Supporting Information
A Study of the Mechanism of Electron Transfer Quenching by Boron-Nitrogen
Adducts in Fluorescent Sensors
Stefan Franzen*, Weijuan Ni, Binghe Wang*
Department of Chemistry
North Carolina State University
Raleigh, NC 27695
*Authors to whom correspondence should be addressed:
Email: [email protected], [email protected]
Phone: (919)-515-8915, (404)-651-0289
Current address for BW: Department of Chemistry, Georgia State University, Atlanta,
GA 30303
1
Supporting Information
The boron- nitrogen interaction in phenyl boronic esters is weak even by the
standards of boron-nitrogen adducts in the literature. (Umeyama and Morokuma 1976;
Brint, Sangchakr et al. 1989; Glendening and Streitweiser 1994; Jonas, Frenking et al.
1995; Leboeuf, Russo et al. 1995; Rablen and Hartwig 1996; Anane, Boutalib et al. 1997;
Anane, Jarid et al. 1998; Anane, Jarid et al. 2000) We have performed density function
theory (DFT) calculations for various models of the type BXY3 -NZ3 . The calculations
have been carried out using DMol3 with the generalized gradient approximation (GGA)
(Delley 2000) and also using Gaussian98 with a 6-31** basis set using the B-LYP
exchange and correlation functional. (Becke 1988; Becke 1999) The calculations were
carried out at the North Carolina Supercomputer Center on the IBM SP and SGI/Cray
Origin 2000. The potential energy surfaces (PESs) for the intermolecular adducts shown
in Figure S1 were obtained by plotting single point energy at displacements of the donor,
NZ3 and acceptor, BXY2 along the B-N bond from 1 Å to 10 Å (BXY2 ßR(Å)à NZ3 ).
Such a potential energy surface approach uses the calculation of the donor and acceptor at
10 Å as the reference state. The energy of the reference state is used as the zero of
energy to determine the relative interaction energy of the adduct. Energies were also
calculated using thermodynamic equilibria of the type shown below:
BXY2 :NZ3 à BXY2 + NZ3
The energy of adduct formation is calculated as E(BXY2 ) + E(NZ3 ) - E(BXY2 :NZ3 ),
where the energies E are the binding energies determined by DFT. For boronate ester
formation with ethylene glycol the hydrolysis equilibrium was considered.
BXC2 H4 O2 :NZ3 + 2 H2 O à BX(OH)2 :NZ3 + C2 H6 O2
where X = OH, CH3 or C6 H5 , Y2 = (OH)2 , C2 H4 O2 or C3 H6 O2 and Z = H or CH3 . Each
molecule was geometry optimized separately in these calculations leading to possible
artifacts in the energy calculation.
Intermolecular models: Units of adducts
B
OH
O
B
OH
PBA
O
B
O
O
PBE
PBP
OH
H3C B
OH
MBA
O
H3C B
O
MBE
Adducts
PBA : NH3
PBA : TMA
PBE : NH 3
PBE : TMA
PBP : NH 3
PBP : TMA
MBA : NH3
MBA : TMA
MBA : NH3
MBE : TMA
Figure S1. Structures of intermolecular adducts studied by DFT methods. The structural
units consist of phenyl and methyl boronic acids and esters. The adducts of these
molecules with ammonia and trimethylamine (TMA) were studied to determine the
strength of the boron-nitrogen bond and conformation of the boronic acid/ester relative to
the phenyl ring. The biosensor models include both boronic acid/ester moiety and
trimethyl amine base covalently attached to a phenyl ring.
2
The comparison of ethane and borane is
a well-studied example of the relative strength
of the B-N relative to the C-C bond. The bond
length BH3 -NH3 is 1.657 Å compared to 1.524
Å for CH3 -CH3 as determined from a DFT
calculation shown in Figure S2 (Umeyama and
Morokuma 1976; Anane, Jarid et al. 1998). DFT
has been shown to give excellent geometries,
dipole moments and vibrational frequencies for
dative bond adducts shown in Figure S1,
Figure S2. Potential energy surfaces for
although it tends to overestimate the binding
and boron-nitrogen analog. The
energy (Holme and Troung 1993). The potential ethane
curves were calculated by changing the C-C
energy surfaces for BH3 -NH3 and CH3 -CH3 shown or B-N bond length for a fixed geometry of
in Figure S2 indicate that the bond energies are – the remaining atoms.
