Establishing the Hydride Donor Abilities of Main Group Hydrides

Article pubs.acs.org/Organometallics

Establishing the Hydride Donor Abilities of Main Group Hydrides Zachariah M. Heiden* and A. Paige Lathem Department of Chemistry, Washington State University, Pullman, Washington 99164, United States S Supporting Information *

ABSTRACT: Interest in reductions with main group hydrides has been reinvigorated with the discovery of frustrated Lewis pairs. Computational analysis showed that the borohydride of the commonly used Lewis acid B(C6F5)3 was determined to be 15 kcal/mol less reducing than borohydride ([BH4]−), 22 kcal/mol less reducing than aluminum hydride ([AlH4]−), and 41 kcal/mol less reducing than superhydride ([HBEt3]−). In addition to [HB(C6F5)3]−, a hydride donor ability scale with an estimated error of ∼3 kcal/mol includes 132 main group hydrides with gradually changing reducing capabilities spanning 160 kcal/mol. The scale includes representatives from organosilanes, organogermanes, organostannanes, borohydrides, boranes, aluminum hydrides, NADH analogues, and CH hydride donors. The large variety of reducing agents and the wide span of the scale (ranging from 0.5 to 160 kcal/mol in acetonitrile) make the scale a useful tool for the future design of metal-based or main group reducing agents.

■

noncoordinating base.10 DuBois and co-workers have worked extensively in this field, employing the hydride donor ability of metal hydride complexes to better understand the activation of dihydrogen with metals.11 Using the pKa for the deprotonation of a metal hydride (eq 1), the two-electron oxidation of a metal complex (eq 2), and the reduction of a proton to a hydride (eq 3), the hydride donor ability of a metal hydride can be estimated (eq 4).11

INTRODUCTION Sodium borohydride has become the most important hydridebased reductant used on the industrial scale, with a market share exceeding 50%.1 Some of the key drivers governing its interest are that it is the least expensive commercially available metal hydride (on a hydride equivalent basis), it is safe with regard to storage and use, the industrial implementation requires no or limited equipment investment, the ease of workup (boron salts), solvents such as water and methanol can be employed, and both chemo- and diastereoselectivity can be achieved.1 Although [BH4]− is the reducing agent of choice from an industrial standpoint, several other main group hydrides have been widely used in the reduction of organic substrates,2 as [BH4]− is not the ideal reducing agent for every chemical reaction. Main group hydrides are primarily used as stoichiometric reagents in both industrial and academic applications, but with the advent of frustrated Lewis pairs (FLPs)3 the possibility of catalytic main group reducing agents has been realized. The emphasis on chemical reactions avoiding precious metals has become a concern of the pharmaceutical industry due to increased restrictions on trace-metal impurities.4 The discovery of catalytic main group reducing agents opens up the possibility of their use in reductions on the industrial scale. Chase and co-workers have recently been able to show that the use of a chemical scavenger promotes the catalytic hydrogenation of imines using a sterically bulky borane at catalyst loadings as low as 0.5 mol %.5 Although the use of main group Lewis acids to promote the catalytic reduction of organic substrates seems promising, their reducing capacity for organic substrates has not been quantified beyond the exploration of substrate scope.6−9 Interest in the quantification of hydride transfer from elemental hydrides has originated through the investigation of the heterolytic activation of H2 with a metal complex and a © XXXX American Chemical Society

LnMH+ → LnM + H+ LnM → LnM2 + + 2e−

H+ + 2e− → H−

G° = 1.37(pK a) G° = 46.1[E°(II/0)]

G° = − 46.1[E°(H+/H−)]

(1) (2) (3)

sum: LnMH+ → LnM2 + + H− G° = 1.37(pK a) + 46.1[E°(II/0)] + 79.6

(4)

Utilizing eqs 1−4, DuBois and co-workers have been able to show metal hydrides with hydride donor abilities ranging from 26 to 89 kcal/mol (see Supporting Information, Table S4). Although, these values are helpful in the determination of H2 activation, they provide little insight into the capability of these metal hydride complexes as reducing agents in organic reactions. This discrepancy becomes problematic, as evaluating the reducing capacity of metal hydrides versus main group hydrides, which are widely used in organic and industrial reductions, would greatly aid in the design of metal-based and metal-free reduction catalysts. Received: November 14, 2014

A

DOI: 10.1021/om5011512 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

Geometry optimizations were undertaken using the M06-2X/631G(d,p) level of theory in the gas phase.27 The M06-2X functional has been shown to be one of the more accurate DFT functionals for main group and FLP chemistry.17,20,28,29 More accurate energies were obtained using single-point energy calculations at the M06-2X/6-311G++(d,p) level of theory in MeCN, employing the optimized molecular geometries at the M06-2X/6-31G(d,p) level of theory. Acetonitrile was chosen as the solvent for computational analysis to evaluate the calculated hydride donor abilities versus values previously reported for metal complexes (see Supporting Information). Although the coordination of acetonitrile to the resulting boranes was considered, CD3CN solutions of HSiEt3 and [lutidinium][HB(C6F5)3] showed no hydride loss over the course of a week and were assumed to have no effect on the hydride donor ability of the examined main group hydride. A similar assumption is used in the determination of the hydride donor ability of metal hydride complexes.30−32 The PCM-UFF solvation model (default for Gaussian 09) was used.27 The IEFPCM-UA0 solvation model has been previously employed in the determination of hydride donor abilities of transition metal complexes,30,32 but the PCM-UFF solvation model was found to give slightly better results in the calculated hydride donor ability of [HBEt3]− (ΔGH− = 26 kcal/mol, experimentally);16 see below. Common Main Group Hydrides in Organic Reductions. The most common reducing agents employed in industrial scale organic synthesis are NaBH4, NaBH3CN, LiAlH4, and NaHBEt3 (superhydride).2 From a reactivity standpoint, superhydride is considered a stronger reducing agent than lithium aluminum hydride, and lithium aluminum hydride is a stronger reducing agent than borohydride,2,33,34 but to date, measurement of the reducing power of these main group hydrides has been qualitative as opposed to quantitative. In this study, we aimed to quantify the reducing power of main group hydrides through computational analysis. To determine the validity of our computational values, we initially determined the hydride donor ability for [HBEt3]−, which was computed to be 24 kcal/mol. The computed value of 24 kcal/ mol agrees fairly well with the experimentally determined value of 26 kcal/mol,16 which is estimated to have an experimental error of ±2 kcal/mol. To further validate the computational model, we compared a computational versus experimental relative Lewis acidity of a borane of interest versus B(C6F5)3. The ratio of the computed hydride donor abilities of the borohydride of interest to [HB(C6F5)3]− was determined to yield a computational relative Lewis acidity to B(C6F5)3 (see Supporting Information, Table S2). Comparison of the experimentally determined Lewis acidities, determined from the Gutmann−Beckett method,23−25,35 generated comparable Lewis acidities (within about 4%) to the computational values (see Supporting Information, Table S3). The resulting differences between the experimental and computational values yield an estimated error of about 3 kcal/mol for the calculated hydride donor abilities. Upon achieving satisfactory results with the chosen level of theory, we expanded this analysis to the common reducing agents in synthetic laboratories and borohydrides relevant to the chemistry of frustrated Lewis pairs. When comparing the reducing ability of borohydride, computations showed that [BH4]− (ΔGH− = 50 kcal/mol) was about 26 kcal/mol less likely to transfer a hydride than [HBEt3]−. The ability to transfer a hydride became less favorable by 18 kcal/mol than

