Intermolecular Interactions of a Chiral Amine Borane Adduct Revealed

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Intermolecular Interactions of a Chiral Amine Borane Adduct Revealed by VCD Spectroscopy Tobias Osowski, Julia Golbek, Klaus Merz, and Christian Merten J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.6b03955 • Publication Date (Web): 02 Jun 2016 Downloaded from http://pubs.acs.org on June 7, 2016

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The Journal of Physical Chemistry

Intermolecular Interactions of a Chiral Amine Borane Adduct Revealed by VCD Spectroscopy Tobias Osowskia, Julia Golbekb, Klaus Merzb, Christian Mertena,* a)

Ruhr-Universität Bochum, Organische Chemie 2, 44801 Bochum, Germany

b)

Ruhr-Universität Bochum, Anorganische Chemie 1, 44801 Bochum, Germany

* corresponding author: [email protected]; +49 234 32-24529

ABSTRACT Amine boranes feature strong hydrogen bonding acceptor and donor moieties in close proximity, leading, for instance, to dihydrogen bonding driven self-aggregation. In this work, the IR and VCD spectra of the bulky bis(-phenylethyl)amine borane 1 in chloroform and acetonitrile solution are reported. By comparison with calculated spectra, the VCD spectral features observed in chloroform solution can clearly be associated with the presence of monomeric species. A shift of the conformational preferences occurs when changing the solvent to acetonitrile, which can only be deduced from the VCD spectral signatures but not from the IR spectrum. Using variable-temperature IR and VCD spectroscopy, the dihydrogen bonded dimeric species is characterized experimentally at -50 °C and theoretically by means of DFT calculations.

1 ACS Paragon Plus Environment


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INTRODUCTION Amine boranes are well known as hydroborating and reducing agents,1 but more recently, their dehydrocoupling became extensively explored. In context of amine boranes, dehydrocoupling describes the simultaneous release of H2 and the formation of a new B-N bond. It has attracted significant interest due to the possible use of the resulting materials in the design of new B-N polymers such as poly(aminoboranes) ([RNH-BH2]n)2 or as hydrogen storage materials.3 Dehydrocoupling (or dehydrogenation) can be achieved either thermally, or by using organometallic4-6 or metal-free7 catalysts at reduced to ambient temperatures.

The key structural motif allowing for an efficient release of H2 is the simultaneous presence of a partially positively charged hydrogen (N-H) close to a hydridic one (B-H). The coordination of the hydride to free coordination sites of metal centers is often considered to be an important step in catalytic mechanisms,8 and the driving force for dehydrogenation.9 The hydride can also act as acceptor in a very subtle type of hydrogen bond, the so-called dihydrogen bond (DHB). DHBs were first discovered as short intermolecular distances in metal hydrides as NHH-M and O-HH-M interactions.10 A later comprehensive survey of the Cambridge Crystal Structure Database (CCSD) on N-HH-B distances established a general set of geometric parameters.11 Accordingly, a DHB is classified by an average HH distance of 1.96 Å (ranging from 1.7-2.2 Å), an N-HH angle of 149° (117-171°), and an average B-HH angle of 120° (90171°).12


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The fact that both the hydrogen bond donor and acceptor are hydrogen atoms makes 1H-NMR and IR spectroscopy (in the N-H and B-H stretching regions) attractive methods to study the intermolecular interactions of amine boranes in solution.3, 12 However, (di-)hydrogen bonds are chiefly characterized in an indirect approach by intensity changes of vibrational bands,13 or by shifts of proton signals in NMR spectra.12, 14 With these techniques, conformational preferences of more complex amine boranes often cannot be resolved as the spectral changes are not distinct enough.

