Effects of Charged Ligand Substituents on the Properties of the Formally Copper(III)-Hydroxide ([CuOH]2+) Unit

Article abstract of DOI:10.1021/acs.inorgchem.8b01529

With the goal of understanding how distal charge influences the properties and hydrogen atom transfer (HAT) reactivity of the [CuOH]²⁺ core proposed to be important in oxidation catalysis, the complexes [M]₃[^SO3LCuOH] (M = [K(18-crown-6)]⁺ or [K(crypt-222)]⁺) and [^NMe3LCuOH]X (X = BAr^F₄^- or ClO₄^-) were prepared, in which SO₃^- or NMe₃⁺ substituents occupy the para positions of the flanking aryl rings of the supporting bis(carboxamide)pyridine ligands. Structural and spectroscopic characterization showed that the [CuOH]⁺ cores in the corresponding complexes were similar, but cyclic voltammetry revealed the E_1/2 value for the [CuOH]²⁺/[CuOH]⁺ couple to be nearly 0.3 V more oxidizing for the [^NMe3LCuOH]²⁺ than the [^SO3LCuOH]^- species, with the latter influenced by interactions between the distal -SO₃^- substituents and K⁺ or Na⁺ counterions. Chemical oxidations of the complexes generated the corresponding [CuOH]²⁺ species as evinced by UV-vis spectroscopy. The rates of HAT reactions of these species with 9,10-dihydroanthracene to yield the corresponding [Cu(OH₂)]²⁺ complexes and anthracene were measured, and the thermodynamics of the processes were evaluated via determination of the bond dissociation enthalpies (BDEs) of the product O-H bonds. The HAT rate for [^SO3LCuOH]^- was found to be 150 times faster than that for [^NMe3LCuOH]²⁺, despite finding approximately the same BDEs for the product O-H bonds. Rationales for these observations and new insights into the roles of supporting ligand attributes on the properties of the [CuOH]²⁺ unit are presented.

Full text of DOI:10.1021/acs.inorgchem.8b01529

Article
pubs.acs.org/IC
Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX
Effects of Charged Ligand Substituents on the Properties of the
Formally Copper(III)-Hydroxide ([CuOH]2+) Unit
Debanjan Dhar,† Gereon M. Yee,† and William B. Tolman*
Department of Chemistry, Washington University in St. Louis, Campus Box 1134, 1 Brookings Drive, St. Louis, Missouri 63130,
United States
S
* Supporting Information
ABSTRACT: With the goal of understanding how distal charge
influences the properties and hydrogen atom transfer (HAT)
reactivity of the [CuOH]2+ core proposed to be important in
oxidation catalysis, the complexes [M]3[SO3LCuOH] (M = [K(18-
crown-6)]+ or [K(crypt-222)]+) and [NMe3LCuOH]X (X = BArF −
4
−
−
+
or ClO4 ) were prepared, in which SO3 or NMe3 substituents
occupy the para positions of the flanking aryl rings of the
supporting bis(carboxamide)pyridine ligands. Structural and
spectroscopic characterization showed that the [CuOH]+ cores
in the corresponding complexes were similar, but cyclic
voltammetry revealed the E1/2 value for the [CuOH]2+/
[CuOH]+ couple to be nearly 0.3 V more oxidizing for the
[
NMe3LCuOH]2+ than the [SO3LCuOH]− species, with the latter
−
influenced by interactions between the distal −SO3 substituents
and K+ or Na+ counterions. Chemical oxidations of the complexes generated the corresponding [CuOH]2+ species as evinced by
UV−vis spectroscopy. The rates of HAT reactions of these species with 9,10-dihydroanthracene to yield the corresponding
[Cu(OH2)]2+ complexes and anthracene were measured, and the thermodynamics of the processes were evaluated via
determination of the bond dissociation enthalpies (BDEs) of the product O−H bonds. The HAT rate for [SO3LCuOH]− was
found to be ∼150 times faster than that for [NMe3LCuOH]2+, despite finding approximately the same BDEs for the product O−
H bonds. Rationales for these observations and new insights into the roles of supporting ligand attributes on the properties of
the [CuOH]2+ unit are presented.