The use of the thermodynamic cycle described in eqs 1−4 can be very helpful in describing the hydride donor abilty of transition metal complexes, but becomes more problematic in the analysis of main group hydrides. A problem that arises in attempting to use eqs 1−4 for main group hydrides is that most main group hydrides cannot be deprotonated in the presence of a base. Also, electrochemical analysis of the parent Lewis acids often results only in a one-electron reduction wave,12−14 which does not make eq 2 valid for main group hydrides. To remedy the inability to use a thermodynamic cycle for the determination of the hydride donor ability of main group hydrides, one could examine the equilibrium constant between a metal complex with a known hydride donor ability and a main group Lewis acid. This technique has been employed in the analysis of the hydride donor ability of superhydride ([HBEt3]−), where a solution of BEt3 was titrated with HRh(dmpe)2, resulting in an estimated hydride donor ability of 26 kcal/mol for [HBEt3]−.15,16 This technique can be problematic with main group hydrides, as the hydride donor ability of metal complexes tends to be measured in polar solvents, which are not compatible with highly Lewis acidic main group complexes. To date, [HBEt3]− is the only main group hydride species where the hydride donor ability has been experimentally determined. Equation 4 shows the hydride donor ability of metal hydrides; in a similar fashion, eq 5 can be applied to determine the hydride donor ability of main group hydrides (ΔGH−). [Lewis Acid H]− → Lewis Acid + H−

ΔG H −

(5)

To investigate the hydride donor ability of main group metal hydrides and obtain reasonable values to be utilized in experimental design, computational methods were employed with the advantage of avoiding experimental dilemmas as described above. Pápai and co-workers have previously computed the ability of boranes to accept a hydride to aid in the understanding of H2 activation with FLPs,17,18 but this study was limited to the analysis of only 12 boron-based Lewis acids. Krossing and co-workers have very recently investigated the effect of alkoxy substitution on a boron-based Lewis acid both experimentally and computationally, in addition to computing the gas phase hydride, fluoride, chloride, and methyl ion affinity of 34 other main group compounds using high-level calculations.19 Gilbert (computationally)20 and Piers (experimentally)21 have probed the effect of fluorine substitution on the Lewis acidity of fluorinated boranes, but generated contradicting conclusions. To experimentally quantify the Lewis acidity of newly synthesized Lewis acids, the Gutmann−Beckett or Childs’ Lewis acidity tests are utilized.22−25 Although these tests can provide insight into the Lewis acidity of a Lewis acid of interest with respect to B(C6F5)3, often conflicting results can be obtained.26 Herein we expand upon the current reports and describe the hydride donor abilities of main group hydrides with reducing capabilities spanning 160 kcal/mol, including representatives from silanes, boranes, borohydrides, aluminum hydrides, NADH analogues, and CH hydride donors. In addition to the described main group hydrides, we also investigate chiral main group hydrides to gain insight into the observed enantioselectivities in reduction reactions.

■

RESULTS AND DISCUSSION The reducing capacity of 132 main group hydrides was investigated using ab initio calculations and Gaussian 09. B


Article

Organometallics [BH4]− when a hydride was replaced with a cyanide ligand in [BH3CN]−. As expected, the addition of a more electronwithdrawing group, such as [CN]−, resulted in a reduced hydride donor ability, while the addition of an electrondonating ethyl group showed an increase in the calculated hydride donor ability. Aluminum hydride, which is widely considered to be more reducing than [BH4]−,2 was found to be only 7 kcal/mol more reducing than [BH4]− and 19 kcal/mol less reducing than [HBEt3]−. To experimentally verify the computationally determined hydride donor ability of [AlH4]−, a tetrahydrofuran solution of LiAlH4 (ΔGH− = 43 kcal/mol) and BPh3 (ΔGH− = 36 kcal/mol) resulted in no hydride transfer to the boron center, even up to a temperature of 50 °C, indicating that the hydride prefers to reside on the aluminum center. In the case of reductive aminations, some organic chemists have preferred the use of NaHB(OAc)3, as opposed to the potential cyanide source, NaBH3CN.36 Computational analysis showed that [HB(OAc)3]− has a similar reducing power to [BH4]− and is about 19 kcal/mol more favorable to transfer a hydride than [BH3CN]−. Hydride Donor Ability of [HB(C 6 F 5 ) 3 ] − and [HBR(C6F5)2]−. The recent discovery of frustrated Lewis pairs has led to the interest in the reducing capacity of the sterically bulky borohydride, [HB(C6F5)3]−.6 Computational analysis of the hydride donor ability of [HB(C6F5)3]− yielded a hydride donor ability of 65 kcal/mol, which is 15 kcal/mol less reducing than [BH4]− and 3 kcal/mol more reducing than [BH3CN]−. The diminished ability of [HB(C6F5)3]− to transfer a hydride when compared to [BH4]− is attributed to the electronwithdrawing nature of the pentafluorophenyl groups, which increases the Lewis acidity at the boron center. Several modifications of B(C6F5)3 have been recently undertaken to alter the ability to heterolytically activate H2 with a sterically bulky base and to promote the reduction of organic substrates. Erker and co-workers recently showed that the 1,1-carboboration of diphenylacetylene or 3-hexyne with B(C6F5)3 results in alkenylboranes.37 Computational analysis of the respective borohydrides shows that the 1,1-carboboration of B(C6F5)3 with diphenylacetylene or 3-hexyne leads to borohydrides that are about 14 or 9 kcal/mol more reducing, respectively, than [HB(C6F5)3]−. The generation of less electron-withdrawing ligands and increased steric bulk38 through carboboration suggests the production of a more reducing borohydride species. Other groups have approached reducing the Lewis acidity by substitution at one of the pentafluorophenyl groups of [HB(C6F5)3]− with a phenyl group ([HB(C6F5)2(Ph)]−).26 This substitution resulted in a borohydride complex about 10 kcal/mol more favorable to transfer a hydride than [HB(C6F5)3]−. Subsequent phenyl for pentafluorophenyl group exchanges continued to exhibit complexes more reactive toward hydride loss by 10 and 9 kcal/mol for [HB(C6F5)(Ph)2]− and [HBPh3]−, respectively. Yoon and co-workers have recently employed K[HBPh3] in the reduction of ketones.39,40 The increase in reactivity of the respective borohydrides toward hydride loss is expected by changing from the strongly electronwithdrawing pentafluorophenyl group to the less electronwithdrawing phenyl group. Soos and co-workers have recently demonstrated the use of increased steric bulk to promote the selective activation of H2 with small Lewis bases.41 Computational analysis showed that [HB(C6F5)2(Mes)]− was 14 kcal/mol more reducing than [HB(C6F5)3]−. Addition of subsequent mesityl groups resulted

