxide (DMSO) listed in Table 3). The data fit the expected sequence for the latter three alcohols, but the situation is less clear for methoxyethanol. The latter is the stronger acid in the series (pKa = 26.5), yet the value of K1(OCH2CH2OMe) for Pd is similar to that of methanol (pKa = 27.5), and for Ni it is between those of methanol and ethanol (pKa = 28.2). Moreover, since 2methoxyethanol is a stronger acid than water (at least in DMSO), it would be reasonably expected that 2-methoxyethoxide complexes would be quite resistant to hydrolysis. In fact, the K1 value not only for 2-metoxyethanol but also for methanol seems surprisingly low considering that water (pKa = 27.9) is a weaker acid than both these alcohols in nonprotic organic solvents like DMSO. Note that this “common sense” interpretation of the hydrolytic stability of alkoxides M−OR based on the acid strengths of alcohols gives the right answer for the wrong reasons. Water is not the strongest acid in the series, but, on the contrary, it behaves as a very weak acid in nonprotic solvents. The common perception of water as a significantly stronger acid than most alcohols only becomes true in protic solvents, like water itself.36b,39 DFT Calculations. To gain a better understanding of the factors behind the trends in hydrolytic sensitivity of Ni and Pd

ΔH1° = DM−OH + DH−OR − DM−OR − DH−OH

(6)

ΔH1° = (DH−OR + DH−OH) − (DM−OR − DM−OH) = ΔDH − OR − ΔDM−OR G

(7) DOI: 10.1021/acs.inorgchem.7b01868 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

OMe). Unfortunately, the values of homolytic ΔDM−OR for the alkoxides are so closely similar that no clear trend emerges from their comparison, and the same can be said of ΔDH−OR. The latter finds support in the experimental gas-phase O−H BDE data for MeOH, EtOH, and iPrOH that are the same within 0.5 kcal·mol−1 (105.2, 105.4, and 105.7 kcal·mol−1, respectively).41 In contrast, a clearer trend emerges from the comparison of heterolytic ΔD. As commented before, homolytic BDEs are usually preferred for reaction energy analyses, but in this case a clearer trend emerges from the consideration of the heterolytic parameters. As can be seen in Table 5, both heterolytic ΔDM−OR and ΔDH−OR exhibit the sequence of MeO > EtO > iPrO. This is largely dictated by the relative stabilities of the RO− anions in benzene with regard to OH− (see eqs 8 and 9), which increases as the growing size of R favors negative charge dispersion through hyperconjugative or polarization effects.42 This view is availed by the previously discussed influence of the R group on experimental ν(O−R) frequencies. As ΔDM−OR decreases somewhat faster than ΔDH−OR, the values of ΔHo1, given by the difference ΔDH−OR − ΔDM−OR (eq 7), decrease in the same order. The fact that this relationship emerges plainly for heterolytic BDEs, and less clearly with homolytic parameters, reinforces the concept that Coulombic interactions dominate over steric effects or covalent bonding in defining the influence of the alkoxo R group on the ΔHo1 enthalpy balance. The more pronounced influence of variation of heterolytic BDE of M− OR bonds on Ni alkoxides than in their Pd analogues could be attributed to the slightly more electropositive character of Ni,43 but it is more likely the effect of shorter bond distances for M = Ni than M = Pd. The sketch shown in Figure 4 summarizes the influence of electrostatic forces on the strength of the M−O bonds.

[(iPrPCP)M−OR] + HO·/− → [(iPrPCP)M−OH] + RO·/− ΔDM − OR (homolytic or heterolytic) = ΔH8 (8)

RO−H + HO·/− → RO·/− + HO−H ΔDH−OR (homolytic or heterolytic) = ΔH9

(9)