Recently, we showed that introducing chirality to the amine moiety of an amine borane adduct makes this class of compounds accessible by vibrational circular dichroism (VCD) spectroscopy.15 VCD spectroscopy is a powerful tool to characterize chiral molecular structures in solution as it shows high sensitivity for conformational changes and intra- and intermolecular interactions of chiral molecules.16-17 Besides the determination of absolute configurations,18-21 VCD spectroscopy is increasingly used for structural studies on small molecules, supramolecular assemblies,22-23 and macromolecules.24-26 Intermolecular interactions ranging from solute-solvent interactions27-30 and self-aggregation31-32 to chirality induction phenomena33-35 have been proven using VCD spectroscopy. Taking advantage of this high sensitivity to intermolecular interactions of chiral molecules, our study revealed detailed structural insight into the self-aggregation of chiral -methylbenzyl amine borane (MBA-BH3) through dihydrogen bonds. The dihydrogen bonded dimer could be identified by unique VCD spectral features.15

Due to the strong self-aggregation of the primary amine borane, it was not possible to characterize the monomeric species in the accessible concentration range. In order to further



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explore the potential to use VCD spectroscopy for studies on intermolecular interactions of amine boranes, we herein introduce the chiral amine borane 1 which is based on the secondary amine bis(-phenylethyl)amine (cf. Scheme 1). We show that the steric bulk of a second phenylethyl moiety efficiently prevents dimerization at room temperature, and that the dihydrogen bonding driven self-aggregation can be controlled by variation of the temperature. The VCD spectra also reveal interesting insights into solute-solvent interactions.

Scheme 1 Structure of bis{(S)--phenylethyl}amine borane, (S)-1.

EXPERIMENTAL DETAILS Preparation of amine borane 1. In a dried flask, 300 µL of enantiopure amine were dissolved in 15 ml n-pentane and DMSBH3 (1.1 eq) dissolved in 5 ml n-pentane was slowly added to the solution. During the addition, a white solid began to precipitate, which was collected by filtration after 30 minutes of stirring at room temperature, subsequently washed multiple times with npentane, and dried in vacuo. Isolated yields were generally above 90%. 1H{11B} (C6D6, 250 MHz) δ: 7.05 - 6.94 (8H, m, ArH), 6.65 (2H, m, ArH), 4.08 (1H, m, C*H-NH), 3.74 (1H, s, broad, NH), 3.68 (1H, m, C*H-NH), 2.47 (3H, m, very broad, BH3), 1.54 (3H, d, CH3), 1.39 (3H, d, CH3); 13C (CDCl3, 63 MHz) δ: 141.60 (ArC-C*H-NH), 138.26 (ArC-CH-NH), 129.13 (ArC),


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129.00 (ArC), 128.81 (ArC), 128.46 (ArC), 127.09 (ArC), 62.15 (CH-NH), 57.78 (CH-NH), 19.39 (CH3), 15.75 (CH3); 11B (CDCl3, 80MHz) δ: -19.52 ppm (q, J = 91.2 Hz). Computational details. Geometry optimizations and IR and VCD spectra calculations were performed for the (S)-enantiomers in the DFT framework using Gaussian 09 D.0136 with tight SCF convergence and ultrafine integration grids. Three functionals, namely B3LYP, B3LYPGD3, and B3PW91, were evaluated in combination with the 6-311++G(2d,p) basis set. The results shown in the main text are those from the B3LYP calculations, while the other results are provided in the supporting information. Solvent effects were taken into account implicitly by using the integral equation formalism of the polarizable continuum model (IEFPCM)37-38 for chloroform or acetonitrile. Counterpoise-corrected complexation energies cannot be determined in the IEFPCM environment. Therefore, counterpoise calculations were carried out in gas phase on the optimized IEFPCM structures without further optimization. Vibrational line broadening was simulated by assigning a Lorentzian band shape with half-width at half-height of 6 cm-1 to the calculated dipole and rotational strength. Finally, the calculated wavenumbers were scaled by a factor of 0.975 for comparison with the experimental data. IR and VCD spectroscopy. The IR and VCD spectra were recorded on a Bruker Vertex 70V spectrometer equipped with a PMA 50 module for polarization modulated measurements. Both IR and VCD spectra were recorded at 4 cm-1 spectral resolution by accumulating 32 respectively ~20000 scans. Baseline corrections of the VCD spectra were done by subtraction of the spectra of the corresponding racemic mixtures. Solutions of the samples were held in a sealed transmission cell with BaF2 windows and 100 µm path length. The temperature dependent IR and VCD measurements were carried out using a variable temperature cell holder by Specac. For the IR and VCD spectra taken at -50 °C, the concentration was 50 mg/ml at room temperature. In