in borohydrides 26 and 43 kcal/mol more likely to transfer a hydride than [HB(C6F5)3]−. As expected, the increased electron-donating ability and increased steric bulk of the mesityl ligand38 favor hydride loss and generation of a threecoordinate borane when compared to the respective phenyl derivatives. An alternative approach to increasing the steric bulk surrounding the boron center was recently described by Neu and Stephan, through the insertion of sterically bulky diazomethane groups into a B−C6F5 bond.42 Insertion of (trimethylsilyl)diazomethane or phenyldiazomethane into a B− C6F5 bond of B(C6F5)3 resulted in complexes with the respective borohydrides being 6 or 5 kcal/mol more favorable then [HB(C6F5)3]−, respectively, to transfer a hydride. As anticipated, the increase in steric bulk upon insertion of an electron-rich carbene fragment38 into a B−C6F5 bond resulted in a more reducing borohydride species by reducing the steric bulk around the boron center upon loss of a hydride ligand. In an attempt to further probe the effects of borane substituents in FLP chemistry, several groups have generated a family of highly Lewis acidic boranes with varying degrees of halogenated ligands. One of the simplest approaches to manipulating the chemistry of a Lewis acidic borane is the generation of boranes via hydroboration.43 Hydroboration of norbornene, cyclohexene, and styrene with [HB(C6F5)2]244,45 resulted in boranes with their respective borohydrides being 12, 12, and 10 kcal/mol more likely to give up a hydride than [HB(C6F5)3]−. These results correspond with an increase in electron density on the boron center by the introduction of a greater electron-donating substituent. Stephan and co-workers recently reported that the replacement of the para-fluorine atom of [HB(C6F5)3]− with a hydrogen atom leads to a FLP exhibiting reversible H2 activation and a borane 97% the Lewis acidity of B(C6F5)3.46 Computational analysis showed that the introduction of a hydrogen atom in the para-position of a triarylfluoroborane, [HB(C6F4-p-H)3]−, promotes the loss of a hydride by only 2 kcal/mol when compared to [HB(C6F5)3]−, which is on par with the experimental Lewis acidity determination (see Supporting Information). Piers recently reported the related B(C6F4-o-H)3,21 where computational analysis of the hydride donor ability of the respective borohydride resulted in a borohydride 2 kcal/mol more likely to transfer a hydride than [HB(C6F4-p-H)3]−. The increased ability to transfer a hydride from [HB(C6F4-o-H)3]− than from [HB(C6F4-p-H)3]− verifies the experimental observations where substitution in the para-position as opposed to the ortho-position resulted in a slightly more Lewis acidic borane.21 In an analogous conversion, replacing the para-hydrogen of [HBPh3]− with a fluorine in [HB(C6H4-p-F)3]− results in a borohydride 1.5 kcal/mol less likely to lose a hydride than [HBPh3]−, which is of similar magnitude to the substitution of the para-fluorine atom of [HB(C6F5)3]− with hydrogen. Oestreich and Paradies recently reported the synthesis of tris(2,4,6-trifluorophenyl)borane and tris(2,6-difluorophenyl)borane,47,48 respectively, where the respective boranes were experimentally determined to exhibit 85% and 82% the Lewis acidity of B(C6F5)3. Computational determination of the transfer of a hydride from [HB(2,4,6-F3-C6F2)3]− and [HB(2,4,6-F3-C6F2)3]− resulted in borohydrides more favorable to transfer a hydride by 13 and 14 kcal/mol, respectively, in comparison to [HB(C6F5)3]−, which suggests similar Lewis acidities to the experimental values (see Supporting Information, Table S2). Substitution of a fluorine group of B(C6F5)3 in the ortho-position with a pentafluorophenyl group has been C


Article

Organometallics

in a borohydride 2 kcal/mol more reducing. In the case of placing an electron-withdrawing pentafluorophenyl group on an alkyl borane, such as 9-BBN (where BBN = borabicyclo[3.3.1]nonane), hydride transfer from the resulting borohydride complex decreased by 14 kcal/mol when compared to [HBEt3]−, exhibiting a hydride donor ability similar to [HB(C6H4-p-F)3]− and [HB(Mes)2(C6F5)]− (see Supporting Information). The decrease in ability to transfer a hydride from [(B-C6F5)(9-H-BBN)]− is of similar magnitude to the removal of a pentafluorophenyl group through a 1,1-carboboration of B(C6F5)3. Alkoxy Borohydrides. Although alkoxy and aryloxy boranes have not met much success in the activation of dihydrogen with FLPs, primarily due to ligand redistribution at the boron center,26 we probed their ability to transfer a hydride from their respective borohydrides computationally. As expected, the replacement of a pentafluorophenoxy group for a pentafluorophenyl group on [HB(C6F5)3]− resulted in a borohydride complex more favorable to transfer a hydride by 12, 18, and 21 kcal/mol than [HB(C6F5)3]− for [HB(OC6F5)(C6F5)2]−, [HB(OC6F5)2(C6F5)]−, and [HB(OC6F5)3]−, respectively. The increase in hydride donor ability is attributed to the partial donation of the lone pair of the aryloxy groups, which reduces the Lewis acidity of the parent borane. A similar phenomenon is observed in the CBS catalysts,56 where addition of a Lewis acid is needed to turn on the Lewis acidity of the oxazaborolidine (see below). Substitution of the first pentafluorophenoxy group resulted in the largest increase in the ability to transfer a hydride, where each subsequent pentafluorophenoxy group substitution affected the ability to transfer a hydride half as much as the previous substitution. If a phenoxy group is present on the borohydride complex as opposed to a pentafluorophenoxy group in the case of [HB(OPh)(C6F5)2]−, hydride loss becomes more favorable by about 6 kcal/mol than [HB(OC6F5)(C6F5)2]−. If all of the pentafluorophenoxy groups in [HB(OC6F5)3]− are replaced with phenoxy or thiophenoxy groups, hydride loss becomes more favorable by about 13 and 2 kcal/mol, respectively. DuBois and co-workers have previously shown that borate esters such as B(OtBu)3 and B(OSiMe3)3 do not accept a hydride from [HBEt3]− even in the presence of 10 equivalents of [HBEt3]−.16 Computational analysis showed that [HB(OtBu)3]−, [HB(OSiMe3)3]−, and [HB(OH)3]− are 23, 19, and 16 kcal/mol more favorable to transfer a hydride than [HBEt3]−, thus verifying the experimental observations. Since catecholate boranes have been recently explored in FLP chemistry,57 we exploited this motif to probe the effect of fluorine substituents on an arene ring of the catecholate ligand. Computational analysis showed that hydride transfer from [(C6F5)-BH(catecholate)]− (which is 31 kcal/mol more favorable than for [HB(C6F5)3]−) becomes less favorable by 3 kcal/mol by substitution of fluorine in the 3- or 4-positions of the catecholate ligand (see Supporting Information), with substitution in the 3-position yielding a slightly more reducing borohydride. The substitution of two fluorine atoms of [(C6F5)-BH(catecholate)]− exhibited a difference of about 1.5 kcal/mol between substitution at the 3,6-positions versus the 3,4- or 3,5-positions, with substitution at the 3,6-positions yielding the less reducing borohydride. Substitution of three fluorine atoms in the 3,4,5-positions yielded a borohydride as reducing as the substitution of two fluorine atoms in the 3,6positions. Substitution of three fluorine atoms in the 3,5,6positions of [(C6F5)-BH(catecholate)]− resulted in a borohy-