INTRODUCTION
■
Among the myriad high-valent copper−oxygen motifs
implicated with relevance to catalytic oxidations in biological
and synthetic systems,1 the recently characterized mononuclear
[CuOH]2+ core (formally copper(III)-hydroxide) stands out as
an excellent hydrogen atom abstraction reagent from a wide
range of substrates.2−5 The ability of the [CuOH]2+ unit to
activate robust C−H bonds has led to it being postulated as a
potential intermediate in oxygenase enzymes, such as lytic
polysaccharide monooxygenase.3,6 In previous work, the effects
of ligand electronic perturbations on the spectroscopy and
hydrogen atom transfer (HAT) reactivity of the [CuOH]2+
core were examined through comparison of complexes
supported by the ligands shown in Figure 1.4,5 A conclusion
drawn from these studies was that, in general, making the
supporting ligand framework more electron-deficient makes
the resulting copper hydroxide moiety a stronger oxidant but
also a weaker base. These effects offset each other in defining
the O−H bond dissociation enthalpy (BDE) of the Cu(II)-
aqua ([CuOH2]2+) product of HAT, but only partially, such
that the complex with the highest [CuOH]2+/+ redox potential
exhibits the highest HAT thermodynamic driving force (BDE
trend NO2LCu(OH2) (3b) > pyrLCu(OH2) (3a) > pipLCu-
Figure 1. Previously studied hydroxo/aqua complexes supported by
indicated ligands (in box). R = Et or Bu.
Received: June 3, 2018
© XXXX American Chemical Society
A
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(OH2) (3c)). A similar trend in rates of HAT from
dihydroanthracene by the corresponding [CuOH]2+ cores
was observed (NO2LCuOH (2b) > pyrLCuOH (2a) > pipL-
CuOH (2c)), consistent with an Evans−Polanyi correlation, as
seen with many other metal-oxo/hydroxo complexes.4
Scheme 1. Synthesis of Proligand Salts
[(iPr)2EtNH][SO3LH2] (top) and [NMe3LH2][OTf]2
(bottom)
Inspired by the results of the studies performed using ligands
of the same overall charge but differing electron donating
abilities, we sought to explore the effects of varying the overall
charge in the complexes by using ligands with either cationic
+
−
(−NMe3 ) or anionic (−SO3 ) substituents. The installation
of such charged distal residues is intended to mimic the effects
of distally located charged protein residues on the hydrophobic
active site of metalloenzymes and to probe how such
secondary sphere effects alter reactivity patterns within the
active site.7 Additionally, these charged residues would render
the corresponding complexes to be soluble in water, a solvent
of interest for several applications. Such a strategy of
incorporating charged ligand substituents has been particularly
well-explored in the porphyrin literature, often focusing on the
use of cationic groups in the meso positions to give enhanced
redox and HAT reactivity for metal−oxygen intermediates, as
well as imbuing water solubility.8 It is also worth noting that
attachment of cationic alkylammonium groups has significant
effects on the rates of catalytic copolymerization reactions.9
Analogous examples featuring anionic substituents include one
in which sulfonate-substituted phenyl groups at the meso
positions of a heme-iron system influence both the oxidation
potential of the porphyrin ring as well as the pKa values of the
corresponding iron(III)-bisaqua complexes when studied in
water.10 Additional reports implicate the Lewis basicity of the
pendant sulfonate functionalities as key determinants in
altering reactivity by means of secondary coordination sphere
effects.11
presence of the triflate anion in [NMe3LH2][OTf]2 was
confirmed from the 13C{1H} and 19F NMR spectra (Figures
S3 and S4). Also, [(iPr)2EtNH]2[SO3LH2] was characterized
structurally by X-ray crystallography (Figure S11). The X-ray
crystal structure confirmed the overall formulation, with the
acidic N−H proton on the countercation involved in a
hydrogen-bonding interaction with one of the oxygen atoms of
the sulfonate group (N−H−O distance ∼2.8 Å and N−H−O
bond angle 166.6°).12
With the aforementioned studies as precedents, we
hypothesized that charged ligand substituents could have
significant effects on the spectroscopic properties and reactivity
of the [CuOH]2+ core. Thus, we targeted new complexes of
the proligands [NMe3LH2]2+ and [SO3LH2]2− (Scheme 1)
The copper(II) hydroxide precursor complexes
[M]3[SO3LCuOH] (4a: M = [K(18-crown-6)]+, 4b: M =
[K(crypt-222)]+) and [NMe3LCuOH]X (5a: X = BArF −, 5b: X
4
+
−
−
having −NMe3 and −SO3 groups on the para position of
the carboxamide aryl rings. Herein, we describe the successful
preparation and characterization of the proligands, their
precursor complexes with the [CuOH]+ unit, and the products
of their one-electron oxidation comprising the [CuOH]2+ core.