46.7 kcal/mol (-195.4 kJ/mol) and –118.7
kcal/mol (-496.6 kJ.mol), respectively. The
calculated bond energies for BH3 -NH3 and
BH3 -N(CH3 )3 are –195.4 and –190.0 kJ/mol,
which can be compared to experimental
values of –141.0 kJ/mol and –172.8 kJ/mol,
35 respectively. The B-N bond is significantly
weaker than the C-C bond due to poorer
overlap in the asymmetric B-N structure.
Nonetheless, the B-N bonds in BH3 -NH3 and
BH3 -N(CH3 )3 are by far the strongest in Table
S1.
These calculations and calculations Figure S3. Potential energy surface for the
with boronic acid (B(OH)3 ) complexed with
phenyl boronic acid and ester adducts with
various Lewis bases (NH3 , N(CH3 )3 ) trimethylamine. Both the PBE and PBP
presented in Table S1 are precursors to models model molecules have identical surfaces.
relevant to phenyl boronic esters. The Lewis
acids considered include borane, boric acid (BA), methyl boronic acid (MBA),
phenylboronic acid (PBA) and the corresponding esters BE, MBE and PBE. The Lewis
bases considered are NH3 and N(CH3 )3 (TMA).
3
Table S1. B-N bond lengths, energies and charges obtained from potential energy surface
studies. The phenyl boronic acid and boronate ester derivatives were calculated using a
Cp-Cp-B-O dihedral angle of 90o .
Model
d(B-N) (Å)
Bond energy ESP Charge ESP Charge
(kJ/mol)
Boron
Nitrogen
BH3 : NH3
BH3 : TMA
BA:NH3
BE:NH3
MBA:NH3
MBE:NH3
PBA:NH3 0
PBE:NH3 0
PBP:NH3 0
PBA:NH3 90
PBE:NH3 90
PBP:NH3 90
BA:TMA
BE:TMA
MBA:TMA
MBE:TMA
PBA:TMA0
PBE:TMA0
PBP:TMA0
PBA:TMA90
PBE:TMA90
PBP:TMA90
1.66
1.52
1.79
3.24
3.27
3.18
3.40
2.80
3.25
3.35
2.73
3.20
1.81
4.65
2.66
2.57
3.86
2.70
3.98
3.70
3.85
4.00
195.4
190.0
14.9
14.0
13.4
14.6
18.7
22.2
25.0
14.2
14.2
13.8
22.2
20.3
28.3
30.3
17.1
29.7
16.8
15.5
15.9
16.4
0.22
0.32
0.82
0.80
0.77
0.88
0.49
0.70
0.71
0.77
0.87
0.84
0.69
0.82
0.77
0.88
0.49
0.70
0.73
0.60
0.71
0.81
-0.44
-0.44
-0.64
-0.76
-0.94
-0.93
-0.82
-0.68
-0.73
-0.88
-0.81
-0.86
0.32
-0.06
0.21
0.22
0.20
0.16
0.12
0.11
0.12
0.22
Table S2. Bond lengths and binding energies for boron bonds to nitrogen, oxygen and
carbon.
Molecule
PBA
PBE
PBP
NH3
R(Å)
2.950
3.155
---
E(kJ/mol)
-20.9
-22.2
-25.5
N(CH3 )3
R(Å)
2.538
2.694
2.823
E(kJ/mol)
-33.9
-29.7
-29.7
OH
R(Å)
1.5164
1.4601
1.5010
E(kJ/mol)
-256.5
-297.1
-264.4
To understand the effect of boronate ester formation on boron-nitrogen bonding,
we have considered both intermolecular and intramolecular models shown in Figure S1.
Intermolecular Lewis acid:base complexes include the various alkyl (MBA, MBE) and
phenyl boronic acids (PBA, PBE, PBP) with NH3 or (CH3 )3 N (TMA). The
intermolecular adducts can be indicated as TMA:PBA etc. where the boron-nitrogen
interaction is the only covalent interaction between the amine donor and boron Lewis
acid acceptor. The intramolecular models are denoted OTMA -PBA etc. to indicate that
4
the trimethylamino group is in the ortho position to the boronic acid moiety (Figure S1).
The designation for phenyl boronic acid is PBA and that for the glycol and 1,3propanediol ester are PBE and PBP, respectively.