The results of our computational study are summarized in Table 5. Computed free-energy balances for eq 1, ΔG1o, are shown in the first data block, where these are compared with experimental values (calculated as ΔGo1 = −RT ln(K1)). As can be seen, the agreement is excellent for Ni and somewhat poorer for Pd. At this point we cannot advance a definitive explanation for this, but it could be related to the use of a special basis function (def2-TZVP) to describe Pd atoms in energy calculations. General trends are reasonably well-reproduced in both series of data. Specifically, the larger experimental K1 equilibrium constants observed for Pd are reflected in the ΔGo1 figures that are systematically less positive for Pd than for Ni. The success of the theoretical model in reproducing the main experimental facts encouraged us to analyze in more detail the origin of the different hydrolytic sensitivities of the nickel and palladium alkoxides. Table 5 lists also the reaction energies (assimilated to reaction enthalpies ΔHo1) computed for eq 1, as well as their breakdown in terms of homolytic or heterolytic relative BDEs, ΔD. Compared to free energies, ΔGo1 (either experimental or theoretical), the computed enthalpies ΔHo1 are significantly more favorable (less positive) for Ni and even slightly negative for M = Pd, implying that the characteristic sensitivity of Ni and Pd alkoxides to hydrolysis is, in part, an entropy effect. The entropy contributions are small but not negligible, ranging between −5 and −9 e.u., except for the degenerate exchange processes of the first row in Table 5 that, by definition, are 0. These values are realistic and comparable to the −9 e.u. cited in the literature7a,25 for related exchange reactions. The close to zero enthalpy changes predicted by the theoretical model evidence the fact that the OH and OR group have similar electronegativity. Despite their small values, computational enthalpies parallel those of the experimental ΔG1o. This confirms that the qualitative order of hydrolytic stabilities (e.g., that isopropoxides are more readily hydrolyzed than methoxides) is not “entropy noise” but is largely determined by the R group of the alkoxides. Computed homolytic ΔDM−O for the alkoxide complexes show small variations, being close to −21 kcal·mol−1 for the Ni derivatives and to −19 kcal·mol−1 for the palladium ones. These relatively large and negative values indicate that M−O bonds of alkoxides are substantially weaker than those of hydroxides, although the R groups have all similar influence on homolytic BDEs. Noteworthy, this is the same trend observed for the σ-donor capacities of the OR ligands, reflected in the NMR Δδ(i-C) parameter. However, since alkoxides should be expected to behave as stronger σ-donors than hydroxide, the greater strength of M−OH bonds is more readily explained as an electrostatic effect than by covalent bonding. Thus, the Coulombic attraction between the positive metal fragment and the negative ligands must be more intense for the compact OH than for the larger and more polarizable alkoxide groups. Note that steric effects are unlikely to play a dominant role here; otherwise, BDEs would be expected to decrease steadily on increasing the hindrance of the OR group (i.e., iPrO < OEt
OEt > OiPr, but the 2-methoxyethoxides behave similarly to methoxides or ethoxides, although 2-methoxyethanol is the stronger acid in the series. A computational DFT model reproduced these experimental trends. On the basis of the results of the theoretical study, we propose that the influence of the metal and the R groups on the hydrolysis equilibrium is mainly due to electrostatic effects. Polarizable R groups disperse the negative charge at the oxygen atom, decreasing the ionic component of polar M−O and H−O bonds. However, the partially ionic M−O bonds of alkoxides are more destabilized by R than the prevailingly covalent H−O bonds of alcohols. These effects are more important for Ni than Pd, because electrostatic effects are more intense for the smaller-sized ion. However, in absolute terms, Pd−O bonds are somewhat weaker than Ni−O bonds, which could explain the lower thermal stability of Pd alkoxides.

Figure 5. Plot of the Mulliken charges on the alkoxide oxygen atom of nickel alkoxide complexes vs ΔHo1 calculated for the corresponding hydroxide/alcohol exchange processes.

alcohol is due to the ability of the 2-methoxy group to reduce the ionic contribution by removing negative charge from the metal-bound oxygen atom. In addition to relative BDEs, the last data block in Table 5 lists absolute homolytic BDE values calculated for the M−O bonds of the Ni and Pd hydroxides and alkoxides in benzene solution (DM−O). As commented before, reliable absolute BDE values are difficult to calculate, and the data may have important systematic errors due to the nonisodesmic character of the bond dissociation process. However, these errors are likely to cancel (at least in part) when comparing their values for analogous complexes of Ni and Pd. Noteworthy, covalent Pd−O bonds appear to be ca. 4% weaker than Ni−O bonds, in absolute terms. The difference has no direct influence on the hydrolytic stabilities of the alkoxides, as these are determined by the relative bond strengths ΔDM−OR and ΔDH−OR, but it can be related to the lower thermal stability of palladium alkoxides as compared to their nickel analogues. This explains the apparent contradiction between thermal stability and hydrolytic sensitivity that we mentioned before. Note that this reasoning makes a distinction between hydrolytic stability, a thermodynamic property linked to the relative stability of the M−OR and M−OH bonds, and kinetic properties, which depend on the destabilization of the alkoxide in absolute terms. For example, the lower stability of Pd−OR bonds probably enhances the nucleophilicity of palladium alkoxides in comparison with their nickel analogues (a typical ground-state ef fect).

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EXPERIMENTAL SECTION

All operations were performed under oxygen-free nitrogen atmosphere using conventional Schlenk techniques or a nitrogen-filled glovebox. Solvents were rigorously dried and degassed before use. Microanalyses were performed by the Analytical Service of the Instituto de ́ Investigaciones Quimicas. NMR spectra were recorded on Bruker DPX-300, DRX-400, and DRX-500 models. Chemical shifts (δ) are in parts per million. Solvents were used as internal standards for the reference of 1H and 13C spectra, but chemical shifts are referenced with respect to tetramethylsilane (TMS). 31P spectra are reported with respect to external H3PO4. 31P{1H} NMR in normal, nondeuterated solvents were performed in tetrahydrofuran (THF) using a capillary solution of PPh3 in C6D6 as external standard. Assignment of signals was assisted by combined one-dimensional (gated-13C) and twodimensional (2D) homonuclear 1H−1H COSY and phase-sensitive NOESY, and heteronuclear 13C−1H HSQC and HMBC heterocorrelations. IR spectra were recorded on a Bruker Tensor 27 FTIR spectrophotometer. Abbreviations for multiplicities are as usual, and the letters b and v denote broad and virtual coupling, respectively (e.g., bm, broad multiplet, dvt = doublet of virtual triplets). Apparent constants for virtual couplings are marked with asterisk (*). Complexes [iPrPCP]M-X (M = Ni, X = Br,6b OH (6a),22 OMe (1a);23 M = Pd, X = Cl,44 OH (6b),22 OMe (1b)26) were prepared as previously described. Spectral simulations were performed with gNMR.45 Synthesis of Alkoxide Complexes [(iPrPCP)Ni-OR] (R = Et, 2; n Bu, 3; iPr, 4; CH2CH2OH, 5). These complexes were similarly prepared from the corresponding halide precursor [(iPrPCP)Ni−Br] as an extension of the method described for the synthesis of 1a. The general procedure is illustrated with the preparation of compound 5: A solution of [(iPrPCP)Ni−Br] (238 mg, 0.5 mmol) in 20 mL of THF was transferred to a gastight centrifuge ampule (a specially designed glass ampule fitted with a gastight polytetrafluoroethylene (PTFE) needle valve), which had been previously loaded with an excess (10 equiv) of NaNH2. The ampule was sonicated in an ultrasound bath for 4 h. The resulting mixture was centrifuged, and the clear solution was transferred to a Schlenk tube via cannula. A solution of ethylene glycol (1 mL, 0.5 M, 0.5 mmol) in THF was added to the stirred solution, and the mixture was taken to dryness. The residue was extracted with Et2O, and the solvent was evaporated under vacuum to remove the remaining THF. The solid was dissolved in hexane, some hexamethyldisiloxane was added, and the solution was concentrated under reduced pressure. The product crystallized at −30 °C as dark