The effects of the charged substituents on the properties of
compounds are explored, leading to new understanding of how
the supporting ligand environment influences the spectro-
scopic features and HAT kinetics and thermodynamics of the
[CuOH]2+ core.
= ClO4 ) were prepared by treatment of the proligand salts
with 1 equiv of Cu(OTf)2 under aqueous conditions in the
presence of excess hydroxide (Scheme 2). In the case of
[K(18-crown-6)]3[SO3LCuOH] (4a), a crude solid was initially
precipitated from the concentrated aqueous reaction mixture
by the addition of acetone; subsequent trituration by
tetrahydrofuran (THF) in the presence of excess 18-crown-6
yielded 4a as a blue solid (0.210 g, 58.0%). Elemental analysis
Scheme 2. Synthesis of Copper(II) Hydroxide Precursor
Complexes
RESULTS
■
Synthesis and Characterization of Ligands and
Copper(II) Complexes. The proligands [SO3LH2]2− and
[
NMe3LH2]2+, isolated as their respective ammonium and
triflate salts, were synthesized as outlined in Scheme 1 (details
in Supporting Information). In both cases, treatment of 2,6-
pyridinedicarbonyl dichloride with 2 equiv of the correspond-
ing aniline (or anilinium salt in the case of [SO3LH2]2−) in the
presence of excess base resulted in the formation of the
dicarboxamide products [(iPr)2EtNH]2[SO3LH2] and NMe2LH2
in gram quantities. Further treatment of NMe2LH2 with 2 equiv
of MeOTf led to the rapid formation of [NMe3LH2][OTf]2.
Both ligand salts were characterized by 1H and 13C{1H} NMR
spectroscopy and elemental analyses (Figures S1−S8). The
B
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and X-ray crystallography (see below) support the indicated
formulation. Further treatment of a solution of 4a in 1,2-
difluorobenzene (DFB) with 5 equiv of crypt-222, followed by
recrystallization by vapor diffusion of Et2O, yielded blue
crystals of [K(crypt-222)]3[SO3LCuOH] (4b, 0.105 g, 87.0%).
The complex [NMe3LCuOH]BArF4 (5a) was synthesized by the
addition of 1 equiv of NaBArF to the initial reaction mixture
4
(proligand, base, Cu(OTf)2), followed by extraction of the
aqueous layer with CH2Cl2 and subsequent removal of the
organic solvent to yield 5a as a light purple solid (0.149 g,
57.9%), the formulation of which was supported by elemental
analysis. Attempts to structurally characterize 5a were
unsuccessful, however, so the variant [NMe3LCuOH]ClO4
(5b) was prepared by slow evaporation of the aqueous
reaction mixture in the presence of excess NaClO4, giving
crystals suitable for X-ray diffraction.
Anticipating that [Cu(OH2)]2+ cores would form as
products of HAT by the [CuOH]2+ complexes of NMe3L2−
and SO3L4−, we sought to prepare these aqua complexes
independently. Success was achieved in the case of the former
ligand, with [NMe3LCu(OH2)](ClO4)2 (6, 0.0784 g, 37.0%)
isolated from the reaction of [NMe3LH2][OTf]2, Cu(OTf)2, and
NaClO4 in H2O (a BArF − salt was not accessible). The
Figure 2. Depictions of the X-ray crystal structures of the anionic
portion of [K(18-crown-6)]3[SO3LCuOH] (4a, top) and cationic
portion of [NMe3LCuOH]ClO4 (5b, bottom). Gray, white, red, blue,
green, and yellow atoms represent C, H, O, N, Cu, and S, respectively.