The benchmark calculations provide a basis for evaluating interactions between
phenyl boronic acids or esters and NH3 or N(CH3 )3 . These calculations predict the
relative strength of bonding interactions applicable to mo lecular sensors described in this
proposal. We consider three models phenyl boronic acid (PBA), phenyl boronic ethylene
glycol ester (PBE) and phenyl boronic 1,3-propandiol ester (PBP). The bond lengths and
binding energies for adducts with NH3 and N(CH3 )3 are given in Tables S1 and S2. The
bonding surface for PBP and PBE with NH3 are identical and are represented by the
dotted line. The adduct of each of these with N(CH3 )3 is shown in Figure S3. It is of
interest that the binding energies are significantly lower for adducts of NH3 and N(CH3 )3
with phenyl boronic esters than for adducts with boronic acid and borane. This is also
reflected in the significantly longer bond length for these systems. Although N(CH3 )3
forms a stronger adduct than NH3 with with BH3 , the reverse is true for the phenyl
boronic acid and esters. There is a balance of steric and electronic effects at work. The
increased basicity of N(CH3 )3 over that of NH3 tends to strengthen the adduct. The steric
repulsion of methyl groups with the phenyl ring and ester oxygens tends to weaken the
N(CH3 )3 adduct with PBA, PBE or PBP. However, steric interactions represent a
relatively small contribution and the Lewis acidity of the boron atom and Lewis basicity
of the nitrogen are principally responsible for the boron-nitrogen bonding interaction.
Table S3. The charge on the boron atom as determined by electrostatic potential fitting
(ESP) or Mulliken charge for each of the model phenyl boronic acid/ester adducts.
Molecule
ESP
PBA
PBE
PBP
PBA:TMA0
PBE:TMA0
PBP:TMA0
0.547
0.743
0.732
0.494
0.702
0.726
Mullike
n
0.482
0.567
0.541
0.555
0.633
0.601
Tables S1 – S3 shows that the relative strength of the interaction of TMA with alkyl
boronic acid and boronate ester is opposite to that required for a fluorescent switching
molecule for an alkyl boron substituent (i.e. the B-N bond is stronger for MBA:TMA than
for MBE:TMA ).
The adducts of boric acid B(OH)3 or BA and the glycol ester B(OH)(OCH2 CH2 O)
or BE form Lewis adducts with ammonia that have an interaction energy ~14 kJ/mol.
The interaction energy of ammonia is the same (~14 kJ/mol) with methyl and phenyl
boronic acids and esters. The interaction is somewhat larger for TMA adducts. The
interaction of TMA is strongest for MBA:TMA and MBE:TMA (28.3 and 30.3 kJ/mol),
weaker for BA:TMA and BE:TMA (22.2 and 20.3 kJ/mol) and weaker still in the phenyl
boronic ester adducts PBA:TMA and PBE:TMA (15.5 – 15.9 kJ/mol). The comparison
of Lewis bases NH3 and TMA rules out a role for steric repulsions of the methyl groups
5
in trimethylamine with any of the Lewis acids studied since TMA has a stronger
interaction than ammonia for all of the Lewis acids except BH3 , where steric effects are
smallest. In addition to being a weak interaction the B-N bond formation transition
competes with two conformational coordinates that have similar energetic barriers.
These are denoted as the torsional coordinates τ C-C-B-O and τO-C- C-O in Figure 3 of the text.
In both the intermolecular and intramolecular adducts there is a competition
between boron- nitrogen bonding, d(B-N) and the ring-boron conformations indicated in
o-trimethylamino phenyl boronate glycol ester (τ C-C-B-O in Figure 3 of the text). Table S4
shows that the most stable configuration of a phenyl boronic acid (PBA) or boronate ester
(PBE or PBP) arises when the boron p-orbital has overlap with the benzene π-system.
The energy for each of three molecules PBA, PBE and PBP is the lowest when the
boron-oxygen bonds are nearly coplanar with benzene (τ C-C-B-O = 0o ).
Table S4. Energy of phenyl boronic acid (PBA), phenyl boronic glycol ester (PBE) and
boronic 1,3-propanediol ester (PBP) as a function of Cp-Cp-B-O dihedral angle from 0o
to 90 o .