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order to account for the temperature dependent change of the density of chloroform, a concentration of 54 mg/ml was assumed at -50 °C. Crystallography. X-ray intensity data for (S)-1 was collected on an Agilent Technologies SuperNova diffractometer with an Atlas CCD detector and Cu Kα radiation (λ = 1.5418) from a microfocus X-ray source with multilayer X-ray optics at 123 K. Structure was solved by direct methods, and all non-hydrogen atoms were refined anisotropically on F² (program SHELXTL97, G.M. Sheldrick, University of Göttingen, Göttingen, Germany). All H atoms were visible in difference maps. The hydrogen atoms attached to the boron and nitrogen centers were refined isotropically, while those attached to C atoms were positioned geometrically and refined as riding atoms. Crystallographic crystal data and processing parameters are given in the supporting information.

RESULTS AND DISCUSSION Room-temperature IR and VCD spectra of 1 in chloroform Figure 1 shows the IR spectra of 1 in the NH/CH and BH stretching as well as the fingerprint region and the VCD spectra obtained in the fingerprint region for a 0.21 M solutions of 1 in CDCl3. The IR spectrum features only few strong and clearly resolved bands arising, for instance, from C=C stretching and CH3 deformation vibrations (1500-1400 cm-1), CH deformation vibration (1400-1350 cm-1), and BH3 bending modes (1200-1150 cm-1). In concentration dependent measurements, the fingerprint region remains essentially unchanged (cf. Figure S1). The BH3-stretching modes are observed in the spectral range from 2600-2200 cm-1, and do not show a distinct concentration dependence either. However, it should be noted that, after solvent subtraction, a small band originating from the CD-stretching mode of the solvent


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remains at 2247 cm-1 suggesting a weak hydrogen bond from chloroform to the amine borane (marked with an asterisk in Figure 1A). Solely in the range of the NH stretching modes, changes can be recognized when the concentration is altered. Here, the decrease of the band associated with the stretching vibration of free NH at 3261 cm-1 and subsequent growth of a band at 3203 cm-1 characteristic for hydrogen bonded NH indicates the formation of dimeric (1)2 with increasing concentration. Using the integrals of these two bands, the ratio of monomer to dimer can be estimated by interpolation of the monomer NH stretching mode to an infinite dilution (cf. supporting information for more details). Accordingly, at the concentration of 0.21 M, the monomeric form of 1 is the predominant species with about 89 %. At a concentration of 0.334 M, the concentration closest to saturation which could be measured, the mole fraction of the monomeric form decreases to only 85 %. In the fingerprint region, the two enantiomers yield good-quality mirror imaged VCD spectra which are rich of spectral features (Figure 1B, bottom panel). All bands can be assigned to specific parent IR bands. It is noteworthy that not only the strong IR bands, but also the weak ones in the range from 1350-1200 cm-1 show strong VCD intensity. Upon increasing of the concentration to 0.33 M, no significant changes in the VCD spectrum were observed besides a slight gain in intensity of the lower wavenumber shoulder of VCD band 3.



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Figure 1. (A) Concentration dependence of the NH/CH and BH stretching region of 1 in chloroform and chloroform-d1. The asterisk marked a band assigned to the C-D+H--B hydrogen bond (cf. text); (B) Comparison of experimental IR and VCD spectra of 1 (0.21 M in chloroform-d1) with those calculated for monomeric 1 (G298K).