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CONCLUSIONS In this contribution, the factors that influence the hydrolytic stability of nickel and palladium alkoxides have been analyzed, focusing on monomeric complexes stabilized by the iPrPCP pincer ligand. Both types of compounds are quantitatively generated in solution when the corresponding amido complexes [(iPrPCP)M(NH2)] (M = Ni, Pd) react with the corresponding alcohols. A series of nickel alkoxide complexes were isolated and fully characterized. In contrast, palladium alkoxide complexes display low thermal stability, which combined with their high solubility in hydrocarbon solvents prevented their isolation as pure samples. Spectroscopic and structural data for the nickel complexes provide direct indications of the highly polar character of the Ni−O bonds. Both Ni and Pd alkoxides react readily with water, affording equilibrium mixtures containing the corresponding hydroxides. The reversibility of these reactions allowed to quantify the hydrolytic stability of the different alkoxides, measuring the I


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Inorganic Chemistry Table 6. Summary of Crystallographic Data and Structure Refinement Results for 2, 3, 4, and 5 compound empirical formula formula weight crystal system space group a, Å b, Å c, Å α, deg β, deg γ, deg volume, Å3 temperature, K Z Dcalcd, Mg/m3 μ, mm−1 F(000) crystal size, mm θ range data, deg no. of reflns collected no. of reflns used Rint parameters R1(F) [F2 > 2σ(F2)]a wR2(F2)b (all data) Sc (all data)

2 4(C22H40NiOP2)·C6H18OSi2 1927.13 triclnic P1̅ 10.8837(5) 15.5919(7) 17.2739(8) 81.9570(10) 73.3070(10) 76.7880(10) 2724.8(2) 173(2) 1 1.174 0.864 1042 0.25 × 0.21 × 0.14 2.16−25.25 44 511 9756 0.0341 563 0.0347 0.0890 1.076

3 C24H44NiOP2 469.24 orthorhombic Pca21 10.3283(6) 16.6955(10) 15.2466(9) 90 90 90 2629.1(3) 173(2) 4 1.186 0.871 1016 0.50 × 0.28 × 0.12 2.31−25.24 15 644 4207 0.0329 263 0.0368 0.0989 1.107

4 C23H42NiOP2 455.21 orthorhombic Pbca 10.3560(13) 15.4904(14) 32.133(3) 90 90 90 5154.8(9) 173(2) 8 1.173 0.887 1968 0.32 × 0.22 × 0.18 2.32−25.24 20 975 4668 0.0929 292 0.0684 0.1692 0.907

5 C22H40NiO2P2 457.19 orthorhombic Pbca 16.8895(8) 14.9835(8) 19.2232(9) 90 90 90 4864.7(4) 173(2) 8 1.248 0.943 1968 0.26 × 0.22 × 0.10 2.10−25.24 39 742 4400 0.0616 284 0.0351 0.0959 1.041

R1(F) = ∑(|F0| − |Fc|)/∑|F0| for the observed reflections [F2 > 2σ(F2)]. bwR2(F2) = {∑[w(F02 − Fc2)2]/∑w(F02)2}1/2. cS = {∑[w(F02 − Fc2)2/(n − p)}1/2; (n = number of reflections, p = number of parameters). a

