Lewis Acidic Boranes, Lewis Bases, and Equilibrium Constants: A Reliable Scaffold for a Quantitative Lewis Acidity/Basicity Scale

Article abstract of DOI:10.1002/chem.202003916

A quantitative Lewis acidity/basicity scale toward boron-centered Lewis acids has been developed based on a set of 90 experimental equilibrium constants for the reactions of triarylboranes with various O-, N-, S-, and P-centered Lewis bases in dichloromethane at 20 °C. Analysis with the linear free energy relationship log K_B=LA_B+LB_B allows equilibrium constants, K_B, to be calculated for any type of borane/Lewis base combination through the sum of two descriptors, one for Lewis acidity (LA_B) and one for Lewis basicity (LB_B). The resulting Lewis acidity/basicity scale is independent of fixed reference acids/bases and valid for various types of trivalent boron-centered Lewis acids. It is demonstrated that the newly developed Lewis acidity/basicity scale is easily extendable through linear relationships with quantum-chemically calculated or common physical–organic descriptors and known thermodynamic data (ΔH (Formula presented.)). Furthermore, this experimental platform can be utilized for the rational development of borane-catalyzed reactions.

Full text of DOI:10.1002/chem.202003916

Accepted Article
Accepted Article
Title: Lewis Acidic Boranes, Lewis Bases, and Equilibrium Constants:
A Reliable Scaffold for a Quantitative Lewis Acidity/Basicity Scale
Authors: Robert J. Mayer, Nathalie Hampel, and Armin R. Ofial
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To be cited as: Chem. Eur. J. 10.1002/chem.202003916
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01/2020

10.1002/chem.202003916
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Lewis Acidic Boranes, Lewis Bases, and Equilibrium Constants:
A Reliable Scaffold for a Quantitative Lewis Acidity/Basicity Scale
Robert J. Mayer, Nathalie Hampel and Armin R. Ofial*
[a]
Dr. Robert J. Mayer, Nathalie Hampel, Dr. Armin R. Ofial
Department Chemie
Ludwig-Maximilians-Universität München
Butenandtstr. 5-13, 81377 München
E-mail: ofial@lmu.de
Supporting information for this article is given via a link at the end of the document.
Abstract: In this study, we developed
a
quantitative Lewis
acetophenone, ethyl benzoate,[7] and trimesitylphosphine;[8] in
CD2Cl2: with lutidine[9]).
acidity/basicity scale toward boron-centered Lewis acids based on a
set of 90 experimental equilibrium constants for the reactions of
triarylboranes with various O-, N-, S-, and P-centered Lewis bases in
dichloromethane at 20 °C. Analysis with the linear-free energy
relationship lg KB = LAB + LBB allows to calculate equilibrium constants
KB for any type of borane/Lewis base combination by the sum of two
descriptors, one for Lewis acidity (LAB) and one for Lewis basicity
(LBB). The resulting Lewis acidity/basicity scale is independent of
fixed reference acids/bases and valid toward various types of trivalent
boron-centered Lewis acids. We demonstrate that the newly
developed Lewis acidity/basicity scale is easily extendable through
linear relationships with quantum-chemically calculated or common
physical-organic descriptors and known thermodynamic data (HBF3).
Furthermore, we show that this experimental platform can be utilized
for the rational development of borane-catalyzed reactions.
Instead, it is common practice to characterize Lewis acidity by
following the changes in the spectroscopic features of selected
reference Lewis bases upon their conversion from the free to the
Lewis adduct state.[2b,4,10a,b,11] Alternatively, quantum chemistry
allows to calculate descriptive properties such as orbital energies
or, more advanced, the thermochemistry for the association of
Lewis acids with a certain reference Lewis base, e.g. with fluoride
or hydride.[4] However, to the best of our knowledge, evidence that
such spectroscopy-derived or quantum-chemically calculated
Lewis acidity descriptors correlate with experimental equilibrium
constants for Lewis adduct formations does not exist. Considering
the general lack of experimentally determined equilibrium
constants, it is practically impossible to set up such correlations.
The relative strengths of Lewis bases are more soundly known
from experimentation: Gal and coworkers used calorimetry to
determine the enthalpies ∆H for the reactions of BF3 with a large
set of Lewis bases in dichloromethane.[5] However, it is not
obvious how the tabulated ∆HBF3 value for a certain Lewis base
could assist when one wants to predict the equilibrium constant
for Lewis adduct formation with any other boron-centered Lewis
acid, e.g., the often used B(C6F5)3.