H atoms are omitted for clarity, and non-hydrogen atoms are depicted
as 50% thermal ellipsoids. Selected interatomic distances (Å) and
angles (deg) are as follows. 4a: Cu1−O3, 1.829(4); Cu1−N1,
1.996(4); Cu1−N2, 1.915(4); Cu1−N3, 2.012(4); N1−Cu1−N3,
159.7(2); N2−Cu1−O3, 179.3(2). 5b: Cu1−O3, 1.862(2); Cu1−
N1, 2.015(2); Cu1−N2, 1.932(2); Cu1−N3, 2.008(2); N1−Cu1−
N3, 159.4(1); N2−Cu1−O3, 176.5(1).
4
complex crystallizes as a hydrate, as determined by X-ray
crystallography and elemental analysis. Similar attempts to
synthesize a [Cu(OH2)]2+ complex for SO3L4− led instead to
the isolation of a dimer [(Et4N)2(SO3LCu)]2 (7, 0.120 g,
56.0%), in which the fourth coordination site of the copper
center was found by X-ray crystallography to be occupied by
one of the pendant sulfonate groups from a neighboring
[
SO3LCu]2− unit (see below).
X-ray Crystal Structures. The X-ray structures of the
[CuOH]+ complexes [M]3[SO3LCuOH] (4a: M = K(18-
crown-6), 4b: M = K(crypt-222)) and [NMe3LCuOH]ClO4
(5b) are depicted in Figure 2 (complexes only) and Figures 3
and 4 (including secondary interactions). Selected bond
metrics are provided in Table 1, along with data for the
previously reported complexes [Bu4N][pyrLCuOH] (1a),
[Et4N][NO2LCuOH] (1b), and [Na][pipLCuOH] (1c) for
comparison.4,14 The geometric parameters for the primary
copper coordination sphere for the set of complexes are
similar, with average Cu1−N1/N3, Cu1−N2, and Cu1−O3
distances of 2.01, 1.93, and 1.85 Å, respectively. The complexes
also exhibit closely analogous square-planar geometries
characterized by similar τ4 values of ∼0.15.13 Thus, despite
differences in remote substituents, overall charge, and/or
outer-sphere interactions (see below), the structural features of
the cores in the entire set of complexes are not significantly
perturbed.
top). Notwithstanding what would be expected to be
coordinative saturation by the multidentate cryptand ligand,
the third [K(crypt-222)]+ cation exhibits a bonding interaction
with one of the pendant SO3− groups on the ligand framework
(Figure 4, bottom). This interaction is indicated by disorder of
the central K+ ion, with a 15% disordered component (K3′)
sitting outside the central cryptand cavity and involved in an
ionic interaction with the SO3− in a similar bidentate fashion as
observed in 4a (K3′−O4, 2.781(15) Å; K3′−O6, 2.842(14)
−
Å). Evidently, the SO3 moiety binds strongly enough to the
K+ ion to cause it to be partially dislodged from the rigid,
tightly coordinating cryptand cavity. As noted below, the
strong tendency for the sulfonate unit to be involved in
secondary interactions with K+ as identified in the crystalline
phase is important in solution as well, with implications for
certain properties and reactivity.