red-orange crystals that were collected by filtration and dried under vacuum. Yield: 60%. Similar procedures were applied for the preparation of 2 (65%), 3 (58%), and 4 (63%). Spectroscopic and Analytical Data for 2−5. 2: 1H NMR (500 MHz, C6D6, 25 °C) δ 1.03 (dvt, 12H, 3JHH ≈ *JHP ≈ 6.7 Hz, CHMeMe); 1.42 (t, 3H, 3JHH = 6.6 Hz, OCH2CH3); 1.45 (dvt, 12H, 3 JHH ≈ *JHP ≈ 7.4 Hz, CHMeMe); 2.00 (m, 4H, CHMeMe); 2.73 (bm, 4H, CH2 (PCP)); 3.86 (q, 2H, 3JHH = 6.6 Hz, OCH2CH3); 6.86 (d, 2H, 3JHH = 7.3 Hz, CarHm); 7.01 (t, 1H, 3JHH = 7.3 Hz, CarHp). 13 C{1H} NMR (125.7 MHz, C6D6, 25 °C): δ 2.0 (s, OCH2CH3); 18.2 (s, CHMeMe); 19.0 (s, CHMeMe); 24.0 (vt, *JCP = 9.5 Hz, CHMeMe); 32.5 (vt, *JCP = 12.7 Hz, CH2 (PCP)); 67.9 (s, OCH2CH3); 121.1 (vt, *JCP = 8.4 Hz, CarHm); 124.4 (s, CarHp); 152.2 (vt, *JCP = 13.6 Hz, Car‑o); 158.2 (t, 2JCP = 17.8 Hz, Car‑i). 31 1 P{ H} NMR (202.4 MHz, C6D6, 25 °C): δ 53.1. IR (mineral oil, cm−1) 1054, 1102 (ν C−O). Anal. Calcd for C22H40NiOP2: C, 59.89; H, 9.14. Found, C, 59.90; H, 9.85%. 3: 1H NMR (500 MHz, C6D6, 25 °C) δ 1.04 (dvt, 12H, 3JHH ≈ *JHP ≈ 6.7 Hz, CHMeMe); 1.14 (t, 3H, 3JHH = 7.2 Hz, O(CH2)3CH3); 1.46 (dvt, 12H, 3JHH ≈ *JHP ≈ 7.2 Hz, CHMeMe); 1.75 (m, 2H, −CH2CH3); 1.75 (m, 2H, −CH2CH2−); 2.03 (m, 4H, CHMeMe); 2.74 (m, 4H, CH2 (PCP)); 3.79 (t, 2H, 3JHH = 5.9 Hz, OCH2); 6.86 (d, 2H, 3JHH = 7.3 Hz, CarHm); 7.01 (t, 1H, 3JHH = 7.3 Hz, CarHp). 13 C{1H} NMR (75.4 MHz, C6D6, 25 °C) δ 15.0 (s, O(CH2)3CH3); 18.3 (s, CHMeMe); 19.0 (s, CHMeMe); 20.4 (s, −CH2CH3); 24.1 (vt, *JCP = 9.5 Hz, CHMeMe); 32.5 (vt, *JCP = 12.7 Hz, CH2 (PCP)); 40.8 (s, -CH2CH2−); 72.3 (t, 2H, 3JCP = 4.7 Hz, OCH2); 122.0 (vt, *JCP = 8.2 Hz, CarHm); 124.4 (s, CarHp); 152.2 (vt, *JCP = 13.5 Hz, Car‑o); 158.1 (t, 2JCP = 17.9 Hz, Car‑i). 31P{1H} NMR (202.4 MHz, C6D6, 25 °C) δ 53.5. IR (mineral oil, cm−1) 1085, 1121 (ν C−O). Anal. Calcd for C24H44NiOP2: C, 61.43; H, 9.45. Found, C, 61.36; H, 9.33%. 4: 1H NMR (500 MHz, C6D6, 25 °C) δ 1.02 (dvt, 12H, 3JHH ≈ *JHP ≈ 6.6 Hz, CHMeMe); 1.39 (d, 6H, 3JHH = 5.6 Hz, OCHMe2); 1.46 (dvt, 12H, 3JHH ≈ *JHP ≈ 7.5 Hz, CHMeMe); 1.97 (m, 4H, CHMeMe); 2.72 (bm, 4H, CH2 (PCP)); 3.58 (sept, 1H, 3JHH = 5.5 Hz, OCHMe2); 6.85 (d, 2H, 3JHH = 7.2 Hz, CarHm); 7.00 (t, 1H, 3JHH