Generally established Lewis acidity and basicity scales are
defined toward a fixed reference Lewis base or Lewis acid,
respectively.[6,12] The infinite number of potential reference Lewis
bases/acids results in an infinite, yet not straightforwardly related
number of potential Lewis acidity/basicity scales.[5] Hence, the
further development of borane catalysis would benefit from
replacing the concept of single reference Lewis acidity/basicity
scales by more versatile multi-reference Lewis acidity/basicity
scales, which allow quantitative predictions.[7c,10a,b,12–14]
Introduction
Lewis acidic boranes are frequently used to catalyze reactions
because they enhance the reactivity of organic compounds by
coordination to Lewis basic sites.[1] The search for even stronger
Lewis acids is still ongoing,[2] and concepts to quantitatively
predict the Lewis acidity of triarylboranes would foster the rational
design of borane-catalyzed reactions.[1e,1f,3]
The extent of Lewis adduct formation for a given borane/Lewis
base combination primarily depends on the electron-accepting
ability of the borane (that is, its Lewis acidity) and the electron-
donating property of the Lewis base (that is, its Lewis basicity).[4,5]
To quantify the Lewis acidity of a borane and the Lewis basicity
of the respective reaction partner in a certain solvent, per
definition experimental equilibrium constants KB for the formation
of Lewis adducts are needed (Scheme 1).[6]
We demonstrate that the systematic variation of the strengths
of the Lewis acids and Lewis bases in combination with isothermal
titration calorimetry (ITC) and NMR spectroscopy yielded an array
of equilibrium constants, which constitutes a unique experimental
platform for the straightforward prediction whether a certain
borane will form a Lewis adduct with a certain Lewis base.
Scheme 1. Equilibrium for the Lewis adduct formation by the reaction of a
borane with a Lewis base.
Results and Discussion
However, only a few equilibrium constants for the association of
triarylboranes with Lewis bases are known. To date, B(C6F5)3 is
the only triarylborane for which equilibrium constants for Lewis
adduct formations with a wider range of Lewis bases have been
determined experimentally (in d6-benzene: with benzaldehyde,
Determination of Equilibrium Constants. In order to
parametrize a Lewis acidity/basicity scale that is independent of
fixed reference compounds, a large set of equilibrium constants
1
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covering a broad range of acidity/basicity is required. In a first step, perfluorinated borane 1i, which does not carry 1H nuclei in its
we determined the equilibrium constants for the formation of
adducts of the triarylboranes 1 with the substituted pyridines 2,
the nitriles 3, the triarylphosphines 4, the carbonyl compounds 5–
9, and triethylphosphine oxide (10) in anhydrous dichloromethane
at 20 °C (Figure 1).
structure, the inverse process was applied: the concentration of
the Lewis base was kept constant and the excess of 1i was varied.
0
-25
-50
-75
-100
-125
-150
-175
-200
-225
0
1000
2000
3000
4000
5000
0
-10
-20
-30
-40
-50
-60
Figure 1. (a) Equilibrium for the adduct formation between a triarylborane and
a Lewis base. (b) Lewis acids and Lewis bases used for the equilibrium studies
in this work.
0.0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
Figure 2. (a) ITC based determination of equilibrium constants KB illustrated for
the reaction between tris(p-anisyl)borane (1b) and pyridine (2d) in
dichloromethane. (b) Detected heat flow during the titration of a 1b solution in
CH2Cl2 with a 2d solution in the same solvent. (c) Integrated heat flow Q vs.
the molar ratio of the Lewis base/borane (black dots) and fitted curve (blue line)
giving KB for the individual titration.
Isothermal titration calorimetry (ITC) is a powerful method to
characterize the interactions of boranes with Lewis bases. In a
microcalorimeter, small amounts of Lewis base (typically 40 ×
6 μL portions of a CH2Cl2 solution) were added to solutions of the
triarylboranes (1.8 mL, 1 mM) in dichloromethane at 20 °C, which
gave rise to a heat evolution within the sample cell (Figure 2a,b).
The resulting heat signals were integrated to give the heats per
injection. The correlation of the heats per injection with the molar
ratio of Lewis base and borane gave a binding isotherm (Figure
2c), which was analyzed with a 1:1 interaction model[15] to derive
the associations constants KB. Three individual ITC experiments
for each Lewis acid/Lewis base combination were performed.
These individual results were averaged to determine the
equilibrium constant KB. Thus, ITC allowed to reliably determine
equilibrium constants in the range of 103 < KB < 107 M–1. However,
our ITC instrument cannot be operated under completely inert
conditions, and side reactions with traces of moisture were a
limitation for boranes more acidic than 1e.
The ITC method was, therefore, complemented by 1H NMR
spectroscopic titrations, which were generally used to study
equilibrium constants KB < 103 M–1 (Figure 3). All NMR titration
experiments were performed in
a glovebox under a dry
atmosphere of argon. A series of samples with constant
concentration of borane and variable concentrations of the Lewis
1
base in CD2Cl2 was prepared, sealed, and analyzed by H NMR
spectroscopy (Figure 3b). Plotting the change of the chemical shift
 of a resonance assigned to the Lewis acid 1 as a function of
the concentration of the Lewis base yielded a binding isotherm,
which was analyzed by a 1:1 binding model (Figure 3c).[16] For the
Figure 3. (a) NMR spectroscopic determination of equilibrium constants KB
shown for the reaction between 2h and 1b. (b) 1H NMR spectra (400 MHz,
CD2Cl2) of borane 1b in the presence of variable concentrations of 2h. (c) Plot
of the chemical shift difference  of the methoxy groups of 1b at different
concentrations of 2h, and the result of numerical fitting used to derive KB (red).