Secondary sphere interactions between the sulfonate and
ligand carboxamide groups and the K+ ion(s) are apparent in
the structures of 4a and 4b (Figures 3 and 4, respectively). In
4a, two K+ centers bind to the oxygen atoms (K···O ≈ 2.8 Å)
of the sulfonate groups (and reside within 18-crown-6
macrocycles),15 with the third bridging between ligand
carboxamide oxygens in adjacent complexes (also within an
18-crown-6 molecule). The latter bridges thus result in the
formation of a coordination network in the crystal (Figure 3,
bottom). Such interactions with K+ highlight the Lewis basic
The X-ray structure of the aqua complex [NMe3LCu(OH2)]-
(ClO4)2 (6) is shown in Figure S12. The cationic unit closely
resembles that in [NMe3LCuOH]ClO4 (5b), with the exception
of a Cu−O distance of 1.927(2) Å, which is 0.065 Å longer
than that in 5b, consistent with an aqua ligand in 6. The aqua
ligand hydrogen atoms appear to be hydrogen-bonded to water
solvate molecule(s), but these are disordered such that details
of the interactions are unclear. Finally, the X-ray structure of
[(Et4N)2(SO3LCu)]2 (7) reveals it to be a dimer with sulfonate
from one unit binding to the Cu(II) ion of the other (Figure
5). The S−O bond distance for the oxygen atom in the
sulfonate that is coordinated to the Cu(II) ion is longer,
measuring 1.483(2) Å as opposed to ∼1.44 Å in the other S−
−
nature of the SO3 functionality. In the case of 4b, the more
effective encapsulation of the K+ ion by the cryptand results in
a structure wherein two of the [K(crypt-222)]+ counterions do
not exhibit any significant secondary interactions (Figure 4,
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Table 1. Comparison of Selected Interatomic Distances (Å)
a
for [CuOH]+ Complexes
b
compound
Cu1−OH Cu1−N2 Cu1−N1/N3
τ4
[K(18-crown-
1.829(4)
1.915(4) 1.996(4),
2.012(4)
0.16
6)]3[SO3LCuOH] (4a)
[K(crypt-
1.848(3)
1.937(3) 2.006(3),
2.012(3)
0.15
222)]3[SO3LCuOH]
(4b)
[
NMe3LCuOH]ClO4 (5b) 1.862(2)
1.932(2) 2.015(2),
2.008(2)
1.924(3) 2.016(3),
2019(3)
1.936(2) 2.005(1),
1.998(1)
2.032(3) 1.958(3),
1.963(3)
0.17
0.16
0.17
0.20
0.18
[Bu4N][pyrLCuOH]
1.845(4)
1.859(2)
1.885(2)
1.891(2)
b
(1a)
[Et4N][NO2LCuOH]
c
(1b)
cd
,
[Na][pipLCuOH] (1c)
2.035(3) 1.959(3),
1.948(3)
a
Estimated standard deviations are indicated in parentheses.
Reference 14. From ref 4. Two molecules in the asymmetric
b
c
d
unit, Na+ is coordinated to the hydroxide.
Figure 3. Additional views of the X-ray crystal structure of [K(18-
crown-6)]3[SO3LCuOH] (4a) with non-hydrogen atoms depicted as
50% thermal ellipsoids, highlighting the interactions with the [K(18-
crown-6)]+ units. Hydrogen atoms are omitted for clarity. Gray,
white, red, blue, green, yellow, and purple atoms represent C, H, O,
N, Cu, S, and K, respectively. (top) View of the three [K(18-crown-
6)]+ units associated with each anion. (bottom) The coordination
polymer network in the solid-state structure depicting the repeating
units (faded) and the key coordinating components (highlighted).
Figure 5. Depiction of the anionic portion of the X-ray crystal
structure of [(Et4N)2(SO3LCu)]2 (7), with non-hydrogen atoms
depicted as 50% thermal ellipsoids and H atoms omitted for clarity.
Gray, blue, red, green, and yellow atoms represent C, N, O, Cu, and S,
respectively. Selected interatomic distances (Å) and angles (deg) are
as follows: Cu1−O3, 1.9445(19); Cu1−N1, 1.928(2); Cu1−N2,
1.995(2); Cu1−N3, 1.995(2); S1−O3, 1.483(2); N2−Cu1−N3,
160.63(10); N1−Cu1−O3, 173.93(10).
cores in [K(18-crown-6)]3[SO3LCuOH] (4a), [NMe3LCuOH]-
BArF (5a), and [Bu4N][pyrLCuOH] (1a) appear to be
4
essentially identical on the basis of UV−vis and electron
paramagnetic (EPR) spectroscopy (Figure 6).4,17 Thus, overlay
of the UV−vis (DFB) and EPR spectra (THF or 1:4 DFB/
THF) for 4a, 5a, and 1a reveals closely similar features (EPR
parameters from simulations provided in Tables S1 and S2).
The similarities are also evident in spectra measured of
aqueous solutions of 4a and 5b (Figures S15 and S16),
although the superhyperfine coupling from the N atoms is
largely unresolved, consistent with the known line-broadening
effects of solvent water.18
Figure 4. Depiction of the X-ray crystal structure of [K(crypt-
222)]3[SO3LCuOH] (4b), showing the anionic portion and the
[K(crypt-222)]+ unit that interacts with the sulfonate group. Gray,
white, red, blue, green, yellow, and purple atoms represent C, H, O,
N, Cu, S, and K, respectively. Hydrogen atoms are omitted for clarity,
and the non-hydrogen atoms are shown as 50% thermal ellipsoids.