Cp-Cp-B-O
Dihedral
0
10
20
30
40
50
60
70
80
90
Phenyl boronic acid
(PBA)
Energy
B ESP
(kJ/mol)
Charge
-8128.7
0.556
-8127.8
0.558
-8126.2
0.554
-8123.2
0.559
-8119.5
0.557
-8115.3
0.5
-8111.1
0.6
-8107.3
0.6
-8104.4
0.6
-8104.0
0.6
Phenyl boronic glycol
ester (PBE)
Energy
B ESP
(kJ/mol)
Charge
-10079.7
0.733
-10079.3
0.738
-10077.6
0.755
-10074.7
0.781
-10070.9
0.797
-10067.1
0.811
-10062.9
0.822
-10059.6
0.826
-10057.5
0.8
-10056.2
0.8
Phenyl boronic 1,3propandiol ester (PBP)
Energy
B ESP
(kJ/mol)
Charge
-11362.2
0.723
-11361.7
0.700
-11359.6
0.700
-11357.0
0.800
-11353.3
0.781
-11349.1
0.796
-11344.9
0.808
-11341.2
0.822
-11338.2
0.829
-11337.4
0.835
The electronic coupling matrix element is not likely to be affected by boronate ester
formation
The electron transfer rate constant will be reduced if the distance increases. The
distance dependence of electron transfer has been shown to be exponential and can be
represented V2 = V0 2 exp{-βR} where R is the edge-to-edge distance of the donor and
acceptor. If we consider the formation of a BN +A- charge transfer state, there is a
competition between anthracene and boron for the nitrogen lone pair. Based on the
calculated binding energies as a function of distance (data not shown), the interaction
energy of TMA with anthracene is ~11 kJ/mol, which is only slightly less than the
interaction of the amine with boronate ester. In fact, these values are sufficiently close
that there is a possible change in structure for B-N bond formation that could cause a
change in the edge-to-edge distance R. However, such a conformational switch is
unlikely to provide the robust switching in fluorescence quantum yield that is observed.
6
Thermodynamic equilibria
Binding constants for ammonia and trimethylamine were calculated using DFT
single point energy calculations of donor:acceptor adduct and the individual component
donor and acceptor molecules. There are possible artifacts in these calculations due to
the fact that separate geometry optimizations were carried out for each molecule and
there are possible conformatio nal effects. Nonetheless, some valuable trends are
observed.
Table S5. Interaction energy of Lewis acids consisting of alkyl or phenyl boronic acid or
boronate ester with trimethylamine or ammonia.
Adduct
MBA:TMA
MBA:NH3
MBE:TMA
MBA:NH3
PBA:TMA90
PBA:NH3 90
PBE:TMA90
PBE:NH3 90
PBA:TMA0
PBA:NH3 0
PBE:TMA0
PBE: NH3 0
Energy (kJ/mol)
-22.2
-14.9
-20.3
-14.1
-26.8
5.8
-17.8
24.7
-29.2
-53.3
-27.3
-29.2
Table S6. Energy of hydrolysis reactions for various boronate ethylene glycol ester. A
negative interaction energy indicates that hydrolysis is exothermic and that the ester not
stable. It is remarkable that the trimethylamine moiety provides the least stabilization for
all of the esters considered.
Hydrolysis Reaction
B(OH)C 2 H4 O2 + 2 H2 O à B(OH)3 + C2 H6 O2
B(OH)C 2 H4 O2 :NH3 + 2 H2 O à B(OH)3 :NH3 + C2 H6 O2
B(OH)C 2 H4 O2 :N(CH3 )3 + 2 H2 O à B(OH)3 :N(CH3 )3 + C2 H6 O2
B(CH3 )C 2 H4 O2 + 2 H2O à B(CH3 )(OH)2 + C2 H6O2
B(CH3 )C 2 H4 O2 :NH3 + 2 H2 O à B(CH3 )(OH)2 :NH3 + C2 H6 O
B(CH3 )C 2 H4 O2 :N(CH3 )3 + 2 H2 O à B(CH3 )(OH)2 :N(CH3 )3 + C2 H6 O2
B(φ−0o )C 2 H4O2 + 2 H2O à B(φ−0o )(OH)2 + C2 H6O2
B(φ−0o )C 2 H4O2 :NH3 + 2 H2 O à B(φ−0o )(OH)2 :NH3 + C2 H6 O
B(φ−0o )C 2 H4O2 :N(CH3 )3 + 2 H2 O à B(φ−0o )(OH)2 :N(CH3 )3 + C2 H6O2
B(φ−90o )C 2 H4O2 + 2 H2O à B(φ−90o )(OH)2 + C2 H6O2
B(φ−90o )C 2 H4O2 :NH3 + 2 H2 O à B(φ−90o )(OH)2 :NH3 + C2 H6 O
B(φ−90o )C 2 H4O2 :N(CH3 )3 + 2 H2 O à B(φ−90o )(OH)2 :N(CH3 )3 + C2 H6O2
7
Energy (kJ/mol)
-6.7
3.9
-19.8
1.5
-0.4
0.6
1.1
-0.7
-33.6
17.1
8.1
-1.8
Solvation
Solvation of the cationic and anionic species must be included in order to obtain a
reasonable estimate of the energy of charge transfer reactions. The energy in kJ/mol of
all of the relevant forms of the possible neutral, cations and anions are given in Table 5 of
the text. The analysis was carried out using a vacuum calc ulation and with the inclusion
of continuum solvation (COSMO). The effectiveness of COSMO for the present purpose
was tested using an explicit solvation model for both TMA and Anthracene. This is
shown in Figures S4 and S5.