2
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The reaction of the borane 1b with 4-benzoylpyridine (2e) was
investigated by both the ITC and the NMR titration methods,
which furnished KB = (3.25 ± 0.27) × 103 M–1 by ITC and KB = (3.14
± 0.06) × 103 M–1 by the NMR titration. We, thus, assumed that
equilibrium constants KB originating from both methods can be
interlaced for the further data analysis.[17] Overall, we determined
a set of 90 equilibrium constants KB for various combinations of
boranes and Lewis bases (see Table S1 and the Supporting
Information for the individual measurements).
DMAP (2a), the strongest Lewis base in this study, forms a
Lewis adduct with trimesitylborane (1j). As depicted in Figure 5,
the equilibrium constant KB = 0.454 M–1 was used to calculate the
Lewis acidity parameter LAB = –9.0 for trimesitylborane 1j. We
assume that this low Lewis acidity parameter rather reflects the
steric shielding of the boron atom by the three adjacent mesityl
groups than the intrinsic Lewis acidity of the boron center.
Constructing a Quantitative Lewis Acidity/Basicity Scale.
Existing Lewis acidity scales rely on a single reference Lewis base
to compare the relative strength of different Lewis acids.[4,5,10–12]
Yet, a single reference Lewis base is not sufficient to compare
boranes of widely differing Lewis acidities due to the limited range,
in which equilibrium constants can be determined experimentally:
the moderately Lewis basic 3,4,5-trichloropyridine (2j) establishes
equilibria for Lewis adduct formation with the donor- and weakly
acceptor-substituted triarylboranes 1b-f. However, selecting the
Lewis base 2j as the reference Lewis base would neither allow to
study the association equilibria with the less Lewis acidic 1a nor
with the stronger Lewis acids 1g-i.
Figure 4. (a) Correlation of lg KB for the reactions of 1c with pyridines 2 with lg
exp
KB for the reactions of 1e with pyridines 2. (b) Correlation of lg KBeq 1 with lg KB
for reactions of the boranes 1 with the Lewis bases 2, 3, and 5–9.
Hence, we decided to use a floating scale of Lewis basic
reference compounds, which relies on combining Lewis acids with
overlapping sets of differently strong Lewis bases. Thus, strong
Lewis bases were used to characterize weak Lewis acids, while
weaker Lewis bases were used to study equilibria of Lewis adduct
formation with stronger Lewis acids.
As illustrated by the linear correlation of the equilibrium
constants KB for boranes 1c and 1e with a series of pyridines
(slope = 1.03), the relative values of the equilibrium constants are
the same for the Lewis base-association with different
triarylboranes (Figure 4a). Similar correlations have previously
been observed for equilibrium constants of the reactions of Lewis
bases with diarylcarbenium ions (Ar2CH+).[13] Data analysis shows
that the two-parameter equation (1) is sufficient to calculate the
equilibrium constants KB for the triarylborane/Lewis base
combinations determined in this work.
lg KB = LAB + LBB (in dichloromethane at 20 °C) (1)
For Lewis acid/base reactions in dichloromethane solution eq. (1)
relates the experimentally determined lg KB with a Lewis acidity
parameter LAB specific for a certain triarylborane and an LBB
parameter referring to a specific Lewis base. Defining the Lewis
acidity of triphenylborane LAB(1d) = 0 and then performing a least-
Figure 5. Least squares minimization of the equilibrium constants (lg KB)
according to the two-parameter equation (1) generates Lewis acidity and Lewis
basicity scales.
squares minimization furnished
a set of 9 Lewis acidity
parameters LAB for 1a–1i and 27 Lewis basicity parameters LBB
Figure 6 illustrates that the thus established Lewis acidity/basicity
scales both cover 15 orders of magnitude. Eq. (1) then provides
a reliable tool to predict absolute equilibrium constants for
arbitrary combinations of triarylboranes of known LAB with Lewis
bases of known LBB. Further types of Lewis bases can
straightforwardly be integrated in the LBB scale by measuring the
equilibrium constants of their reactions with some of the
for 2, 3, 5, and 7–10.[18] Applying LAB and LBB in eq. (1) results in
eq 1
calculated equilibrium constants KB
that deviate at maximum
by a factor of 4.5 from KBexp determined by experiment (Figure 4b).
The Lewis basicities LBB for another six Lewis bases (2n, 4a-c,
5f, 6) were estimated on the basis of available KB for association
with only a single reference Lewis acid 1.
In Figure 5, the parallel correlation lines obtained by plotting
lg KB against the Lewis basicity parameter LBB illustrate eq. (1)
graphically. The negative intercepts of the correlation lines with
the abscissa correspond to the Lewis acidity parameters LAB of
the triarylboranes 1.
characterized triarylboranes. For example, an averaged LBB
–0.90 for tetrahydrothiophene (11) was determined from the
equilibrium constants for its reactions with 1f and 1h (Table S1).