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Density functional theory (DFT) based spectra calculations were performed in order to assign the IR and VCD spectral features. The conformational analysis of 1 was carried out at the B3LYP/6-311++G(2d,p) level of theory by systematically varying the torsional angles B-N-CCAr and yielded only five conformations (cf. Table 1). Two of them, denoted as c1 and c2 (cf. Figure 2), contribute almost equally to about 97% of the overall population.

Figure 2. Monomeric structures of (S)-1

Table 1. Characteristic geometric parameters (in degree), relative zero-point corrected (EZPC, in kcal/mol) and Gibbs free energies (G298K, in kcal/mol), and corresponding Boltzmann weights at 298 K (popE and popG, given in %) Conf.

B-N-C -CAr

B-N-C’-C’Ar

EZPC

G298K

popE

popG

c1

168.53

60.08

0.02

0.11

47.90

44.45

c2

92.17

-152.64

0.00a

0.00a

49.65

53.85

c3

-84.24

27.77

1.83

2.10

2.28

1.56

c4

82.22

-63.87

3.44

3.63

0.15

0.12

c5

88.8

70.16

4.51

4.56

0.02

0.02

a

Referenced to the ZPC-energy of c1, Eh = -702.418444 hartree, and the Gibbs free energy Gh= 702.464255 hartree



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Based on the populations derived from the relative Gibbs free energies, G298K, the IR and VCD spectra of monomeric 1 were simulated. The resulting spectra are shown alongside the experimental data in Figure 1B. As indicated by the band assignments, the predicted spectra resemble the experimentally observed ones very well, and almost all bands in the experimental IR and VCD spectrum can be correlated with bands in the calculated spectra. As the IR spectra indicated a weak hydrogen bond between the amine borane and CDCl3 (Figure 1A), we briefly considered solute-solvent complexes [1CDCl3] in our calculations. The strongest possible hydrogen bond with the solvent will be formed with the hydridic hydrogens at the boron. Counterpoise-calculations predict a complexation energy of -3.28 kcal/mol for this dihydrogen bond (-6.00 kcal/mol at B3LYP-GD3 level of theory). In comparison with the calculated fingerprint spectra of the free monomeric form, there are no changes in the VCD spectrum which could clearly be resolved (cf. Figure S5). It is noteworthy that the binding of CDCl3 to the BH3 moiety also causes a slight change in the VCD signature of the BH-stretching vibrations, which however mainly originates from small frequency shifts. It also induces a very weak VCD intensity in the CD-stretching vibration of CDCl3. Unfortunately, the spectral intensities in this range of the VCD spectrum were too weak to observe any of these bands.15

Solvent-induced conformational changes in CD3CN For the analysis of the spectra taken in chloroform, the effect of solute-solvent interactions of 1 with dihydrogen bonded chloroform on its VCD spectrum and its conformational preferences could be neglected. As aprotic hydrogen bond accepting solvent, acetonitrile does not feature any acidic groups, which could interact with the BH3 moiety, nor does it usually contain much water


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which could react with the BH3 moiety. Therefore, it seemed to be a suitable solvent to study the effect of hydrogen bonding to the NH group of 1.

Figure 3. Comparison of experimental IR and VCD spectra of 1 (0.21 M in acetonitrile-d3 at 100 µm path length) with those calculated for the solute-solvent structure [1CD3CN].

In the experimental IR spectrum of 0.21 M solution of 1 in CD3CN, there is no more IR band of non-hydrogen bonded NH (3261 cm-1) but a new band appeared at 3216 cm-1 which can be assigned to the NH stretching vibration of NHNCCD3 (Figure S6). The fingerprint region of the experimental IR spectrum (Figure 3, bottom) does not feature any significant changes compared to the spectrum taken in chloroform-d1 and it also agrees reasonably well with the predicted monomeric spectrum (cf. Fig. 1B). This observation could thus indicate that the conformational preferences of 1 are preserved in acetonitrile. On the other hand, significant VCD spectral changes become apparent when comparing the spectra with those taken in CDCl3 at room temperature (Figure 3, top; direct comparison provided in Figure S7). For instance, the



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VCD bands 6, 8, 11, and 14 became much weaker respectively almost vanished, while the region of bands 1-4 shows a clearly resolved strong bisignate couplet.