recently shown by O’Hare and co-workers to result in a borane 113% the Lewis acidity of B(C6F5)3.49 Computational analysis of [HB(C6F4-o-C6F5)3]− yielded a hydride donor ability 6 kcal/ mol less likely to transfer than [HB(C6F5)3]−, verifying the experimentally determined Lewis acidity and demonstrating that a pentafluorophenyl group is more electron withdrawing than a fluoride. Oestreich and co-workers recently reported the synthesis of tris(5,6,7,8-tetrafluoronaphthalen-2-yl)borane and experimentally determined a Lewis acidity 98% of B(C6F5)3.50 Computational analysis of the hydride donor ability of the respective hydride suggested a less Lewis acidic borane than observed experimentally, 16 kcal/mol more likely to transfer a hydride than [HB(C6F5)3]−. We propose that tris(5,6,7,8tetrafluoronaphthalen-2-yl)borohydride is a better hydride donor than experimentally determined due to a similar structure and hydride donor ability to [HBPh2(C6F5)]−. Ashley and co-workers have recently described the use of perchlorinated derivatives of B(C6F5)3 in FLP chemistry.51,52 Analysis of the respective hydrides showed that [HB(C6Cl5)3]−, [HB(C6F5)(C6Cl5)2]−, and [HB(C6F5)2(C6Cl5)]− were 5, 4, and 4 kcal/mol, respectively, more favorable to transfer a hydride than [HB(C6F5)3]−. Intriguingly, the ability to transfer a hydride from [HB(C6Cl5)3]−, [HB(C6F5)(C6Cl5)2]−, and [HB(C6F5)2(C6Cl5)]− was nearly identical even though their steric bulk varies by almost 30° (cone angles of 213°, 202°, and 185°, respectively), indicating more of an electronic effect. Marder and co-workers have recently reported the synthesis of air-stable boranes containing fluoromesityl groups.53 Computational analysis showed that the respective borohydride, [HB(C6H4-p-tBu)(2,4,6-tris(trifluoromethyl)phenyl)2]−, resulted in a hydride donor ability about 19 kcal/mol more reducing than [HB(C6F5)3]− and on the order of [BH4]−. Ashley and co-workers recently reported that tris(3,5-bis(trifluoromethyl)phenyl)borane is more Lewis acidic than B(C6F5)3,54 but computational analysis of the respective borohydride [HB(3,5-bis(trifluoromethyl)phenyl)3]− showed that it was 12 kcal/mol more favorable to transfer a hydride than [HB(C6F5)3]−. The increase in the ability to transfer a hydride is attributed to the diminished electron-withdrawing nature of a 3,5-bis(trifluoromethyl)phenyl group when compared to a pentafluorophenyl group, as the respective borohydrides exhibit similar steric bulk.38 Investigation of the electron-withdrawing nature of halogens instead of pentafluorophenyl groups in [HB(C6F5)3]− resulted in the transfer of a hydride from [HBCl3]− or [HBF3]− to be 1 kcal/mol less favorable or 16 more favorable than [HB(C6F5)3]−, respectively. Reducing Ability of Alkyl Borohydrides. Reduction chemistry employing alkyl boranes is primarily governed by reactions with selectride ([HB(sec-butyl)3]−) and superhydride ([HBEt3]−).34,55 Computational analysis of the hydride donor ability of selectride and superhydride showed that selectride was about 4 kcal/mol more favorable to transfer a hydride than superhydride. This observed reactivity is attributed to the increased steric bulk surrounding the borohydride center in selectride.38 To examine the effect of changing from an alkyl group to an arene group, we probed the hydride donor abilities of the respective borohydrides upon replacing ethyl groups with phenyl groups. Replacement of an ethyl group with a phenyl group of [HBEt3]− yielded borohydrides about 2−6 kcal/mol less reducing with each subsequent replacement for [HBEt2Ph]− and [HBEtPh2]−, respectively. Replacement of the three ethyl groups of [HBEt3]− with methyl groups resulted D


Article

Organometallics

Figure 1. List of calculated hydride donor abilities for chiral boranes relevant to the chemistry of frustrated Lewis pairs. In the case of two possible diastereomers, only the dominant diastereomer is shown. The numbers shown are in units of kcal/mol. The hydride ligand that undergoes hydride transfer is bold and colored in red.

pyrrolidine nitrogen center, the reducing capacity of the borohydride decreases by about 19 kcal/mol (Figure 1), resulting in a borohydride with a similar reducing capacity to [HB(Mes)2(C6F5)]−. In the catalytic reductions of ketones with the CBS catalysts, hydride transfer occurs from the bound BH3 center to the substrate as opposed to a hydride that is bound to the boron center in the oxazaborolidine ring.56 Coordination of BH3 to the pyrrolidine ring (3) resulted in the hydride donor ability of the BH3 group of the three-coordinate oxazaborolidine (Figure 1) to have a similar reducing ability to NH3BH3 (see below). Recent developments of chiral boranes have focused on the metal-free reduction of imines using FLPs. The first chiral borane (4) used in asymmetric hydrogenations was developed by Klankermayer and co-workers through the hydroboration of (+)-α-pinene with Piers borane ([HB(C6F5)2]2),59 resulting in enantioselectivities of 13%.60 Computational analysis of these borohydrides showed that the most thermodynamically favorable borohydride was 10 kcal/mol more reducing than [HB(C6F5)3]−, which is of similar magnitude to the related hydroboration reactions (see above). The minor diastereomer (3 kcal/mol less stable than the major diastereomer) exhibited a hydride donor ability 1 kcal/mol more reducing than the borohydride of the major diastereomer, suggesting that a greater energy difference is needed to result in reductions with greater ee. Klankermayer and co-workers then expanded on the previous result with the use of a camphor-derived borane (5), resulting in enantioselectivities of up to 83%.61 Computational analysis showed that the major diastereomer (exo-product) generated in the FLP H2 reactions had a reducing ability about 10 kcal/mol more favorable than [HB(C6F5)3]− and is less susceptible to retrohydroboration. The minor diastereomer (endo-product) was about 6 kcal/mol less stable, and the respective borohydride was 3 kcal/mol less reducing than the major diastereomer’s respective borohydride, suggesting that the higher enantioselectivity is attributed to the energy differences between the two diastereomers as opposed to steric bulk (cone angles of 197° and 200°).38 Several other groups have generated chiral boranes, but have not yet employed them in FLP chemistry. Jäkle and co-workers have synthesized a planar chiral borane generated from

dride about 1.5 kcal/mol less reducing than substitution in the 3,4,5-positions. The substitution of four fluorine atoms on the catecholate ligand results in a borohydride that is 11 kcal/mol less likely to transfer a hydride than the parent catecholate complex, [(C6F5)-BH(catecholate)]−. Chiral Boranes. In asymmetric reductions, chiral boranes are used to promote the interaction of the borohydride with one face of the substrate over another. To achieve enantioselectivities greater than 90% ee, transition-state energies between the approaching hydride and the two faces of the prochiral substrate need to be 1.8, 2.0, and 2.2 kcal/mol for reactions at 25, 60, and 100 °C, respectively. To promote reductions with high enantioselectivities, a combination of steric bulk and a regulated hydride donor is needed in the design of a desired reductant to achieve the required energy difference between the prochiral faces of a substrate. We have recently reported that an elemental hydride exhibiting a cone angle greater than 165° promotes enantioselective reductions of >90% ee.38 Although the steric bulk of an elemental hydride can be used to promote asymmetric reductions, the thermodynamics of hydride transfer also play a part in the resulting enantioselectivity. To investigate the magnitude of the ability of chiral borohydrides to transfer hydrides, we examined the hydride donor ability of nine chiral borohydrides related to the chemistry of FLPs. Chiral oxazaborolidine (CBS) catalysts developed by Corey, Bakshi, and Shibata have been widely used in the asymmetric reduction of ketones to chiral alcohols.56 Given that the Bpentafluorophenyl derivative has been reported and is capable of reducing ketones to alcohols in 88% ee,58 we investigated computationally its reducing capacity versus other main group hydrides. As expected, the parent CBS catalyst (1) is a poor Lewis acid and a strong hydride donor (ΔGH− = 19 kcal/mol), exhibiting a reducing capacity similar to [HB(sec-butyl)3]−, with the opposite diastereomer being 15 kcal/mol more reducing. Although we initially probed the addition of a hydride ligand to the boron center, the reduction of ketones with CBS catalysts is proposed to undergo a mechanism involving BH3.56 To activate the CBS catalysts, BH3 binds to the pyrrolidinyl nitrogen center to increase the Lewis acidity at the metal center. Computational analysis shows that upon coordination of BH3 (2) to the E


Article

Organometallics

Figure 2. List of calculated hydride donor abilities for amine-boranes of borenium cations relevant to the chemistry of frustrated Lewis pairs.