=
3
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Figure 6. Experimental Lewis acidity and basicity scales for boranes and N-, O-, S-, and P-centered Lewis bases derived on the basis of equation (1). Compounds
located on the same vertical level combine with an equilibrium constant KB = 1 M–1 in dichloromethane at 20 °C. [a] LAB(1j) not generally valid. [b] LAB for BX3 (X =
F, Cl, Br) were derived computationally and linked to experiment via isodesmic reactions (see text). [c] Only a single equilibrium constant KB was available for the
determination of LBB. [d] LBB(4a–c) are only valid toward Ar3B without ortho-substituents (see text). [e] LBB(Et2O) was calculated from data in ref [19] (see text).
Structure-affinity relationships for Lewis adduct formations of the
characterized acids and bases can now be analyzed on an
experimentally relevant, quantitative basis. Systematic deviations
are expected when Lewis adduct formation is hampered by steric
impositions, and some of such examples will be discussed below.
For a robust averaging, we used three structurally diverse Lewis
bases to estimate the LAB values of boron trihalides. Based on
the experimental G(II) for the formation of Lewis adducts of
1d/2d, 1f/3a, and 1h/7b, we estimated LAB of BF3, BCl3, and BBr3
(in CH2Cl2) by calculating G(I) for the transfer of pyridine (2d),
acetonitrile (3a), and benzaldehyde (7b). The individual LAB
parameters obtained from Giso for the three different reference
bases were then averaged to give LAB = 8.4 ± 2.0 for BF3, LAB =
9.3 ± 1.8 for BCl3, and LAB = 10.1 ± 1.3 for BBr3 (see Supporting
Information). Though we have to acknowledge significant errors
for these LAB values, the qualitative ordering of Lewis acidities
LAB with 1i < (or ≈) BF3 < BCl3 < BBr3 accords with Lewis acidity
rankings based on spectroscopic data.[10] The results of this
analysis are included in the Lewis acidity scale shown in Figure 6.
Access to LAB of BF3, BCl3 and BBr3 by DFT Calculations.
Next, we explored whether the experimental Gibbs reaction
energies rG (= –RT ln KB) for a representative set of B-N, B-O,
and B-P Lewis acid/base combinations (that is, 1d/2d, 1f/3a,
1h/7b, and 1f/4b) could be reproduced by using density functional
theory (DFT) methods for quantum-chemical calculations. For this
purpose, several DFT methods were tested. Optimum agreement
between calculated and experimental Gibbs reaction energies
rG was reached when we applied the SMD(DCM)/MN15/def2-
TZVP method[20] (average deviation: −5.4 ± 2.4 kJ/mol, see Figure
S6, Supporting Information for details).
Subsequently, quantum-chemically calculated thermodynamics
was utilized to estimate LAB values of further boranes BX3.[4,21] As
outlined in Figure 7, the Gibbs reaction energies G(I) for
isodesmic Lewis base transfer reactions between the Lewis
adduct Ar3B–LB and a borane BX3 were computed by using the
SMD(DCM)/MN15/def2-TZVP method. For anchoring the
Figure 7. Combining the isodesmic reaction (I) with an experimental reference
boranes BX3 at the LAB scale,
a known, experimentally
reaction (II) allows one to determine the Lewis acidities of boranes BX3 from
determined Gibbs energy G(II)exp for the formation of a Lewis
adduct of Ar3B and the investigated Lewis base was added to
G(I) to give ∆Giso. The Gibbs energy ∆Giso in eq. (III) now
describes the thermodynamics of the Lewis adduct formation
between BX3 and a Lewis base, which can be used to estimate
the Lewis acidity LAB of BX3 if LBB of the Lewis base is known.
Giso
.
LAB and LBB from Competition Reactions. The parallel
correlation lines in Figure 5 imply that the relative Lewis basicity
of a pair of Lewis bases is independent of the Lewis acid used for
the comparison, and vice versa. From this it follows that reported
4
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equilibrium constants for borane complexation, as shown in eq (2), regard to reflecting the experimental Lewis acidities LAB of the
also reflect the relative Lewis basicity of the two competing Lewis
bases LB1 and LB2.
entire set of investigated boranes: LAB for the boranes 1a-1l, BF3,
BCl3, and BBr3 correlated linearly with FIADCM (Figure 8a).
The shared linear correlation of different types of boranes
toward one specific reference Lewis base, that is, the fluoride ion,
highlights that the LAB scale presented herein is generally valid,
not only toward triarylboranes but also toward other types of
boranes.