= 7.2 Hz, CarHp). 13C{1H} NMR (75.4 MHz, C6D6, 25 °C) δ 18.3 (s, CHMeMe); 19.0 (s, CHMeMe); 24.1 (vt, *JCP = 9.5 Hz, CHMeMe); 31.1 (s, OCHMe2); 32.6 (vt, *JCP = 12.9 Hz, CH2 (PCP)); 73.4 (t, 3JCP = 5.7 Hz, OCHMe2); 121.9 (vt, *JCP = 8.1 Hz, CarHm); 124.2 (s, CarHp); 152.1 (vt, *JCP = 13.3 Hz, Car‑o); 157.7 (t, 2JCP = 18.4 Hz, Car‑i). 31P{1H} NMR (202.4 MHz, C6D6, 25 °C) δ 53.0. IR (mineral oil, cm−1) 974, 1134, 1147 (ν C−O). Anal. Calcd for C23H42NiOP2: C, 60.68; H, 9.30. Found, C, 60.68; H, 9.29%. 5: 1H NMR (300 MHz, C6D6, 25 °C) δ 0.94 (dvt, 12H, 3JHH ≈ *JHP ≈ 6.9 Hz, CHMeMe); 1.39 (dvt, 12H, 3JHH ≈ *JHP ≈ 7.5 Hz, CHMeMe); 1.90 (m, 4H, CHMeMe); 2.67 (vt, 4H, *JHP = 3.8 Hz, CH2 (PCP)); 3.66 (m, 2H, CH2OH); 3.83 (m, 2H, OCH2); 4.01 (bt, 1H, 3JHH = 2.5 Hz, OH); 6.82 (d, 2H, 3JHH = 7.5 Hz, CarHm); 6.99 (t, 1H, 3JHH = 7.4 Hz, CarHp). 13C{1H} NMR (75.4 MHz, C6D6, 25 °C) δ 18.2 (s, CHMeMe); 19.0 (s, CHMeMe); 24.0 (vt, *JCP = 9.8 Hz, CHMeMe); 32.2 (vt, *JCP = 13.0 Hz, CH2 (PCP)); 65.6 (s, CH2OH); 70.6 (t, 3JCP = 4.1 Hz, OCH2); 122.1 (vt, *JCP = 8.2 Hz, CarHm); 124.8 (s, CarHp); 152.2 (vt, *JCP = 13.2 Hz, Car‑o); 156.1 (t, 2JCP = 18.1 Hz, Car‑i). 31P{1H} NMR (202.4 MHz, C6D6, 25 °C) δ 54.3. IR (mineral oil, cm−1) 1101, 1084, 1047 (ν C−O). Anal. Calcd for C22H40NiO2P2: C, 57.80; H, 8.82. Found, C, 57.66; H, 8.65%. Attempted Isolation of Palladium Alkoxide Complexes. The same methodology described for Ni alkoxide complexes was applied to [(iPrPCP)Pd−Cl]. Although, as reported before, this method allowed the isolation and full characterization of the correspondig methoxide 1b,26 no crystalline materials were isolated from alcohols different than MeOH. Prior to the workup procedure, the NMR spectra of crude reaction mixtures (THF) showed a largely prevailing signal shifted 2− 4 ppm upfield of that of the ubiquitous [(iPrPCP)Pd−OH] impurity, which was assumed to correspond to that of the corresponding alkoxide complex [(iPrPCP)Pd-OR]. In the case of the isopropoxide complex, a sample was evaporated and directly dissolved in C6D6. The 1 H NMR spectrum was somewhat broad but showed distinct signals for the OiPr ligand at positions that closely resemble those of its Ni analogue (see below). Crude materials (OR = OEt, OiPr) proved extremely soluble in hexane and hexane/hexamethyldisiloxane J


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Inorganic Chemistry mixtures. These solutions slowly darkened on prolonged storage in the freezer at −20 °C, suggesting that partial decomposition, as reported for 1b.9b The difficulty of growing crystals of these compounds probably arises from a combination of factors: low thermal stability, hydrolytic sensibility, high solubility, and that smaller amounts of Pd complexes are regularly used in comparison with Ni. [(iPrPCP)Pd−OiPr]. 1H NMR (300 MHz, C6D6, 25 °C) δ 0.93 (dvt, 12H, 3JHH ≈ *JHP ≈ 7.1 Hz, CHMeMe); 1.31 (dvt, 12H, 3JHH ≈ *JHP ≈ 8.0 Hz, CHMeMe); 1.45 (d, 6H, 3JHH = 6.0 Hz, OCHMe2); 1.97 (bm, 4H, CHMeMe); 2.81 (bm, 4H, CH2 (PCP)); 4.27 (sept, 1H, 3JHH = 5.8 Hz, OCHMe2); 7.00 (d, 2H, 3JHH = 7.2 Hz, CarHm); 7.05 (t, 1H, 3 JHH = 7.2 Hz, CarHp). 31P{1H} NMR (202.4 MHz, C6D6, 25 °C)δ 54.9. Titration of Hydroxide Complexes 6a and 6b with Alcohols. These reactions were performed in NMR tubes, using C6D6 as solvent. Alcohols used in this protocol were dried by distillation from the corresponding sodium alkoxides, generated in situ by dissolving pieces of sodium in the alcohol prior to the distillation. The following experimental protocol was applied: In the glovebox, a certain amount close to 0.025 mmol (6a: 10.33 mg; 6b 11.53 mg) of the corresponding hydroxide was accurately weighed in an analytical balance. The sample was dissolved in 0.6 mL of C6D6 and transferred to an NMR tube capped with a rubber septum. Another septumcapped NMR tube was used to store the required alcohol. The tubes were taken out of the box, and 31P{1H} and 1H NMR spectra of the complex sample solution were recorded on the 300 MHz NMR spectrometer. With gastight syringes of suitable volume, increasing amounts of the alcohol were taken from the reservoir tube and successively added to the sample of the complex (e.g., 0.5, 1.5, 3.5, 6.5 μL, etc). After each addition, the sample was gently shaken, returned to the NMR probe, and the 31P{1H} spectra were recorded. Control experiments were performed to check that the equilibration of the species was essentially instantaneous. The equilibrium constant K1 was calculated as [MOR]/[MOH] × [H2O]/[ROH] for each addition. The alkoxide to hydroxide ratio was directly obtained from the 31P integrals, and the ratio of water to alcohol was deduced from the total amount of alcohol added using stoichiometry relationships. For Pd complexes, the small amount of crystallization water (ca. 0.5 equiv, observed at δ 4.13 ppm) was taken into account in the calculations. X-ray Structure Analysis for 2, 3, 4, and 5. A summary of the crystallographic data and structure refinement results for 2, 3, 4, and 5 crystal compounds is given in Table 6. Crystals of these new compounds coated with dry perfluoropolyether were mounted on a glass fiber and fixed under a cold nitrogen stream. The intensity data were collected on a Bruker-Nonius X8ApexII CCD area detector diffractometer using Mo Kα radiation source (λ = 0.710 73 Å) fitted with a graphite monochromator. The data collection strategy used was ω and ϕ rotations with narrow frames (width of 0.50°).46 Instrument and crystal stability were evaluated from the measurement of equivalent reflections at different measuring times, and no decay was observed. The data were reduced using SAINT and corrected for Lorentz and polarization effects, and a semiempirical absorption correction was applied (SADABS).47 The structures were solved by direct methods using SIR-200248 and refined against all F2 data by fullmatrix least-squares techniques using SHELXL-2016/649 minimizing w[Fo2 − Fc2]2. All the non-hydrogen atoms were refined with anisotropic displacement parameters. The hydrogen atoms of these compounds were included from calculated positions and allowed to ride on the attached atoms with isotropic temperature factors (Uiso values) fixed at 1.2 times (1.5 times for methyl groups) those of Ueq values of the corresponding attached atoms. 2 CCDC 1541332, 3 CCDC 1541333, 4 CCDC 1541334 and 5 CCDC 1541335. Computational Details. All DFT calculations were performed using the Spartan16 software package.50 A preliminary screening study was performed to select an efficient method for geometry optimization and single-point energy calculations. We found that a combination of the GGA functional PBE and 6-31G* leads to results comparable to those of B3LYP/6-31+G* at a fraction of the computation time. Vibrational analyses were systematically performed to confirm that the optimized geometries correspond to true stationary points (no