The explicit model of solvatio n involves calculation of the energy of TMA
relative to the TMA + cation in the presence of a number of water molecules. The
extrapolation to a sea of solvent molecules is consistent with a >200 kJ/mol reduction in
the energy of ionization of TMA (Figure S4). Similarly the energy of Anth relative to
Anth- was calculated in the presence of a number of water molecules as shown in Figure
S5.These calculations show that solvation by water lowers the energy of the charge
transfer state to a range that permits an exothermic electron transfer reaction as required
for a mechanism that involves quenching of anthracene fluorescence. The presence of
boronic acid or boronate also lowers the energy of the amino oxidation reaction TMA à
TMA + + e− .
Figure S4. Calculation of the ionization potential of trimethylamine in a solvent shell
consisting of a discrete number of water molecules. Two calculations are shown. The
calculation of the reorganization energy (filled squares) has the same geometry as neutral
TMA after geometry optimization. The reoptimized geometry (open squares)
corresponds to the solvated form. Note that the energy is lowered systematically by
approximately 185 kJ/mol with respect to the fixed geometry calculation.
8
Figure S5. Calculation of the electron affinity of anthracene in a solvent shell consisting
of a discrete number of water molecules. Two calculations are shown. The calculation
of the reorganization energy (filled squares) has the same geometry as neutral Anth after
geometry optimization. The reoptimized geometry (open squares) corresponds to the
solvated form. Note that the energy is lowered systematically by approximately 40
kJ/mol with respect to the fixed geometry calculation.
It is a challenging task to calculate solvation of a molecule the size anthracene
using a discrete solvation shell. The calculations do suggest that inclusion of a molecular
solvation shell will lead to a lowering of the energy of a dipole created in an electron
transfer reaction. This explicit calculation can be compared to that obtained using
dielectric continuum approach provided by COSMO. In the particular case shown for
TMA and Anth we calculate that the energy of TMA +Anth− dipole is 234 kJ/mol above
the neutral ground state, which shows good agreement with the COSMO calculation
(using a dielectric constant of 78.4 for H2 O). The energy of the lowest energy transition
of anthracene is calculated to 337 kJ/mol (and this approximately in agreement with
experimental values). Both COSMO and explicit solvent predicts that the charge transfer
will occur because the energy of the charge transfer state is ca. 100 kJ/mol lower than the
energy of the excited singlet state. This is in agreement with experiment.
Inner Sphere Reorganization Energy
The reorganization energy can be divided conceptually into inner sphere
(molecular) and outer sphere (solvent) contributions. The inner sphere reorganization
energy is the energy required for distortion along the product potential energy surface
until the equilibrium geometry of the reactant is reached. DFT can be used to calculate
the inner sphere reorganization energy in two ways. First, the overall inner sphere
reorganization energy can be calculated by comparing the energy of a geometry
optimized charged molecule with the energy of the cation or anion calculated in the
geometry of the neutral molecule. Calculations of this type are given in Table S6.
Second, the total reorganization energy can be calculated by imagining a distortion along
9
the energy surface of the charge-separated state until the geometry of the neutral state is
reached.
The inner sphere reorganization energy can be obtained by carrying out a
vibrational frequency calc ulation of the neutral molecule followed by projection of the
eigenvectors obtained from matrix diagonalization of the Hessian matrix onto the
geometry of the cation or anion. Calculations of this type give the specific electron
phonon couplings responsib le for the inner sphere reorganization energy. The distortion
along each mode is considered a contribution to a dimensionless displacement ∆i along
that mode. The ∆i contributions are converted to electron-phonon couplings, S = ∆i2 /2
and then inserted into Eqn. 6 to obtain a molecular or inner sphere contribution. The
results of a calculation of this type are presented in Table S6. One caveat for such a
calculation is the charged and neutral components must be appropriate superposed to
minimize spurious contributions from translation or rotation between the two structures.