For example, the thermodynamics for the reactions of BF3∙OEt2
with benzaldehyde (7b, LBB = –1.78) and p-anisaldehyde (7a, LBB
= –0.18) has been investigated.[19] The equilibrium constant for
the reaction of 7a with BF3∙OEt2 to yield 7a·BF3 and diethyl ether
was determined to be K = 7.2 (in CDCl3 at rt). This indicates that
diethyl ether is by 0.86 LBB units less Lewis basic than 7a. The
analogous BF3 transfer from BF3∙OEt2 to 7b gave an equilibrium
constant of K = 0.16 (in CD2Cl2 at rt; K = 0.19 in CDCl3 at rt). Thus
diethyl ether is 0.8 LBB units more Lewis basic than 7b. By
Quantum-chemical descriptors, such as LUMO energies (LUMO
)
or global electrophilicity indices (GEI), were proposed as
alternative measures of Lewis acidity.[1f,4,24] The experimental
Lewis acidities LAB of triarylboranes 1a-1i correlated with LUMO
and GEI (both calculated for the gas phase and for the SMD
solvent model for dichloromethane). However, neither LUMO
energies nor GEIs are generally applicable to a broader scope of
boranes because boron trihalides as well as 1k-1l gave separate
correlations of their estimated LAB with LUMO or GEI (Figures
S1b,c,e,f). We conclude that LUMO energies and GEI can be
used to compare the relative Lewis acidities of structurally
analogous boranes. However, both quantum-chemical
descriptors are limited when a comparison of structurally diverse
boranes is intended.
averaging both results and neglecting the solvent effect LBB
–1.0 is assigned to diethyl ether.
=
In a similar manner, also the Lewis acidities LAB of further
boranes can be estimated from equilibrium constants for the
exchange of Lewis bases between two competing Lewis acids, as
depicted in equation (3).
Figure 8b illustrates that a linear correlation of good quality
exists for Lewis acidity parameters LAB of triarylboranes with the
sum of the Hammett substituent parameters Σσ.[25] This
correlation allows to extrapolate LAB for further triarylboranes with
various substitution patterns without the need for quantum-
chemically calculated descriptors.
Based on the equilibrium constant for the complexation of
acetonitrile (3a) by B(C6F5)3 (LAB = 7.24) and a competing borane,
the weaker Lewis acid tris(perfluoro--naphthyl)borane (1k, K =
1/2.6 = 0.39 at 20 °C in d8-toluene, LAB = 6.8)[22] and the
marginally stronger Lewis acidic perfluorinated 9-phenyl-9-
borafluorene (1l, K = 1.3, LAB = 7.35)[23a] can be characterized.
The boranes 1k and 1l are, therefore, tentatively positioned to the
left and the right, respectively, of B(C6F5)3 in the Lewis acidity
scale in Figure 6.[23b]
Figure 8. (a) Correlation of LAB with fluoride ion affinities FIADCM in
dichloromethane [∆G at the SMD(DCM)/MN15/def2-TZVP level of theory, for
individual data see Supporting Information]. The entries for BF3, BCl3, BBr3, 1k
and 1l (grey triangles) and trimesitylborane 1j (red circle) were not used for the
calculation of the correlation line. (b) Correlation of LAB parameters with
Hammett σ parameters.
It is straightforward that, once integrated in the Lewis
acidity/basicity scales of Figure 6, equation (1) enables to
estimate equilibrium constants for many further reactions of the
newly characterized Lewis acids or bases.
Accessing Further Lewis Basicities LBB. The Lewis basicity
parameters LBB cover almost 15 orders of magnitude from highly
basic pyridines to weakly basic aldehydes and ketones.
Within different classes of Lewis bases (pyridines,
benzaldehydes, and acetophenones), Hammett σ constants[25a]
can be used to extrapolate further Lewis basicities (Figure S2).
Access to Further Lewis Acidity Parameters by Correlations.
Gas-phase fluoride ion affinities (FIA, calculated as H) have
often been considered to be a measure for Lewis acidity.[4]
Accordingly, we observed a linear correlation of LAB with gas-
phase FIAs for para-substituted triarylboranes (Table S2 and
Figure S1a, Supporting Information). However, LAB for the
boranes 1i and 1g with ortho-fluoro substituents deviate from the
correlation line, and the estimated LAB values for BX3 (X = F, Cl,
Br) also departed significantly. We anticipated that the
consideration of the solvent and entropy corrections would be
crucial to improve the scope of the LAB vs FIA correlation. Indeed,
calculations of FIAs as ∆G in dichloromethane (FIADCM) by the
SMD(DCM)/MN15/def2-TZVP method were more general with
LBB parameters for pyridines correlate linearly with their Brønsted
[25b-e]
basicities pKaH
(Figure 9a) and with Lewis basicities LB
determined from equilibrium reactions with diarylcarbenium ions
(Figure 9b). The slope of 0.45 of the linear relationship in Figure
9b indicates a weaker bonding of pyridines toward boron-
centered than toward carbon-centered Lewis acids in accord with
the lower bond dissociation energies for B-N bonds if compared
to C-N bonds.[26]
5
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A consistent linear correlation of Lewis basicities LBB over
different classes of Lewis bases (except for the S-centered 11)
exists with Gal’s experimental enthalpies for BF3-complexation
HBF3 (Figure 10a),[5]. Insofar, Figure 10a corroborates what is
already indicated in Figure 8a: the multi-referenced Lewis basicity
ranking, which is based on a scaffold of reactions toward
triarylboranes, also holds toward further types of boranes (here:
BF3).