imaginary frequencies found). A set of representative bond distances and angles was chosen to compare with the experimental geometry data (see Table S2, Supporting Information). In addition, we observed that IR frequencies of C−O bonds were reproduced by the PBE functional, especially in combination with the 6-31+G* basis set. Definitive geometries used in the model were optimized with the PBE functional and 6-31G* basis set for organic compounds and Ni complexes. For palladium compounds, the basis set LACVP* was used. This comprises the same 6-31G* basis functions for all light atoms and the LANL2DZ effective core for elements beyond Kr. Solvent effects (benzene) were approximated with the Conductor-Like Polarizable Continuum Model (CPCM). Bondi-type PCM Radii for Ni and Pd, not included in the Spartan package, were set as 2.28 and 2.48 Å, respectively, based on an estimation of the covalent radii in the X-ray structures, incremented by 20%. Depending on the type of molecule, the starting geometries were built using different strategies. For organic molecules, the most stable conformer was selected after a search with the Merck Molecular Force Field (command FINDBEST = MMFF); for nickel hydroxide and alkoxide complexes, experimental X-ray structures were used, and for the palladium complexes the starting geometries were built using the nickel complexes as the starting point. Initial points for the 2-methoxyethyl derivatives were built replacing the OH in the X-ray structure of compound 5 with a OCH3, and then allowing the organic ligand to relax to the all-anti configuration. Energies were computed performing a single-point calculation on the previous geometries at the PBE/6-311++G(3df,2p)/CPCM level, adding the thermal correction from the double-ζ calculation (thermal correction = ΔG° − E(SCF)). At this level of the theory, Spartan uses the all-electron def2-TZVP basis set to describe the Pd atom. Calculation times were significantly reduced using the “dual” option, in which self-consistent field (SCF) convergence is first achieved at the 6-311G* level, and then corrected perturbatively for the effect of the basis set extension. Additionally, this procedure has the advantage of extending the application to the element Pd, as the extended version of the def2-TZVP basis set is currently not available. Independent tests showed that in the nickel case, the use of the “dual” option does not introduce significant variations in the results. A modification of the above calculation method was used to improve the description of intramolecular hydrogen bond in 5. In this case, the geometry was optimized with the PBE functional including Grimme’s empirical correction for dispersive forces (PBE-D3), together with the larger basis set 6-31+G** and the PCM solvent model (benzene). Nonessential iPr substituents were replaced by Me to reduce the calculation time. The stabilization gained by the hydrogen bond was estimated based on ΔG° comparison with that of the conformer with the Ni−O−CH2CH2OH group set in a linear conformation without hydrogen bond. Further details and complete molecular coordinates for optimized geometries can be found in the Supporting Information.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.7b01868. Further computational details and description of DOSY experiments (PDF) Accession Codes