The results in Table S7 agree reasonably well with the results obtained in the manuscript
except for the contribution of anthracene, which is found to be very small using the
normal mode approach and substantial using the relative binding energy of the cation in
the two geometries (cation and neutral geometry optimized structures).
The molecular methods based on normal mode analysis would be very difficult to
carry out in a DFT calculation that included a discrete solvent shell. When there are
numerous water molecules the lowest frequency normal modes will be poorly
characterized. We can estimate the outer sphere or solvent component of the
reorganization energy by another method. We geometry optimize the solvent shell and
solute as a neutral species. Once a stable minimum has been found, the cation energy is
calculated. This is equivalent to calculating the vertical energy between the neutral
reactants complex DA and the charge separated products PES, D+ A−.
Table S7. Inner sphere reorganization energy for donors and acceptors involved in
electron transfer quenching of the anthracene excited state fluorescence.
TMA
PBA-OTMA
PBE-OTMA
PBP-OTMA
Anth
Reorganization
Energy (kJ/mol)
42.7
22.2
5.9
15.0
0.42
Model studies based on boronic acid biosensors
Our approach will lead to an understanding of the factors that tune boron-nitrogen
bonding. Our hypothesis is that these interactions are a major determinant of the
molecular recognition of the phenyl boronic acid moiety. In addition the energy levels of
the nitrogen play a role in its ability to quench fluorophores. Study of the energy levels
will also lead to an understanding of the addition fluorescent molecules that may be used
to improve the characteristics of biosensors based on boronic ester formation.
10
Figure S6. The reaction of phenyl boronic acid with a X-Y 1,2-diol or a X-Y-Z 1,3diol to form phenyl boronic ester is shown. The dashed line indicates a putative B-N
interaction that arises due to increased Lewis acidity in the ester form. The ortho- and
para-substituents on the phenyl ring are indicated as R1 and R2 , respectively. In all model
calculations presented here R1 = H and R2 = H.
In phenyl boronic ester biosensors the trimethylamine moiety is covalently found to
the phenyl ring ortho to the boronic acid/ester as shown in Figure S6. Model calculations
of molecules 1, 2 and 3 shown in Figure S6 will modulate the B-N distance as the
molecule rotates about the axis of the phenyl carbon bond. In this case the optimum bond
length is ≈3 Å and the bonding energy is ≈5 kcal/mol. These models reveal a dilemma for
the proposed mechanism of action of the phenyl boronic ester biosensor. The increase in
Lewis acidity of the boron atom upon ester formation has been mentioned as the origin of
the effect on charge transfer. Although the intermolecular models in Figure S1 show a
substantial increase in charge on the boron atom, the increased Lewis acidity of boron
does not result in a change in ground state structure. That is, the bond length for the B-N
bond in models with structure relevant to boronate ester formation (e.g. the reaction in
Figure S6) are virtually identical.
The energetic considerations for electron transfer, which are the main focus of the
text of the manuscript are also inconsistent with B-N bond formation as the cause of the
switching of fluorescence quenching. The energies of the charge transfer states of models
shown as 1, 2 and 3 in Figure S6 are given in Table S8. The corresponding charge
transfer states are depicted in Figures S7 – S9.
11
Table S8. Relative energy of the charge transfer state for the models.
Neutral (kJ/mol )
-12636.9
-14592.5
-15867.0
-4990.7
-15593.3
-8123.2
-10073.8
-11355.8
OTMA-PBA
OTMA-PBE
OTMA-PBP
TMA
Anth
PBA
PBE
PBP
CH 3
Anion(kJ/mol)
-12613.5
-14575.8
-15842.3
NA
-15663.6
-8006.9
-9971.3
-11242.0
CH 3
OH
Cation (kJ/mol )
-12036.1
-14000.9
-15287.1
-4230.4
-14992.1
NA
NA
NA
-
+
OH
B
+
OH
+
B
OH
CH 3
CH 3
dH = (-3583.159-1913.745) - ( 3726.891 - 1947.461 ) = 171.448 kcal/mol
CH 3
CH 3
-
+
O
O
B
+
+
O
B
O
+
CH
3
CH 3
dH = (-3583.159 - 2383.193 ) - ( -3726.891 - 2407.693 ) = 168.234 kcal/mol
CH 3
CH 3
O
O
B
+
-
+
O
+
CH 3
B
O
CH 3
dH = ( -3583.159 - 2686.860) - ( -3726.891 - 2714.125 ) = 170.997 kcal/mol
Figure S7. Phenyl boronic acid (PBA) and boronate esters (PBE and PBP) in charge
transfer states with dimethylanthracene (Anth).