For Lewis bases that lack experimental HBF3, borane affinities
BA, that is, the Gibbs reaction energy for the addition of BH3 to a
Lewis base in dichloromethane, provide an equally reliable tool
for estimating LBB. In this work, borane affinities BA were
quantum-chemically calculated from isodesmic reactions, which
were anchored to available experimental rG (Supporting
Information). A uniform linear correlation of LBB vs BA is obtained
(Figure 10b) indicating that the change of the borane reference
Lewis acids from BAr3 to BH3 does not have a significant impact
on the relative Lewis basicity of sterically unbiased O-, S- and N-
centered Lewis bases.
For example, the linear LBB vs pKaH relationship for pyridines in
Figure 9a predicts LBB = 7.3 for 2,6-lutidine (2n, pKaH 6.72). The
experimentally observed K = 54.7 M–1 for the reaction of 2n with
1i gives a Lewis basicity LBB = –5.5 for 2n, however. This LBB
shows that adduct formation of 2n with B(C6F5)3 (1i) is in fact
attenuated by almost 13 orders of magnitude because of the steric
imposition of the pyridine’s methyl substituents with the aryl
groups of the BAr3 fragment (Figure 11a). While H+ is the smallest
possible Lewis acid, BH3 and BF3 are large enough to gradually
experience steric repulsion with the methyl groups of 2n. As a
result, the data for 2n deviate by only 7 to 8 orders of magnitude
from the linear correlations observed for sterically unbiased Lewis
bases in Figure 10.
Reactions of triarylboranes 1 and triarylphosphines 4a-c yield
hexaaryl phosphonium-boronates, which are isoelectronic to the
hypothetical hexaarylethane.[27] It is hence unsurprising that the
triarylphosphines 4a-c deviate from the correlation of LBB for other
Lewis bases with borane affinities BA. The boron/phosphorous
analogs are formed with equilibrium constants 5 to 6 powers of
ten lower than expected based on the correlation for the formation
of less strained Ar3P-BH3 adducts (Figure 10b/11b).[28]
Figure 11. (a) Repulsive steric interactions in Lewis adducts of triarylboranes
with 2,6-lutidine (2n) and of (b) triarylphosphines.
Figure 9. (a) Correlation of LBB for pyridines with (a) their pKaH (from ref.[25a-e]
and gathered in Table S3) and (b) their Lewis basicities LB (toward Ar2CH+ in
CH2Cl2, from ref [13]).
The fact that contacts of the Lewis acid’s and the Lewis base’s
aryl rings exist in the Lewis adducts formed by combining
triarylboranes
with
triarylphosphines[8,18c]
makes
the
triarylphosphines sensitive probes for subtle changes in the steric
demand of the triarylboranes. While the almost identical Lewis
acidities of 1g (LAB = 4.08) and 1h (LAB = 3.98) indicate that the
pattern of fluoro-substitution at the borane has a negligible impact
on the affinity toward sp- or sp2-centered (rodlike) reference Lewis
bases, this is not true toward the triarylphosphines 4. For the
reaction of tris(p-anisyl)phosphine (4a) with the borane 1e (p-F)
we determined KB = 360 M–1, which was used to estimate LBB =
2.43 for 4a. Accordingly, the association of 4a with the 3,4,5-
fluorinated borane 1h proceeded quantitatively (KBeq 1 = 2.4 × 106
M–1). The analogous reaction of 4a with the 2,4,6-trifluoro-
substituted triphenylborane 1g, however, gave an experimental
KBexp = 861 M–1 which is more than 3 logarithmic units lower than
predicted for this Lewis adduct formation (KBeq 1 = 3.3 × 106 M–1).
The latter finding indicates that ortho-fluorine atoms in
triarylboranes can cause repulsive effects if these triarylboranes
associate with sterically demanding Lewis basic counterparts.[29]
Moreover, GEI or LUMO energies, which inform about the sheer
electronic properties of triarylboranes, may be applied to quantify
bigger effects of steric hindrance on the acidity of boranes. The
equilibrium constant for the reaction of trimesitylborane (1j) with
DMAP (2a) yields LAB = –9.0 for 1j, which is 8 to 9 orders of
magnitude lower than that predicted by its GEI or εLUMO values
(Figure S1e,f). The significant deviations from the correlation lines
Figure 10. Correlation of LBB parameters with further types of experimental or
quantum-chemically calculated basicity descriptors (open data points not used
for the construction of the correlation line). (a) Correlation of LBB parameters
with experimental BF3 affinities in dichloromethane (–∆HBF3) from ref [5]. (b)
Correlation of LBB parameters with quantum-chemically calculated BH3 affinities
BA (at the SMD(DCM)/MN15/def2-TZVP level of theory, referenced to the
experimental ∆rG of the reaction of 1d with 2d as outlined in the Supporting
Information).