CCDC 1541332−1541335 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. K


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(6) (a) Smith, D. A.; Beweries, T.; Blasius, C.; Jasim, N.; Nazir, R.; Nazir, S.; Robertson, C. C.; Whitwood, A. C.; Hunter, C. A.; Brammer, L.; Perutz, R. N. The Contrasting Character of Early and Late Transition Metal Fluorides as Hydrogen Bond Acceptors. J. Am. Chem. Soc. 2015, 137, 11820−11831. (b) Martínez-Prieto, L. M.; Melero, C.; Del Río, D.; Palma, P.; Cámpora, J.; Á lvarez, E. Synthesis and Reactivity of Nickel and Palladium Fluoride Complexes with PCP Pincer Ligands. NMR-Based Assessment of Electron-Donating Properties of Fluoride and Other Monoanionic Ligands. Organometallics 2012, 31, 1425−1438. (7) (a) Holland, P. L.; Andersen, R. A.; Bergman, R. G.; Huang, J.; Nolan, S. P. Monomeric Cyclopentadienylnickel Methoxo and Amido Complexes: Synthesis, Characterization, Reactivity, and Use for Exploring the Relationship between H−X and M−X Bond Energies. J. Am. Chem. Soc. 1997, 119, 12800−12814. (b) Holland, P. L.; Andersen, R. A.; Bergman, R. G. Application of the E-C Approach to Understanding the Bond Energies Thermodynamics of Late-Metal Amido, Aryloxo and Alkoxide complexes: An Alternative to pπ/dπ Repulsion. Comments Inorg. Chem. 1999, 21, 115−129. (8) (a) Hanley, P. S.; Hartwig, J. F. Migratory Insertion of Alkenes into Metal−Oxygen and Metal−Nitrogen Bonds. Angew. Chem., Int. Ed. 2013, 52, 8510−8525. (b) Hartwig, J. F. Electronic Effects on Reductive Elimination to Form Carbon-Carbon and CarbonHeteroatom Bonds from Palladium(II) Complexes. Inorg. Chem. 2007, 46, 1936−1947. (c) Hartwig, J. F. Covalent (X-type) Ligands Bound Trough Metal-Heteroatom Bonds. In Organotransition Metal Chemistry, From Bonding to Catalysis; University Science Books, 2010; pp 147−247. (9) (a) Matas, I.; Cámpora, J.; Palma, P.; Á lvarez, E. Decomposition of Methylnickel(II) Amido, Alkoxo, and Alkyl Complexes by βHydrogen Elimination: A Comparative Study. Organometallics 2009, 28, 6515−6523. (b) Martínez-Prieto, L. M.; Á vila, E.; Palma, P.; Á lvarez, E.; Cámpora, J. β-Hydrogen Elimination Reactions of Nickel and Palladium Methoxides Stabilised by PCP Pincer Ligands. Chem. Eur. J. 2015, 21, 9833−9849. (10) Noyori, R.; Hashiguchi, S. Asymmetric Transfer Hydrogenation Catalyzed by Chiral Ruthenium Complexes. Acc. Chem. Res. 1997, 30, 97−102. (11) Hamid, M. H. S. A.; Slatford, P. A.; Williams, J. M. J. Borrowing Hydrogen in the Activation of Alcohols. Adv. Synth. Catal. 2007, 349, 1555−1575. (12) (a) Morton, D.; Cole-Hamilton, D. J.; Utuk, I. D.; PanequeSosa, M.; Lopez-Poveda, M. Hydrogen Production from Ethanol Catalysed by Group 8 Metal Complexes. J. Chem. Soc., Dalton Trans. 1989, 489−495. (b) Johnson, T. C.; Morris, D. J.; Wills, M. Hydrogen generation from formic acid and alcohols using homogeneous catalysts. Chem. Soc. Rev. 2010, 39, 81−88. (13) (a) Nishimura, T.; Onoue, T.; Ohe, K.; Uemura, S. Palladium(II)-Catalyzed Oxidation of Alcohols to Aldehydes and Ketones by Molecular Oxygen. J. Org. Chem. 1999, 64, 6750−6755. (b) Muzart, J. Molecular Oxygen to Regenerate PdII Active Species. Chem. - Asian J. 2006, 1, 508−515. (c) Parmeggiani, C.; Cardona, F. Transition Metal Based Catalysts in the Aerobic Oxidation of Alcohols. Green Chem. 2012, 14, 547−564. (14) Hartwig, J. F. Carbon-Heteroatom Bond Formation Catalysed by Organometallic Complexes. Nature 2008, 455, 314−322. (15) (a) Hartwig, J. F.; Richards, S.; Barañano, D.; Paul, F. Influences on the Relative Rates for C−N Bond-Forming Reductive Elimination and β-Hydrogen Elimination of Amides. A Case Study on the Origins of Competing Reduction in the Palladium-Catalyzed Amination of Aryl Halides. J. Am. Chem. Soc. 1996, 118, 3626−3633. (b) Widenhoefer, R. A.; Buchwald, S. L. Electronic Dependence of C−O Reductive Elimination from Palladium (Aryl)neopentoxide Complexes. J. Am. Chem. Soc. 1998, 120, 6504−6511. (16) (a) Trend, R. M.; Stoltz, B. M. Structural Features and Reactivity of (Sparteine)PdCl2: A Model for Selectivity in the Oxidative Kinetic Resolution of Secondary Alcohols. J. Am. Chem. Soc. 2008, 130, 15957−15966. (b) Kapteijn, G. M.; Spee, M. P. R.; Grove, D. M.; Kooijman, H.; Spek, A. L.; van Koten, G. Chemistry of

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Juan Cámpora: 0000-0002-5596-0178 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The Spanish Ministerio de Economiá y Competitividad (MINECO) and the FEDER funds of the European Union (CTQ2015-68978-P) supported this work.