Figure 4. Optimized structures of [1CD3CN]. No concentration dependent changes were observed in the IR spectra, and the presence of dimers in acetonitrile seemed unlikely. Therefore, we carried out additional calculations assuming explicit solvation. Solute-solvent complexes [1CD3CN] were generated starting from the two lowest energy monomer conformations (Figure 4). Interestingly, while the conformers c1 and c2 were almost equally populated if no explicit solvation or only interaction with CDCl3 was assumed, the conformation c1 becomes clearly favored in the complex [1CD3CN]. In fact, the solvent-complex of c2 is found to be over 4 kcal/mol less stable than the one of c1. This preference probably originates from the better accessibility of the NH proton in the more open c1-conformation. In Figure 3, the IR and VCD spectra corresponding to [1-c1CD3CN] are shown alongside the experimental spectra. The fingerprint region of the calculated IR spectrum is very similar to the experimental but also to the calculated spectra of the chloroform solution. Hence, based on the IR spectra, it is not possible to differentiate between an equal population of both conformers and the sole existence of the solvated [1-c1CD3CN] adduct. In contrast, the calculated VCD spectrum explains all spectral differences observed between the chloroform and acetonitrile


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measurements. Thus, only by considering the VCD spectrum, the predicted shift of conformational preferences can be confirmed. Although the solute solvent adduct is predicted to have a counterpoise-corrected complexation energy of -4.44 kcal/mol, it is worth noting that the binding of acetonitrile to 1 virtually does not affect the fingerprint region of the single-conformer IR and VCD spectra. Solely the NH stretching modes are influenced by the hydrogen bond. Therefore, it can be concluded that binding of acetonitrile indeed only changes the energetic preferences and thus the location of the conformational equilibrium, but it does not cause any significant structural changes to the amine borane conformations themselves.

Dimeric structures of 1 As the VCD spectra could be used to confirm the change in conformational preferences upon changing the solvent, we wanted to elucidate whether the formation of dimeric (1)2 also affects the VCD spectrum. Therefore, based on the two highly populated monomer conformations, we attempted to optimize dimeric structures (1)2. In analogy to ammonia borane, H3N-BH3, and our previous study, we initially assumed that the organic amine borane HRR’N-BH3 would form head-to-tail dimers HT-(1)2.9, 39 In this orientation, however, only two dimers, HT-(c1-c1) and HT-(c1-c2), could be obtained due to the steric demands of the two phenyl ethyl moieties of each monomer unit. In contrast to what has been found for primary amine boranes,15 the symmetric head-to-tail dimer (Figure 5) does not feature bifurcated dihydrogen bonds, but a single interaction between N-H+H-B (H-H distance of 2.10 Å, and an N-HH angle of 155.58°). In comparison to the corresponding monomeric structure, the bonds involved in the DHB are elongated (B-H by 0.003 Å, N-H by 0.004 Å), while the N-B bond is shortened (0.007 Å).



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Furthermore, natural bond orbital (NBO) analysis reveals intermolecular interactions between the (B-H) orbital and the *(N-H) orbital of the respective other monomer in the dimeric structure, as well as characteristic changes in orbital occupations (see Tab. 2). The dimerization energy is calculated to -4.15 kcal/mol per mole monomer.