naphthylferrocene (6).62 Computationally, we determined that the generated borohydride is about 17 kcal/mol more reducing than [HB(C6F5)3]−. Computational analysis also showed that upon oxidation of the iron center, the ability to transfer a hydride decreases by about 5 kcal/mol. Zhong and co-workers recently described the synthesis of a phosphine-borane derived from [2.2]paracyclophane,63 and Paradies recently described a chiral borane containing the [2.2]paracyclophane scaffold.64 Given that this is a chiral scaffold, we probed the hydride donor ability of a paracyclophane-derived borohydride (7) in the 4position (Figure 1),64 which resulted in a borohydride complex about 10 kcal/mol more reducing than [HB(C6F5)3]− and similar to [HB(C6F5)2(Ph)]−. Piers and co-workers have previously synthesized 2-B(C6F5)2-1,1′-binaphthyl (8) and 2B(C6F5)2-2′-methyl-1,1′-binaphthyl (9).65 [2-BH(C6F5)2-1,1′binaphthyl]− was determined to be about 12 kcal/mol more reducing than [HB(C6F5)3]− and 1 kcal/mol less reducing than [2-BH(C6F5)2-2′-methyl-1,1′-binaphthyl]−. To further exploit the chemistry of binaphthyl groups, Oestreich and co-workers have recently shown that the axially chiral borane (10) containing one pentafluorophenyl group can be used for the asymmetric hydrosilylation of imines, achieving 33−63% ee.66,67 Computational analysis showed that the thermodynamically favored borohydride was about 19 kcal/mol more reducing than [HB(C6F5)3]− and similar to [HB(C6F5)(Ph)2]−. Boranes Derived from Borenium Cations. In an effort to generate Lewis acidic boranes employing commercially available reagents, focus has moved to the generation of airand moisture-stable borenium cations.68 Crudden and coworkers recently discovered that a DABCO-HB(pinacolate) (11) is more reducing than [HB(C6F5)3]−.69 In agreement with the experimental observations, computational analysis showed that DABCO-HB(pinacolate) is about 23 kcal/mol more likely to lose a hydride than [HB(C6F5)3]−. The related 2,6-lutidineHB(pinacolate) (12) complex exhibited a hydride donor ability about 6 kcal/mol greater than the DABCO adduct. In addition to Crudden, Stephan and Ingleson have employed borenium cations in FLP chemistry. Stephan and co-workers have employed borenium cations derived from N-heterocyclic carbenes and 9-BBN for the catalytic hydrogenation of imines.70 Computational analysis showed that 1,3-bis(isopropyl)imidazol-2-ylidene-H-BBN (13) and the related 1,3-bis(2,6-diisopropylphenyl)imidazol-2-ylidene-H-BBN (14) complexes resulted in boranes that were 18 and 17 kcal/mol, respectively, more likely to give up a hydride than [HB(C6F5)3]−. Ingleson has recently investigated the chemistry

of 2,6-lutidine adducts of BHCl2 (15).29 Computational analysis revealed the hydride donor ability to be 8 kcal/mol less favorable than [HB(C6F5)3]− and 7 kcal/mol less favorable than [HBCl3]−. In an attempt to examine the effect of Lewis basicity on the hydride donor ability of catecholborane, we determined the hydride donor ability of the NH3 (16), DABCO (17), NEt3 (18), P t Bu 3 (19), 2,6-lutidine (20), 1,3-bis(2,6diisopropylphenyl)imidazol-2-ylidene (21 (R = Dipp)), and 1,3-bis(isopropyl)imidazol-2-ylidene (22 (R = iPr)) Lewis base adducts. The hydride donor abilities ranged 15 kcal/mol (Figure 2), with the weakest hydride donor being the DABCO adduct and the strongest hydride donor being the 1,3bis(isopropyl)imidazol-2-ylidene adduct. No clear correlation between Lewis basicity (pKa in MeCN)71,72 and hydride donor ability of the Lewis base-catecholborane adduct could be determined (Figure 2). We have previously shown that about half of the cone angle of the Lewis base contributes to the cone angle of the Lewis base-borane adduct,38 but no clear correlation between steric bulk and hydride donor ability could be determined. The absence of a correlation indicates that a combination of steric and electronic effects needs to be taken into account when tuning the hydride donor ability of a Lewis base-borane adduct. With recent interest in the use of ammonia-borane as a possible hydrogen storage material,73 we included the examination of hydride transfer chemistry of NH3BH3 in this study. Hydride transfer from the neutral NH3BH3 was 8 kcal/ mol less favorable to transfer a hydride than [HB(C6F5)3]−. If deprotonation of the amine center occurs prior to hydride transfer, the ability to transfer a hydride from [NH2BH3]− becomes more favorable than NH3BH3 by about 60 kcal/mol and is as reducing as [HB(sec-butyl)3]−. Aluminum Hydrides. Although most of the synthesis of new Lewis acids for FLP chemistry has focused around the implementation of boron, several examples of Lewis acids employing aluminum have been reported. Stephan and coworkers have recently described FLP chemistry with the aluminum analogue of B(C6F5)3.74,75 Computational analysis showed that the hydride donor ability of [HAl(C6F5)3]− was 6 kcal/mol more reducing than [HB(C6F5)3]−. A similar observation was observed by Timoshkin and Frenking, where they found that B(C6F5)3 was a stronger Lewis acid than Al(C6F5)3 for the addition of a hydride.76 The related phenyl complex [HAlPh3]− was determined to have a hydride donor ability 20 kcal/mol less than [HAl(C6F5)3]− and about 3 kcal/ F


Article

Organometallics mol less likely to transfer a hydride than [HBPh3]−. Although [HAl(C 6 F 5 ) 3 ] − appears to be more reducing than [HB(C6F5)3]−, the aluminum hydride experimentally exists with two Lewis acids coordinated to the hydride.75 As expected, the coordination of the second aluminum complex decreased the ability to transfer a hydride by 27 kcal/mol, indicating a stabilization by the addition of a second Lewis acid. A similar observation was perceived in the determination of the hydride donor abilities of [HAlCl3]−, [HAlBr3]−, and [HAlI3]−, where the ability to transfer a hydride became less likely by 18, 13, and 17 kcal/mol, respectively, upon addition of the second Lewis acid. Krossing and co-workers recently described the synthesis of a strong Lewis acid, Al(OC(CF3)3)3, for the protonation of mesitylene in the presence of an alcohol.77,78 Computational analysis of the respective hydride resulted in an aluminum hydride species that was 3 kcal/mol less favorable to transfer a hydride than [HB(C6F5)3]−; similar results were obtained by Krossing and co-workers.19 Phosphorus Cations. Interest in the development of Lewis acids that exhibit greater air and moisture stability have led Stephan and co-workers to recently demonstrate the dehydrofluorination of fluoroalkanes and the hydrosilylation of olefins with phosphorus(V) cations.79,80 The proposed catalytic cycles consist of the formation of HFP(C6F5)3. Computational analysis of HFP(C6F5)3 resulted in a hydride species that was 14 kcal/mol less reducing than [HB(C6F5)3]− for the phosphorus(V) species with HF in the cis-configuration; the trans-configuration was 1 kcal/mol less stable and more reducing. The reduced ability to transfer a hydride is on par with the experimental observations that [FP(C6F5)3]+ is a stronger Lewis acid than B(C6F5)3.80 The related cis-HFPPh3 derivative was 35 kcal/mol more reducing than the pentafluorophenyl derivative. Tetrel Hydride Donors. To explore the scope of the hydride donor ability of main group hydrides beyond triel and pnictogen hydrides, we looked to examine tetrel hydrides. Silanes have often been considered analogues of H2,81 and triaryl silylium ions have been recently shown to heterolytically cleave H2 in the presence of trimesitylphosphine.82 The most reducing silane analyzed in this study was HSi(C6Me5)3, which was computed to be about 9 kcal/mol less reducing than [HB(C6F5)3]−. The increased ability to transfer a hydride of HSi(C6Me5)3 over other silanes is attributed to an increased steric bulk around the silicon center.38 The more commonly used silanes for hydrosilylation reactions, HSi(iPr)3, Me2SiHPh, and Et3SiH,83−85 resulted in hydride donor abilities about 20 kcal/mol less reducing than [HB(C6F5)3]−. Intriguingly, the organosilanes commonly used in hydrosilylation reactions were calculated to be about 10 kcal/mol less reactive than H2, suggesting their propensity to model sigma complexes of H2.81,86 Substitution of a phenyl group in HSiPh3 with a hydride decreased the ability of the organosilane to transfer a hydride by 7, 4, and 13 kcal/mol for Ph2SiH2, PhSiH3, and SiH4, respectively. Replacement of a phenyl group in Me2SiHPh with a chloride yielded a silane 13 kcal/mol less reducing, which is expected with the introduction of a greater electron-withdrawing substituent. Piers and co-workers recently reported the formation of a Et3SiH-borole adduct.87 Computational analysis of hydride transfer from 4,5,6,7-tetrafluoro-1,2,3trispentafluorophenyl-1H-benzo[b]borol-1-uide resulted in a hydride that was 4 kcal/mol less likely to transfer than [HB(C6F5)3]− and 17 kcal/mol less likely to transfer than