Both disordered K+ centers are shown (occupancies: K3, 85%; K3′,
15%).
−
+
Cyclic Voltammetry. While the para-SO3 and -NMe3
substituents have little effect on the UV−vis and EPR
spectroscopic features of [K(18-crown-6)]3[SO3LCuOH] (4a)
and [NMe3LCuOH]BArF (5a), the charged residues dramat-
4
ically influence their electrochemical behavior. Cyclic voltam-
metry measurements made in DFB solvent19 (Figure 7, Table
2) revealed quasi-reversible single-electron redox events with
E1/2 values of −0.135 and 0.140 V (vs Fc+/Fc) for [K(18-
O bonds in the sulfonate. Such asymmetry in coordinated
sulfonates has been reported previously in the literature.16
UV−Vis and EPR Spectroscopy. Despite their different
overall charges, the electronic structures of the copper(II)
crown-6)]3[SO3LCuOH] (4a) and [NMe3LCuOH]BArF (5a),
4
respectively. The shown cyclic voltammogram (CV) of the
parent complex [Bu4N][pyrLCuOH] (1a) (E1/2 = −0.076 V),
D
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Figure 7. CVs of [K(18-crown-6)]3[SO3LCuOH] (4a) (blue),
[
NMe3LCuOH]BArF (5a) (red), and [Bu4N][pyrLCuOH] (1a)
4
(black dotted). Conditions: 2 mM analyte in DFB solvent, 0.2 M
Bu4NPF6 (0.2 M Bu4NBArF in case of 5a; data for this electrolyte
4
provided in Figures S9 and S10).
Table 2. E1/2 Values for the [CuOH]2+/[CuOH]+ Redox
a
Couples of 4a, 5a, and 1a−c
complex
E1/2 (V vs Fc+/Fc)
[K(18-crown-6)]3[SO3LCuOH] (4a)
−0.135
0.140
−0.076
0.124
[
NMe3LCuOH]BArF ] (5a)
4
b
[NBu4][pyrLCuOH] (1a)
[NBu4][NO2LCuOH] (1b)
b
b
[NBu4][pipLCuOH] (1c)
−0.260
a
Conditions: 2 mM analyte in DFB solvent, 0.2 M Bu4NPF6 (0.2 M
b
Bu4NBArF in case of 5a). From ref 3.
4
4b. To probe such effects, 100 equiv of [Na(15-crown-5)]OTf
were added to solutions of [K(18-crown-6)]3[SO3LCuOH]
(4a) in DFB, which were analyzed by UV−vis and EPR
spectroscopy, as well as cyclic voltammetry. The addition of
[Na(15-crown-5)]OTf had little influence on the UV−vis and
EPR spectra of 4a, with virtually no changes to the UV−vis
spectrum (Figure S17) and only subtle shifts in the g∥
component of the EPR spectrum (Figure 8, top; Table S2
and Figure S14). In contrast, the electrochemical differences
were stark, with the CV showing a large (∼180 mV) positive
shift in the E1/2 for the [CuOH]2+/[CuOH]+ couple from
−0.135 to 0.043 V (vs Fc+/Fc; Figure 8, bottom). A similar
shift was also observed upon addition of [Na(15-crown-
5)]OTf to [K(crypt-222)]3[SO3LCuOH] (4b) (Figures S13
and S19; the CVs for 4a and 4b are essentially identical in
water; see Figure S18). The magnitude of this shift to a higher
potential by mere alteration of the counterion is comparable to
the difference in the redox potentials between [Bu4N]-
Figure 6. (top) Overlay of UV−vis spectra for 4a, 5a, and 1a taken in
DFB at 298 K. (bottom) Overlay of EPR spectra of 4a (1:4 by volume
DFB/THF), 5a (THF), and 1a (THF) at 30 K.