Figure 5. Dimeric structures of (S)-1. A) calculated structure of a head-to-tail dimer; B) section of the crystal structure showing the helical arrangement of 1 (carbon-bond hydrogen atoms omitted); C) calculated open dimer structure


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Table 2. Comparison of characteristic changes in bond length (given in Å), and bond angles (given in degree) obtained from calculations on 1 in its monomeric (c1) and dimeric HT-(c1-c1) and open-(c1-c1) form, as well as changes in orbital occupations obtained from NBO analysis. For atom numbering, see Figure 3. bond

monomer

HT-dimer (change)

open dimer (change) solid state

B-H1

1.2090

1.2123 (0.0033)

1.2092b) (0.0002)

1.1418

B-H2

1.2062

1.2051 (-0.0011)

1.20703b) (0.0011)

1.1390

B-H3

1.2084

1.2076 (-0.0008)

1.2086b) (0.0002)

1.1573

N-H4

1.0195

1.0237 (0.0042)

1.0232b) (0.0037)

0.9301

B-N

1.6457

1.6391 (-0.0066)

1.6379b) (-0.0078)

1.6367

H1H4

2.1024

2.1064

2.2594

H2H4

2.8789

BH1H4-N

155.58

BH2H4-N

148.56

H1-B-N-H4 162.23

159.864

164.15b)

(B-H1)

1.98435

1.97729 (-0.00706)

1.97902 (-0.00533)

(B-H2)

1.98417

1.98418 (10-5)

1.98409 (8x10-5)

(B-H3)

1.98775

1.98602 (-0.00173)

1.98726 (-0.00049)

*(N-H4)

0.02865

0.04901 (0.02036)

0.04082 (0.01217)

-4.15 kcal mol-1

-2.32 kcal mol-1

Edimera)

2.4587 161.666

149.009 156.977 156.106

a) BSSE corrected complexation energy divided by 2 b) Measured for the BN bond of the DHB-acceptor



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Figure 5B shows a section of the crystal structure of 1 which could be obtained from crystals grown from toluene solution. It features a helical spatial arrangement of monomer units in a NHiHB(i+1)-NH(i+1) HB(i+2) network. This open network does not feature formal DHBs, as the shortest H-H distance was found to be 2.2594 Å, thus slightly longer than twice the van der Waals radius of hydrogen (2.2 Å). However, it indicates that a head-to-tail dimerization might not be as favourable as suggested by its large computed dimerization energy. Hence, for comparison, we also calculated open dimer structures, denoted as open-(1)2. Although more stretched and thus presumable less affect by the steric effects of the -phenyl ethyl groups, again only two stable structures were found based on the c1 and c2 monomer conformation. In contrast to the solid state structure, the calculated open-(c1-c1) structure (Figure 5C) features a true dihydrogen bond with an H-H distance of 2.10 Å and an N-HH angle of 161.67° (further changes in bond lengths and orbital occupation are given in Table 2). As the open dimer features only one hydrogen bond instead of two, the decrease in calculated dimerization energy by a factor of about two to -2.32 kcal/mol is not unexpected. Interestingly, the relative energies of the four dimer structures vary significantly (Table 3). On the one hand, according to the EZPC energies, the head-to-tail dimers are highly preferred and together make up for 84.4 % of the conformational distribution. On the other hand, taking into account the entropic contributions and thus comparing the G298K values, the open dimers are clearly preferred with more than 90 %. At lower temperature, the conformational preferences shift towards the c1-c2 dimer structures of both the HT- and the open-dimer form.


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Table 3. Relative zero-point corrected (EZPC) and Gibbs free energies (G298K), and corresponding Boltzmann weights at 298 K and 223 K (popE and popG, given in %) of the HTand open-dimer forms of 1. Conf.

EZPC

G298K

G223K

popE

popG298K

popG223K

HT(c1-c1)

0.00a

1.41

0.50

59.11

6.25

15.28

HT(c1-c2)

0.50

1.82

0.02

25.32

3.09

34.16

open(c1-c1)

1.12

0.00b

0.50

8.94

67.42

15.15

open(c1-c2)

1.30

0.63

0.00c)

6.63

23.24

35.41

a

Referenced to the ZPC-energy of HT(c1-c1): Eh = -1404.840446 hartree

b

Referenced to the Gibbs free energy of open(c1-c1) at 298 K: Gh= -1404.914872 hartree

c

Referenced to the Gibbs free energy of open(c1-c1) at 223 K: Gh= -1404.914872 hartree