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Inorganic Chemistry the Hydrolytic Stability of Zinc Alkoxide Species. J. Am. Chem. Soc. 2002, 124, 9970−9971. (39) The pKa of water is strongly influenced by solvation effects. In protic solvents, H2O dissociation is highly favored due to the efficient solvation of the small hydroxyl anion; conversely, formation of a “naked” OH− in a low-polarity solvent is energetically prohibitive. See, for example: Cox, B. G. Acids and Bases. Solvents Effects on Acid-Base Strength; Oxford University Press: Oxford, England, 2013. (40) Hehre, W. J. In A Guide to Molecular Mechanics and Quantum Chemical Calculations; Wavefunction, Inc.: Irvine, CA, 2003; pp 221− 252. (41) Luo, Y. R. Comprehensive Handbook of Chemical Bond Energies; CRC Press: Boca Raton, FL, 2007. (42) Note that heterolytic H−O BDEs for water or alcohol anion are defined as the reversal of protonic affinity for the hydroxide or alkoxide anions, a concept equivalent to Brönsted basicity in the gas phase. As a consequence, the hydrolytic stability order discussed in this work (i.e., M−OMe > M−OEt > M−OiPr) is the opposite that one would expect from alcohol “acid strength” in the gas phase (and probably in nonprotic, low-polarity solvents such as benzene, as our DFT calculations show). Note, however, that our interpretation should be restricted to the comparison of relative stabilities of aliphatic alkoxides, or containing weakly polar groups, such as a CH2CH2OMe fragment, because the polarity of the M−O bonds is very similar. When the M− O bonds arise from stronger Brönsted acids, such as perfluoroalcohols, phenols, or caboxylic acids, the ionic contribution to bond energy is dominated by the different electronegativities of the R groups or extensive electronic delocalization effects. In these cases, the usual qualitative relationships between acid strengths and hydrolytic stability hold plainly. (43) Electronegativities of divalent Ni and Pd in square-planar tetracoordinate environments are 1.470 and 1.483: Li, K.; Xue, D. Estimation of Electronegativity Values of Elements in Different Valence States. J. Phys. Chem. A 2006, 110, 11332−11337. (44) Frech, C. M.; Shimon, L. J. W.; Milstein, D. Redox-Induced Collapse and Regeneration of a Pincer-Type Complex Framework: A Nonplanar Coordination Mode of Palladium(II). Angew. Chem., Int. Ed. 2005, 44, 1709−1711. (45) Budzelaar, P. H. M. NMR Simulation: gNMR; https://home.cc. umanitoba.ca/~budzelaa/gNMR/gNMR.html. (46) Bruker APEX2; Bruker AXS, Inc.: Madison, WI, 2007. (47) Bruker Advanced X-ray Solutions. SAINT and SADABS programs; Bruker AXS, Inc.: Madison, WI, 2004. (48) Burla, M. C.; Camalli, M.; Carrozzini, B.; Cascarano, G. L.; Giacovazzo, C.; Polidori, G.; Spagna, R. SIR2002: the Program. J. Appl. Crystallogr. 2003, 36, 1103−1104. (49) Sheldrick, G. M. Crystal Structure Refinement with SHELXL. Acta Crystallogr., Sect. C: Struct. Chem. 2015, 71, 3−8. (50) (a) Spartan16, 2.0.3; Wavefunction, Inc.: Irvine, CA, 2016. (b) Shao, Y.; Fusti-Molnar, L.; Jung, Y.; Kussmann, J.; Ochsenfeld, C.; Brown, S. T.; Gilbert, A. T. B.; Slipchenko, L. V.; Levchenko, S. V.; O’Neill, D. P.; DiStasio, R. A., Jr; Lochan, R. C.; Wang, T.; Beran, G. J. O.; Besley, N. A.; Herbert, J. M.; Yeh Lin, C.; Van Voorhis, T.; Chien, S. H.; Sodt, A.; Steele, R. P.; Rassolov, V. A.; Maslen, P. E.; Korambath, P. P.; Adamson, R. D.; Austin, B.; Baker, J.; Byrd, E. F. C.; Dachsel, H.; Doerksen, R. J.; Dreuw, A.; Dunietz, B. D.; Dutoi, A.; Furlani, T. R.; Gwaltney, S. R.; Heyden, A.; Hirata, S.; Hsu, C. P.; Kedziora, G.; Khalliulin, R. Z.; Klunzinger, P.; Lee, A. M.; Lee, M. S.; Liang, W. Z.; Lotan, I.; Nair, N.; Peters, B.; Proynov, E. I.; Pieniazek, P. A.; Rhee, Y. M.; Ritchie, J.; Rosta, E.; Sherrill, D.; Simmonett, A. C.; Subotnik, J. E.; Woodcock, H. L., III; Zhang, W.; Bell, A. T.; Chakraborty, A. K.; Chipman, D. M.; Keil, F. J.; Warshel, A.; Hehre, W. J.; Schaefer, H. F., III; Kong, J.; Krylov, A. I.; Gill, P. M. W.; HeadGordon, M. Advances in Methods and Algorithms in a Modern Quantum Chemistry Program Package. Phys. Chem. Chem. Phys. 2006, 8, 3172−3191.

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