Et3SiH, suggesting that the increase in Lewis acidity favors adduct formation as opposed to hydride transfer. Hantzsch esters have been recently utilized in asymmetric transfer hydrogenation reactions using chiral Brønsted acids88,89 and have been shown to transfer a hydride to B(C6F5)3.90 Calculation of the hydride donor ability of the neutral Hantzsch ester showed that it is more likely to give up a hydride than [HB(C6F5)3]− by 6 kcal/mol and has a similar hydride donor ability to [HAl(C6F5)3]−. Deprotonation of the nitrogen center leads to activation of the hydride, with the hydride complex becoming 33 kcal/mol more reducing. Whereas a Hantzsch ester exhibited a hydride donor ability similar to [HAl(C6F5)3]−, another organic hydrogen source, formate, displayed a similar hydride donor ability to [HAlPh3]−. Some of the interest in the reduction chemistry of Hantzsch esters is related to their structural similarity to the biological hydride transfer agent NADH. A hydride source related to NADH is 5,10-methylenetetrahydromethanopterin, which in combination with a monometallic iron cofactor activates hydrogen in the Hmd hydrogenase.91 The hydride donor ability of 5,10-methylenetetrahydromethanopterin was found to have almost an identical hydride donor ability to [BH4]−. [9-H10-Me-acridinium]+ has been recently employed as a carbonbased Lewis acid for the promotion of the hydrosilylation of imines and heterolytic cleavage of H2.83 Computational analysis showed 10-methylacridane to be about 3 kcal/mol less reducing than [HB(C6F5)3]−. To examine the trend of hydride donor abilities upon moving down the periodic table, we investigated the H(E)Ph3 motif, where E = C, Si, Ge, and Sn. Trityl cation was found to be the most Lewis acidic tetrel cation, with triphenylmethane being about 6 kcal/mol more stable than triphenylsilane toward the loss of a hydride. Determination of the hydride donor ability of triphenylgermane and triphenylstannane resulted in a hydride complex more favorable to transfer a hydride than the previous tetrel by about 5 kcal/mol. The observed trend suggests that the Lewis acidity of the tetrels decreases as one moves down the periodic table.

■

CONCLUSIONS The quantification of the reducing ability of main group hydrides is critical to catalyst design. To give insight into the reducing capacity of main group compounds that are widely used in reduction chemistry, a computational study was conducted. We described the hydride donor abilities of 132 main group hydrides consisting of borohydrides, boranes, organosilanes, organostannanes, and carbon-based hydrides spanning a range of about 160 kcal/mol (Figure 3, Supporting Information). The commonly used Lewis acid, B(C6F5)3, in FLP chemistry yielded a borohydride 15 kcal/mol less reducing than [BH4]−, 22 kcal/mol less reducing than [AlH4]−, and 41 kcal/mol less reducing than superhydride ([HBEt3]−). Analysis of periodic trends showed that as one goes down the tetrels, the Lewis acidity decreases and the ability to transfer a hydride increases by 5 kcal/mol. Further analysis of the hydride donor abilities of organosilanes resulted in hydrides that are less reducing than dihydrogen by about 10 kcal/mol, thus verifying their use as hydrogen analogues and the propensity to form sigma complexes with metals. Investigation of chiral borohydrides, related to the chemistry of FLPs, showed that most of the chiral borohydrides exhibited a reducing capacity similar to [BH4]−. Further scrutiny of the hydride donor abilities of chiral borohydrides suggested that the less stable diastereomer results G


Article

Organometallics

description of main group compounds and the way it accounts well for the dispersion interactions of weakly bound complexes.27,28,92 In all calculations, the ultrafine integration grid was employed to ensure the stability of the optimization procedure for the molecules of interest. Each stationary point was confirmed by a frequency calculation at the same level of theory to be a real local minimum on the potential energy surface without an imaginary frequency to be a minimum structure or through a potential energy surface scan in the case of [HB(3,5-bis(trifluoromethyl)phenyl)3]−, [DABCO-B(catecholate)]+, and [n-Bu3Sn]+, where the optimized structures consistently exhibited one imaginary frequency.28 The 6-311++G(d,p) basis set was used to compute more accurate electronic energies for the optimized geometries.97,98 All reported free energies are for acetonitrile solutions at the standard state (T = 298 K, P = 1 atm, 1 mol/L concentration of all species in MeCN) as modeled by a polarized continuum model.99 The energy values given here correspond to solvent-corrected Gibbs free energies that are based on the M06-2X/6-311++G(d,p) electronic energies and all corrections calculated at the M06-2X/6-31G(d) level. For structures with greater than one possible conformer, each possible conformer (ranging from two to six structures) was optimized independently, and only the lowest energy conformer is reported within. An estimated error for the computed hydride donor abilities is 3 kcal/mol.

■

ASSOCIATED CONTENT

* Supporting Information S

3D coordinates and energies of all calculated structures, table of computed and experimentally determined Lewis acidities, tables of free energies and enthalpies for the hydride donor ability of all computed main group hydrides. Complete ladder diagram showing the hydride donor abilities of the main group hydrides described within compared to known metal complexes. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/om5011512.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■ ■

ACKNOWLEDGMENTS This research was supported with startup funds from Washington State University.

in a more reducing borohydride, and the larger energy difference between the two borane diastereomers resulted in the highest enantiomeric excess in catalytic hydrogenation reactions. The investigation and implementation of main group elements in reduction chemistry is a topic of interest in our laboratory.

■

REFERENCES

(1) Sodium Borohydride; Dow Chemical Company; accessed September 11, 2014. http://www.dow.com/sbh/products/sodiumborohydride.htm. (2) Reductions. In Ullmann’s Encyclopedia of Industrial Chemistry, 7th ed.; John Wiley & Sons, Inc., 2010. (3) Stephan, D. W.; Erker, G. Angew. Chem., Int. Ed. 2010, 49, 46−76. (4) USP Elemental Impurities - Limits [online early access]. Published online: 2014. http://www.usp.org/sites/default/files/usp_ pdf/EN/USPNF/key-issues/c232_final.pdf. (5) Thomson, J. W.; Hatnean, J. A.; Hastie, J. J.; Pasternak, A.; Stephan, D. W.; Chase, P. A. Org. Process Res. Dev. 2013, 17, 1287− 1292. (6) Stephan, D. W.; Greenberg, S.; Graham, T. W.; Chase, P.; Hastie, J. J.; Geier, S. J.; Farrell, J. M.; Brown, C. C.; Heiden, Z. M.; Welch, G. C.; Ullrich, M. Inorg. Chem. 2011, 50, 12338−12348. (7) Stephan, D. W.; Erker, G. Frustrated Lewis Pair Mediated Hydrogenations. In Frustrated Lewis Pairs I; Erker, G., Stephan, D. W., Eds.; Springer: Berlin, 2013; Vol. 332, pp 85−110. (8) Mahdi, T.; Stephan, D. W. J. Am. Chem. Soc. 2014, 136, 15809− 15812.