Temperature dependent IR and VCD measurements In order to experimentally verify if 1 can form dimers in chloroform solution, and if so, which dimer topology might be preferred, further experiments were carried out focusing on trying to increase the degree of dimerization. Concentration dependent IR measurements have shown that the solubility of 1 limits the control over the amount of dimers by variation of the concentration. By performing variable temperature (VT)-IR measurements for the 0.21 M solution, however, a significantly larger degree of dimerization could be reached by decreasing the temperatures. In fact, a simultaneous decrease in intensity of the free NH stretching mode and an increase in the intensity of the hydrogen bonded NH groups can be observed (cf. Figure S8). By fitting the experimental spectra and comparison with the previously determined intensity of the free NH



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stretching vibration at infinite dilution, it can be estimated that approximately 36 % of the amine boranes are found in a dimeric form at -50 °C (mole fraction of 21 %). At about -55 °C, 1 began to precipitate. The temperature dependence of the NH stretching vibrations can be used to estimate the formation enthalpy of the dimeric species.40 Using the fitted band intensities, a van’t Hoff plot can be prepared from which the formation energy of one mole dimer can then be estimated experimentally to -2.79 kcal/mol. This accounts for -1.39 kcal/mol per mole monomer involved in the dimer (see SI for details). The self-association enthalpy can also be determined directly from the IR spectra by analysis of str(NH) frequency shifts using Iogansen’s empirical relationship (eq. 1)41-42 Δ

.

(1)

Here, H (in kcal/mol) is determined from  (in cm-1), which is the difference of the frequencies of the free NH stretching vibration (3261 cm-1) and the NH stretching vibration of the dimer (3206 cm-1), i.e.  = (freeNH-HbondedNH). This correlation yields a binding enthalpy of -1.35 kcal/mol, which is in excellent agreement with the association enthalpy obtained from the van’t Hoff plot. In comparison with the predicted complexation energies (cf. Table 2), these experimentally determined binding energies are significantly lower. This suggests that the binding is generally not strong, and that the energetically less stabilized open-dimer structure might be the more favored in solution.


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Figure 6. Comparison of experimental IR and VCD spectra of 1 (0.21 M in CDCl3 at 100 µm path length) taken at 23 °C and -50 °C with those calculated for monomeric 1 and its dimeric species (1)2.

While the NH stretching region of the VT-IR spectrum at -50 °C clearly indicates an increased self-aggregation, the changes in the fingerprint region, although more prominent than in the concentration dependent spectra, are relatively small (cf. Figure 6 bottom). Solely, the IR band 15 at 1082 cm-1 and the shoulder 12 of the BH3 stretching mode at 1198 cm-1 gain some intensity upon cooling, whereas the broad bands 5 and 6 between 1400-1350 cm-1 show a general sharpening. As the comparison with the room temperature spectrum shows, the changes in the VCD spectra measured at -50 °C are also very small (Figure 6, top). Solely a certain broadening



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of the VCD band 3, a slight decrease in intensity of band 6, or the shift of band 11, might be noted. Nevertheless, both the fingerprint regions of both the IR and the VCD spectra do not indicate any really significant structural changes.

Figure 7. (A) Comparison of dimer VCD spectra (1)2 calculated based on various conformational energy differences compared to the spectra of monomeric 1 and the single c1conformer; (B) Comparison of HT-dimer VCD spectra with spectra obtained by simple summation of the respective single conformer spectra

As already noted above, the preferences for dimer structures shift from HT-dimers (EZPC), to open-dimers (G298K), or to the c1-c2 dimers (G223K). Interestingly, although the EZPC and G298K spectra arise from different dimer topologies, the fingerprint regions of the calculated IR and VCD spectra are almost identical, and very similar to the spectrum of monomer