Steric Effects. Correlations of LBB, which were parameterized
toward triarylboranes in this work, with Lewis basicity descriptors
that were determined toward relatively small Lewis acids such as
H+ (i.e., pKaH) or BH3 (i.e., borane affinities BA) may be used to
quantify the magnitude of repulsive steric effects in Lewis base-
borane adducts.
6
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in Figures S1e,f clearly indicate that equilibrium constants for
Lewis adducts that contain trimesitylborane (1j) will strongly
depend on the structure of the Lewis base.[30]
triarylboranes 1i (LAB = 7.2), 1g (LAB = 4.1), and 1h (LAB = 3.9),
but not when 1f (LAB = 1.3) was used as the catalyst (Figure 13a).
The Lewis acidities determined in this work also reflect the
different efficacies of boranes to catalyze the Diels-Alder reaction
of methyl vinyl ketone with cyclopentadiene (Figure 13b, Figure
S4) and the Michael addition of 1-methylindole to methyl vinyl
ketone[32] (Figure 13c, Figure S5). The quantum-chemically
calculated BA allowed to estimate the Lewis basicity LBB = −2.5
Scope of NMR-Based Lewis Acidity Rankings. NMR assays
are often used to quantify Lewis acidity. For example, the Childs
method[11a] utilizes proton 3-H in trans-crotonaldehyde (8,
reference Lewis base) to derive a measure for Lewis acidity from
the difference of the NMR chemical shifts in the free 8 and in the
corresponding Lewis adduct (Figure 12a). To be meaningful, the
Childs method assumes quantitative formation of the Lewis
adduct. With LBB(8) = –0.73, eq. (1) can be used to derive that
of methyl vinyl ketone. Accordingly, minor conversion was still
eq 1
obtained with the catalyst 1f, for which KB
= 6.8 × 10–2 M–1 is
calculated, suggesting that per mille levels of Lewis adduct are
sufficient to enable the reactions. While faster Diels-Alder
reactions and Michael additions were observed with borane
catalysts that are more Lewis acidic than 1f, the less acidic
triphenylborane (1d) was inapt to catalyze both reactions even
after extended reaction times (Figure 13b,c).
this assumption is true for the highly Lewis acidic boranes 1h-1i
eq 1
(KB
= 103–106 M–1, calculated with the LAB and LBB values in
Figure 6). However, many less Lewis acidic boranes cannot be
ranked by the Childs method. For example, quantitative Lewis
adduct formation of tris(4-chlorophenyl)borane (1f) with 8 is
unlikely to be achieved (KBeq 1 = 4.0 M–1). Moreover, Lewis adduct
formation of the aldehyde 8 with the even less Lewis acidic
triphenylborane (1d) is endergonic (KBeq 1 = 0.19 M–1).
Product inhibition may impair catalytic transformations. At
present and based on our results, it is not possible to define a
certain threshold value for the difference in Lewis basicity (LBB)
between the substrate and the product. Figure 13a shows,
however, that LBB = +2 was still tolerated in Nazarov cyclizations.
Notably, not only highly Lewis acidic fluorinated triarylboranes
are utilized as catalysts. Already triphenylborane (1d) forms
adducts with tertiary amides, such as N,N-dimethylacetamide (9,
exp
LBB = 0.97) (KB = 7.6 M–1), enabling their selective reduction
with hydrosilanes as hydride donors in acetonitrile (LBB = –0.48),
which is only a slightly weaker Lewis base than the amide 9.[33]
Less Lewis acidic boranes as catalysts may be beneficial in future
applications because of their advantageous chemoselectivity in
reactions at highly functionalized substrates, which tend to react
unselectively when highly Lewis acidic boranes are used as
catalysts.[33,34]
Figure 12. Probe molecules for (a) the Childs method[11a] and (b) the Gutmann-
Beckett method.[11b]
The Gutmann-Beckett method follows the 31P NMR chemical shift
of triethylphosphine oxide (Figure 12b, 10),[11b] which is a stronger
Lewis base (LBB = 2.51) toward boranes than 8. The practical use
of the Gutmann-Beckett method is impeded, however, by
significant scatter in published data of 31P NMR chemical shifts for
Lewis adducts of 10. For example, reported 31P NMR chemical
shifts for the Et3P=O/BPh3 adduct (10 + 1d) cover a range from
65.9 to 72.5 ppm (in C6D6) owing to different direct and indirect
methods applied for their determination.[10] Again, the single
reference Gutmann-Beckett method is not able to characterize
weak Lewis acids with LAB < –1.
Conclusion
From a set of 88 experimental equilibrium constants, we
constructed
a Lewis acidity/basicity scale that allows the
quantitative prediction of absolute equilibrium constants for the
associations between triarylboranes and various O-, N-, S-, and
P-centered Lewis bases covering 15 orders of magnitude in
acidity/basicity. As a consequence, eq (1) enables chemists to
predict whether at all and if ‘yes’, to which extent Lewis adduct
formation of a certain Lewis acidic borane with a certain Lewis
base in dichloromethane will be possible (Figure 14).