Figure 3. Calculated hydride donor abilities for selected main group hydrides in acetonitrile at the M06-2X/6-31G(d)//M06-2X/6-311G+ +(d,p) level of theory. Please see the Supporting Information for a more complete ladder diagram.

EXPERIMENTAL SECTION

Computational Methods. All structures were fully optimized without symmetry constraints using the M06-2X92 functional as implemented in Gaussian 09,93 using the 6-31G** basis set.94,95 The Stuttgart basis set with effective core potentials was employed for all tin, germanium, and iodine atoms.96 Exchange−correlation functional M06-2X was chosen for the good overall performance observed in the H


Article

Organometallics (9) Scott, D. J.; Fuchter, M. J.; Ashley, A. E. J. Am. Chem. Soc. 2014, 136, 15813−15816. (10) Curtis, C. J.; Miedaner, A.; Ellis, W. W.; DuBois, D. L. J. Am. Chem. Soc. 2002, 124, 1918−1925. (11) DuBois, M. R.; DuBois, D. L. Chem. Soc. Rev. 2009, 38, 62. (12) Cummings, S. A.; Iimura, M.; Harlan, C. J.; Kwaan, R. J.; Trieu, I. V.; Norton, J. R.; Bridgewater, B. M.; Jäkle, F.; Sundararaman, A.; Tilset, M. Organometallics 2006, 25, 1565−1568. (13) Kwaan, R. J.; Harlan, C. J.; Norton, J. R. Organometallics 2001, 20, 3818−3820. (14) Lawrence, E. J.; Oganesyan, V. S.; Wildgoose, G. G.; Ashley, A. E. Dalton Trans. 2013, 42, 782−789. (15) DuBois, D. L.; Blake, D. M.; Miedaner, A.; Curtis, C. J.; DuBois, M. R.; Franz, J. A.; Linehan, J. C. Organometallics 2006, 25, 4414− 4419. (16) Mock, M. T.; Potter, R. G.; Camaioni, D. M.; Li, J.; Dougherty, W. G.; Kassel, W. S.; Twamley, B.; DuBois, D. L. J. Am. Chem. Soc. 2009, 131, 14454−14465. (17) Rokob, T. A.; Hamza, A.; Papai, I. J. Am. Chem. Soc. 2009, 131, 10701−10710. (18) Rokob, T. A.; Papai, I. Top. Curr. Chem. 2013, 332, 157−212. (19) Bohrer, H.; Trapp, N.; Himmel, D.; Schleep, M.; Krossing, I. Dalton Trans. 2015, 44, 7489−7499. (20) Durfey, B. L.; Gilbert, T. M. Inorg. Chem. 2011, 50, 7871−7879. (21) Morgan, M. M.; Marwitz, A. J. V.; Piers, W. E.; Parvez, M. Organometallics 2013, 32, 317−322. (22) Childs, R. F.; Mulholland, D. L.; Nixon, A. Can. J. Chem. 1982, 60, 809−812. (23) Gutmann, V. Coord. Chem. Rev. 1976, 18, 225−255. (24) Mayer, U.; Gutmann, V.; Gerger, W. Monatsh. Chem. 1975, 106, 1235−1257. (25) Beckett, M. A.; Brassington, D. S.; Coles, S. J.; Hursthouse, M. B. Inorg. Chem. Commun. 2000, 3, 530−533. (26) Neu, R. C.; Ouyang, E. Y.; Geier, S. J.; Zhao, X.; Ramos, A.; Stephan, D. W. Dalton Trans. 2010, 39, 4285−4294. (27) Zhao, Y.; Truhlar, D. G. Theor. Chem. Acc. 2008, 120, 215−241. (28) Goerigk, L.; Grimme, S. Phys. Chem. Chem. Phys. 2011, 13, 6670−6688. (29) Clark, E. R.; Del Grosso, A.; Ingleson, M. J. Chem.Eur. J. 2013, 19, 2462−2466. (30) Nimlos, M. R.; Chang, C. H.; Curtis, C. J.; Miedaner, A.; Pilath, H. M.; DuBois, D. L. Organometallics 2008, 27, 2715−2722. (31) Raugei, S.; DuBois, D. L.; Rousseau, R.; Chen, S.; Ho, M.-H.; Bullock, R. M.; Dupuis, M. Acc. Chem. Res. 2015, 48, 248−255. (32) Kovacs, G.; Papai, I. Organometallics 2006, 25, 820−825. (33) Paquette, L. A.; Ollevier, T.; Desyroy, V. Lithium Aluminum Hydride. In e-EROS Encyclopedia of Reagents for Organic Synthesis; John Wiley & Sons, Ltd, 2001. (34) Zaidlewicz, M.; Brown, H. C. Lithium Triethylborohydride. In eEROS Encyclopedia of Reagents for Organic Synthesis; John Wiley & Sons, Ltd, 2001. (35) Adamczyk-Woźniak, A.; Jakubczyk, M.; Jankowski, P.; Sporzyński, A.; Urbański, P. M. J. Phys. Org. Chem. 2013, 26, 415−419. (36) Abdel-Magid, A. F.; Carson, K. G.; Harris, B. D.; Maryanoff, C. A.; Shah, R. D. J. Org. Chem. 1996, 61, 3849−3862. (37) Kehr, G.; Erker, G. Chem. Commun. 2012, 48, 1839−1850. (38) Lathem, A. P.; Treich, N. R.; Heiden, Z. M. Isr. J. Chem. 2015, 55, 226−234. (39) Yoon, N. M.; Kim, K. E.; Kang, J. J. Org. Chem. 1986, 51, 226− 229. (40) Yoon, N. M.; Kim, K. E. J. Org. Chem. 1987, 52, 5564−5570. (41) Eros, G.; Mehdi, H.; Papai, I.; Rokob, T. A.; Kiraly, P.; Tarkanyi, G.; Soos, T. Angew. Chem., Int. Ed. 2010, 49, 6559−6563. (42) Neu, R. C.; Stephan, D. W. Organometallics 2012, 31, 46−49. (43) Brown, H. C. Hydroboration; W. A. Benjamin, Inc.: New York, NY, 1962. (44) Jiang, C.; Blacque, O.; Fox, T.; Berke, H. Dalton Trans. 2011, 40, 1091−1097.