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conformation c1 as well (Figure 7A). In contrast, the spectrum obtained from the G223K populations appears to be very similar to the calculated averaged monomer spectra (cf. Figure 6 and 7A). Hence, the calculations clearly suggest that there would not be any significant observable changes in the fingerprint region spectra at -50 °C, and thus describe the experimental observations reasonably well. The spectra presented in Figure 7B further showcase the insensitivity of the fingerprint region for the intermolecular interactions of 1. The comparison of the HT-(c1-c1) dimer and monomeric c1 reveals that the hydrogen bonding interaction between two molecules of 1 does not influence the VCD spectral response in the fingerprint region. The HT-(c1-c2) dimer spectrum can also be approximated by simply adding the single conformer spectra of c1 and c2 in a one-to-one ratio. Clearly, both solvation and dimerization solely affect the NH and BH stretching frequencies, but not so much their deformation modes. Furthermore, the influence on the conformations in the vicinity of the N-B bond or of the -phenylethyl side groups is not large enough to cause any conformational changes which would give rise to a different VCD spectral pattern. Hence, it can be concluded that the observable VCD spectral signatures in the fingerprint region thus solely relate to a change in conformer population, i.e. the ratio of c1 and c2, but they are independent of the state - monomeric, solvated, dimeric - the conformer is in. This is in contrast to other examples, where solvation24, 27-28, 43-44 or aggregation31-32, 45 have shown much stronger effects on the VCD spectral signatures.

CONCLUSIONS In conclusion, the present study showed that the secondary amine borane 1 features different conformational preferences in chloroform and acetonitrile solution. In the non-polar solvent, two



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conformations, c1 and c2 (Figure 2), are almost equally populated, and hydrogen bonding of chloroform to the BH3 moiety was not found to affect the conformational preferences. In contrast, hydrogen bonding to the NH by acetonitrile shifts the preferences towards c1, the conformation with the more accessible NH group (Figure 4). This shift in conformational preferences could not be deduced from the IR spectra, and only by analyzing the unique VCD spectral features of the conformers the calculated preference of c1 could be confirmed. Variable temperature IR and VCD spectra were used to also characterize the dihydrogen bonding driven self-aggregation of 1. The NH stretching region of the VT-IR spectra could be utilized to quantify the degree of dimerization and to estimate the association energy to be -1.39 kcal/mol. In contrast, changes of the IR and VCD spectral features in the fingerprint region were not prominent enough to clearly identify any structural changes. Analysis of the VCD spectra revealed that the interaction with acetonitrile or dimerization does not affect the actual spectral pattern of the conformers of 1, but only their relative contribution to the conformational equilibrium. Hence, hydrogen bonding involving the NH or BH bonds of 1 obviously solely changes the vibrational frequencies of the stretching modes, but none of the deformation modes observed in the fingerprint region. This insensitivity of the fingerprint region for hydrogen bonding interactions is attributed to the rather rigid structure of 1. Although the VCD spectrum could not be used to reveal more insight into the dimerization behavior of 1, the observation of the solvent-induced shift of conformational preferences nicely showcases the interesting potential application of VCD spectroscopy as a tool to characterize interactions of amine boranes with other reactants.46


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ASSOCIATED CONTENT Supporting Information. Concentration and temperature dependence of IR spectra, additional calculated IR and VCD spectra, crystal structure of 1, Cartesian coordinates . This material is available free of charge via the Internet at http://pubs.acs.org.”

AUTHOR INFORMATION Corresponding Author * Dr. Christian Merten, Ruhr-Universität Bochum, Organische Chemie 2, Bochum, Germany; [email protected]; +49 234 32-24529 Funding Sources Deutsche Forschungsgemeinschaft (EXC 1069) Fonds der Chemischen Industrie (Liebig fellowship to CM)

ACKNOWLEDGMENT CM thanks the Fonds der chemischen Industrie for a Liebig fellowship. Furthermore, we acknowledge support by the Deutsche Forschungsgemeinschaft through the Cluster of Excellence RESOLV (“Ruhr Explores SOLVation”, EXC 1069).

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TABLE OF CONTENTS IMAGE


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Intermolecular Interactions of a Chiral Amine Borane Adduct Revealed

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