12
_
CH3
+
CH3
OH
OH
B
+
B
OH
OH
+
N
CH3
N
CH3
dH= ( -2876.664 - 3743.724) - (-3020.299- 3726.891) = 126..802kcal/mol
CH3
CH3
O
+
_
O
B
B
O
+
N
CH3
O
+
N
CH3
dH= ( -3346.258 - 3743.724) - (- 3487.732 - 3726.891) = 124.641kcal/mol
CH3
O
CH3
B
_
+
O
O
+
B
+
O
N
CH3
N
CH3
dH= ( -3653.671 - 3743.724) - (-3792.336- 3726.891) = 121..832kcal/mol
_
CH3
+
CH3
+
N
+
CH3
N
CH3
dH= ( -1011.05 - 3743.724) - (-1192.803- 3726.891) = 164.92kcal/mol
Figure S8. O-benzyl dimethamino phenyl boronic acid (OTMA -PBA) and boronate esters
(OTMA-PBE and OTMA -PBP) in charge transfer states with dimethylanthracene (Anth).
The amino group acts as the electron donor in this series. The adduct of the trimethyl
amine (TMA) and Anth is also shown.
13
+
CH3
-
CH3
OH
OH
B
+
B
OH
OH
+
N
CH3
N
CH3
dH = (-3583.159 -3014.717) - (-3726.891 - 3020.299 ) = 149.314 kcal /mol
CH3
CH3
O
-
+
O
B
B
O
+
N
CH3
O
+
N
CH3
dH = (-3583.159 -3483.699) - (-3726.891 - 3487.732 ) = 147.765 kcal /mol
CH3
O
CH3
B
+
CH3
+
O
O
B
+
O
N
CH3
N
dH = (-3583.159 -3786.3577) - (-3726.891 - 3792.336 ) = 149.71 kcal /mol
Figure S9. O-benzyl dimethamino phenyl boronic acid (OTMA -PBA) and boronate esters
(OTMA-PBE and OTMA -PBP) in charge transfer states with dimethylanthracene (Anth).
The boron acts as the electron acceptor in this series.
The unusual boron-nitrogen bond strength of the cation
In the ester form the covalent interaction of boron and nitrogen appears to be
strengthened by oxidation. Table 2 in the text shows that the equilibrium B-N bond
length is shorter for the cation than for the neutral for all dihedral angles considered. The
HOMO of OTMA-PBE has anti-bonding character consisting mostly of the nitrogen lone
pair and an out-of-phase pσ donation by the phenyl ring to the boron. By contrast, the
bond lengths of neutral and cation of the boronic acid molecule (OTMA-PBA) are 1.98
Å and 2.83 Å, respectively (data not shown). The significance of this is not entirely clear
since model calculations that involve the oxidized (+1) form of intermolecular adducts
OTMA:PBA and OTMA:PBE show longer equilibrium bond lengths than the
corresponding neutral species. Both of the cationic forms have bond lengths in the range
from 4.2 Å to 4.3 Å.
The energies of the cationic forms of OTMA:PBA and
OTMA:PBE are both higher than those of the neutral by ~579 kJ/mol, respectively.
Similar conclusions were reached for the methyl boronic acid (MBA) and ester (MBE)
adducts with trimethyl amine (TMA). Thus, it appears that the bonding interaction in the
14
cation OTMA-PBE+ is a specific effect that relies on the short bond in that species. It is
not clear whether such an observation could be responsible for the fluorescent switch.
The conclusions of the remainder of the manuscript suggest that the observation of a
stronger B-N bond the OTMA-PBE+ relative OTMA-PBE is not important for the
mechanism. For this reason the observation is not reported in the manuscript.
Molecular Orbitals for OTMA-PBE
52
15
53
16
54
17
Homo 55 (-5.538 eV)
18
Lumo 56
19
58
20
59
21
60
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23