(45) Peuser, I.; Neu, R. C.; Zhao, X.; Ulrich, M.; Schirmer, B.; Tannert, J. A.; Kehr, G.; Fröhlich, R.; Grimme, S.; Erker, G.; Stephan, D. W. Chem.Eur. J. 2011, 17, 9640−9650. (46) Ullrich, M.; Lough, A. J.; Stephan, D. W. Organometallics 2010, 29, 3647−3654. (47) Keess, S.; Simonneau, A.; Oestreich, M. Organometallics 2015, 34, 790−799. (48) Greb, L.; Daniliuc, C.-G.; Bergander, K.; Paradies, J. Angew. Chem., Int. Ed. 2013, 52, 5876−5879. (49) Binding, S. C.; Zaher, H.; Mark Chadwick, F.; O’Hare, D. Dalton Trans. 2012, 41, 9061−9066. (50) Mohr, J.; Durmaz, M.; Irran, E.; Oestreich, M. Organometallics 2014, 33, 1108−1111. (51) Ashley, A. E.; Herrington, T. J.; Wildgoose, G. G.; Zaher, H.; Thompson, A. L.; Rees, N. H.; Krämer, T.; O’Hare, D. J. Am. Chem. Soc. 2011, 133, 14727−14740. (52) Zaher, H.; Ashley, A. E.; Irwin, M.; Thompson, A. L.; Gutmann, M. J.; Kraemer, T.; O’Hare, D. Chem. Commun. 2013, 49, 9755−9757. (53) Zhang, Z.; Edkins, R. M.; Nitsch, J.; Fucke, K.; Steffen, A.; Longobardi, L. E.; Stephan, D. W.; Lambert, C.; Marder, T. B. Chem. Sci. 2015, 6, 308−321. (54) Herrington, T. J.; Thom, A. J. W.; White, A. J. P.; Ashley, A. E. Dalton Trans. 2012, 41, 9019−9022. (55) Hubbard, J. L.; Dake, G. Lithium Tri-sec-butylborohydride. In eEROS Encyclopedia of Reagents for Organic Synthesis; John Wiley & Sons, Ltd, 2001. (56) Corey, E. J.; Helal, C. J. Angew. Chem., Int. Ed. 1998, 37, 1986− 2012. (57) Dureen, M. A.; Lough, A.; Gilbert, T. M.; Stephan, D. W. Chem. Commun. 2008, 4303−4305. (58) Korenaga, T.; Kobayashi, F.; Nomura, K.; Nagao, S.; Sakai, T. J. Fluorine Chem. 2007, 128, 1153−1157. (59) Parks, D. J.; von H. Spence, R. E.; Piers, W. E. Angew. Chem., Int. Ed. Engl. 1995, 34, 809−811. (60) Chen, D.; Klankermayer, J. Chem. Commun. 2008, 2130−2131. (61) Chen, D.; Wang, Y.; Klankermayer, J. Angew. Chem., Int. Ed. 2010, 49, 9475−9478. (62) Chen, J.; Venkatasubbaiah, K.; Pakkirisamy, T.; Doshi, A.; Yusupov, A.; Patel, Y.; Lalancette, R. A.; Jakle, F. Chem.Eur. J. 2010, 16, 8861−8867. (63) Wang, G.; Chen, C.; Du, T.; Zhong, W. Adv. Synth. Catal. 2014, 356, 1747−1752. (64) Greb, L.; Paradies, J. Top. Curr. Chem. 2013, 334, 81−100. (65) Morrison, D. J.; Piers, W. E.; Parvez, M. Synlett 2004, 2429− 2433. (66) Hermeke, J.; Mewald, M.; Oestreich, M. J. Am. Chem. Soc. 2013, 135, 17537−17546. (67) Mewald, M.; Fröhlich, R.; Oestreich, M. Chem.Eur. J. 2011, 17, 9406−9414. (68) Piers, W. E.; Bourke, S. C.; Conroy, K. D. Angew. Chem., Int. Ed. 2005, 44, 5016−5036. (69) Eisenberger, P.; Bailey, A. M.; Crudden, C. M. J. Am. Chem. Soc. 2012, 134, 17384−17387. (70) Farrell, J. M.; Hatnean, J. A.; Stephan, D. W. J. Am. Chem. Soc. 2012, 134, 15728−15731. (71) Izutsu, K. Acid-Base Dissociation Constants in Dipolar Aprotic Solvents; Blackwell Scientific Publications: Oxford, 1990; Vol. 35. (72) Kaljurand, I.; Kuett, A.; Soovaeli, L.; Rodima, T.; Maeemets, V.; Leito, I.; Koppel, I. A. J. Org. Chem. 2005, 70, 1019−1028. (73) Stephens, F. H.; Pons, V.; Tom Baker, R. Dalton Trans. 2007, 2613−2626. (74) Menard, G.; Tran, L.; Stephan, D. W. Dalton Trans. 2013, 42, 13685−13691. (75) Menard, G.; Stephan, D. W. Angew. Chem., Int. Ed. 2012, 51, 8272−8275. (76) Timoshkin, A. Y.; Frenking, G. Organometallics 2008, 27, 371− 380. (77) Kraft, A.; Beck, J.; Steinfeld, G.; Scherer, H.; Himmel, D.; Krossing, I. Organometallics 2012, 31, 7485−7491. I


Article

Organometallics (78) Kraft, A.; Trapp, N.; Himmel, D.; Böhrer, H.; Schlüter, P.; Scherer, H.; Krossing, I. Chem.Eur. J. 2012, 18, 9371−9380. (79) Pérez, M.; Hounjet, L. J.; Caputo, C. B.; Dobrovetsky, R.; Stephan, D. W. J. Am. Chem. Soc. 2013, 135, 18308−18310. (80) Caputo, C. B.; Hounjet, L. J.; Dobrovetsky, R.; Stephan, D. W. Science 2013, 341, 1374−1377. (81) Kubas, G. J. Adv. Inorg. Chem. 2004, 56, 127−177. (82) Schäfer, A.; Reißmann, M.; Schäfer, A.; Saak, W.; Haase, D.; Müller, T. Angew. Chem., Int. Ed. 2011, 50, 12636−12638. (83) Clark, E. R.; Ingleson, M. J. Angew. Chem., Int. Ed. 2014, 53, 11306−11309. (84) Parks, D. J.; Piers, W. E. J. Am. Chem. Soc. 1996, 118, 9440− 9441. (85) Gutsulyak, D. V.; Vyboishchikov, S. F.; Nikonov, G. I. J. Am. Chem. Soc. 2010, 132, 5950−5951. (86) Kubas, G. Metal Dihydrogen and σ-Bond Complexes; Kluwer Academic/Plenum: New York, 2001. (87) Houghton, A. Y.; Hurmalainen, J.; Mansikkamäki, A.; Piers, W. E.; Tuononen, H. M. Nat. Chem. 2014, 6, 983−988. (88) Ouellet, S. G.; Walji, A. M.; Macmillan, D. W. C. Acc. Chem. Res. 2007, 40, 1327−1339. (89) You, S.-L. Chem. Asian J. 2007, 2, 820−827. (90) Webb, J. D.; Laberge, V. S.; Geier, S. J.; Stephan, D. W.; Crudden, C. M. Chem.Eur. J. 2010, 16, 4895−4902. (91) Shima, S.; Pilak, O.; Vogt, S.; Schick, M.; Stagni, M. S.; MeyerKlaucke, W.; Warkentin, E.; Thauer, R. K.; Ermler, U. Science 2008, 321, 572−575. (92) Zhao, Y.; Truhlar, D. G. J. Chem. Theory Comput. 2008, 4, 1849−1868. (93) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Keith, T.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, O.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09, Revision D.01; Gaussian, Inc.: Wallingford, CT, 2013. (94) Dill, J. D.; Pople, J. A. J. Chem. Phys. 1975, 62, 2921−2923. (95) Hehre, W. J.; Ditchfield, R.; Pople, J. A. J. Chem. Phys. 1972, 56, 2257−2261. (96) Andrae, D.; Häußermann, U.; Dolg, M.; Stoll, H.; Preuß, H. Theoret. Chim. Acta 1990, 77, 123−141. (97) Krishnan, R.; Binkley, J. S.; Seeger, R.; Pople, J. A. J. Chem. Phys. 1980, 72, 650−654. (98) Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.; Gordon, M. S.; DeFrees, D. J.; Pople, J. A. J. Chem. Phys. 1982, 77, 3654−3665. (99) Cossi, M.; Rega, N.; Scalmani, G.; Barone, V. J. Comput. Chem. 2003, 24, 669−681.

J


Establishing the Hydride Donor Abilities of Main Group Hydrides

Recommend Documents