Quantum-chemical calculations and isodesmic reactions were
used to additionally assess the LAB parameters of halogenated
boranes BX3 (X = F, Cl, Br). A common linear correlation with
fluoride ion affinities in dichloromethane for all types of boranes
was found, highlighting that the acidity/basicity scale on the basis
of eq. (1) is not only valid toward triarylboranes but that it is of
more general applicability for trivalent boron compounds.
Lewis bases have a constant relative basicity toward the
various investigated boranes, as illustrated by the parallel
correlation lines in Figure 5. Linear free energy relationships of
LAB and LBB with physical organic descriptors such as Hammett
σ-parameters, tabulated reaction enthalpies as well as quantum-
chemically calculated thermodynamics offer direct avenues for
chemists to introduce further sterically unbiased substrates to the
Lewis acidity/basicity scales presented herein, which are the first
Applying Quantitative Acidities and Basicities in Synthesis.
To investigate the relevance of our Lewis acidity/basicity scale for
a rational design of organic synthesis, we studied triarylborane-
catalyzed reactions of carbonyl compounds. Product inhibition
may be a limitation in the investigated Nazarov cyclizations, Diels-
Alder reactions and Michael additions. As long as the product is
of comparable Lewis basicity, however, the reactions will work in
the presence of catalytic amounts of the Lewis acid.
For example, Nazarov cyclizations were reported to be
catalyzed by 5-mol% of B(C6F5)3 (1i) whereas BPh3 (1d, LAB = 0)
was ineffective (Figure 13).[31] The quantum-chemically calculated
BA of the divinyl ketone in combination with the correlation in
Figure 10b gives a Lewis basicity LBB of −2.9, which suggests that
activating association might only occur with Lewis acids with LAB
> 3. This is in line with our results (10 mol-% BX3 or BAr3, CD2Cl2,
25 °C, Figure S3), which show that conversion to the
cyclopentenone product was observed for BBr3, BCl3 and the
7
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10.1002/chem.202003916
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Figure 13. Application of the Lewis acidity parameters LAB in borane-catalyzed reactions: (a) Nazarov cyclizations. Borane 1g (with ortho-fluorine atoms) is a less
efficient catalyst than the equally Lewis acidic 1h (with only ortho-hydrogen atoms). (b) Diels-Alder reactions of methyl vinyl ketone with cyclopentadiene. (c) Michael
additions of 1-methylindole to methyl vinyl ketone (ref [32] reports catalysis of this Michael addition by 5 mol-% 1i in CHCl3 at 80 °C).
extensive scales for boranes derived from experimental
equilibrium constants. They can now be used to assess the scope
of methods that characterize Lewis acidity/basicity without
equilibrium data.
It is an often recognized limitation of the Lewis acid/base
concept that the position of a certain compound in Lewis
acidity/basicity scales depends on the choice of the reference
reaction(s).[5,12] Knowledge of equilibrium constants, as
determined in this work, does not solve this multidimensional
issue. For example, for a set of Lewis bases equilibrium constants
determined against diaryliodonium ions (Ar2I+) yield a different
ordering of Lewis basicity than analogous measurements against
diarylcarbenium ions (Ar2CH+).[13,14] As a consequence, also the
construction of a unified Lewis acidity scale that involves a wide
array of different types of atoms at the reactive site is elusive at
present.[12]
Yet, the relative Lewis basicities of substituted pyridines
Figure 14. The two-dimensional arrangement of Lewis acidity and Lewis
followed the same order toward boron-centered (Ar3B) and
carbon-centered (Ar2CH+)[13] Lewis acids, as depicted in Figure 9b.
Owing to little overlap of the studied Lewis bases it remains to be
basicity scales illustrates which Lewis acid/base pair forms a Lewis adduct (in
dichloromethane).
8
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Psaltakis, J. J. Cortes, S. S. Siddiqi, J. J. Devery III, J. Org. Chem. 2020,
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Acknowledgements
[17] We considered the small difference in temperature between the ITC
(20 °C) and NMR measurements (22 °C) to be negligible for the
subsequent analysis.
We thank Amalina Buda for experimental assistance at the
beginning of the project, Prof. Herbert Mayr (LMU) for helpful
discussions and the Department Chemie for financial support.
R.J.M. is grateful to the Fonds der Chemischen Industrie for a
Kekulé fellowship.
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9
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Entry for the Table of Contents
Loose or Lewis? A comprehensive set of experimental equilibrium constants informs about the strength of the interactions between
boron-centered Lewis acids and N-, O-, S-, and P-centered Lewis bases. Data analysis shows that the two-parameter equation (1)
enables to straightforwardly predict Lewis adduct formation for borane/Lewis base couples. The available experimental data can be
supplemented by quantum-chemical calculations which helps to streamline the design of borane-catalyzed reactions.
10
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