in combination with the previously reported data for
[Et4N][NO2LCuOH] (1b; E1/2 = 0.124 V) and [Bu4N]-
[pipLCuOH] (1c; E1/2 = −0.260 V),3,4 serve as precedent for
assignment of the waves for 4a and 5a as analogous
[CuOH]2+/[CuOH]+ couples. As expected, the electron-
+
withdrawing properties of the NMe3 substituents in 5a (σp
= 0.88) result in a positively shifted redox potential, slightly
larger than that seen for 1b (NO2 groups σp = 0.78).20 On the
−
contrary, while the −SO3 substituents in 4a are slightly
electron-withdrawing in nature (σp = 0.09),20 the redox
potential is ∼60 mV lower than that of the parent complex
[Bu4N][pyrLCuOH] (1a). We surmise that this result reflects
the overall trianionic nature of 4a, which renders the one-
electron oxidation thermodynamically more favorable.21
Effect of Counterions on the Spectroscopic and Electro-
chemical Properties of [M]3[SO3LCuOH] (4a or 4b). In view of
the secondary interactions between the K+ ions and the
sulfonate groups observed by X-ray crystallography, we sought
to understand the role of these interactions on the solution-
state spectroscopic and/or electrochemical properties of 4a or
[pyrLCuOH] (1a) and [Et4N][NO2LCuOH] (1b) (ΔE1/2
≈
200 mV).4 We interpret the results to suggest that there is a
significant difference between the interactions of the pendant
−SO3 groups in 4a/4b with the K+-containing counterions
−
compared to [Na(15-crown-5)]+. While this difference causes
significant changes in the redox potential, the UV−vis and EPR
E
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−
−SO3 group interacts more strongly with the more Lewis
acidic Na+ ion22 and exists as a “tighter” ion pair, thus
mitigating the electrostatic effect of the anionic −SO3− groups
(i.e., effectively decreasing the overall charge) and leading to
an anodic shift in the [CuOH]2+/[CuOH]+ redox potential.
Characterization of [CuOH]2+ Complexes. Treatment of
solutions of [K(18-crown-6)]3[SO3LCuOH] (4a) and
[
NMe3LCuOH]BArF (5a) in DFB with 1 equiv of oxidant
4
([Fc]BArF4 for 4a; [AcFc]BArF4 for 5a) at −25 °C resulted in
the immediate formation of intense features in the correspond-
ing UV−vis spectra (Scheme 3, Figure 9).23 By analogy to
Scheme 3. Oxidation Reactions of 4a to 8a and 5a to 9a
Figure 9. Overlay of UV−vis spectra. Red (solid line): [K(18-crown-
6)]2[SO3LCuOH] (8a); blue: [NMe3LCuOH](BArF )2 (9a); black
4
(dotted): NO2LCuOH (2b); red (dashed): [K(18-crown-
6)]2[SO3LCuOH] + 100 equiv of [Na(15-crown-5)]OTf (8a′).
results of previous studies, we attribute these features to ligand
2+
2
2
aryl π to Cu dx −y charge transfer transitions of the [CuOH]
Figure 8. (top) Overlay of EPR spectra of [K(18-crown-
6)]3[SO3LCuOH] (4a) both in the absence (black) and presence
(blue) of 100 equiv of [Na(15-crown-5)]OTf. The dashed line shows
the subtle difference in the gz component (1 mM concentration of 4a
in DFB). (bottom) Overlay of CVs for 4a in DFB in the absence
(black) and presence (blue) of 100 equiv of [Na(15-crown-5)]OTf.
Conditions: 2 mM analyte in DFB solvent, 0.2 M Bu4NPF6, 50 mV/s.
cores in the product complexes [K(18-crown-
6)]2[SO3LCuOH] (8a) and [NMe3LCuOH](BArF )2 (9a),
4
respectively. The broad absorption feature for [K(18-crown-
6)]2[SO3LCuOH] (8a; solid red line, Figure 8) at ∼610 nm (ε
= 12 500 M−1 cm−1) is significantly red-shifted relative to the
sharper feature for [NMe3LCuOH](BArF )2 (9a; solid blue line)
4
centered around ∼498 nm (ε = 14 000 M−1 cm−1). The latter
spectrum is quite similar to that reported previously for
NO2LCuOH (2b; dashed black line) (λmax (ε) = 513 nm
(15 000 M−1 cm−1)), with respect to both peak shape and
position, attesting to similar electron-withdrawing effects of the
-NO2 and -NMe3+ substituents.4 For the oxidation reactions of
spectra are largely unaffected, reflecting little perturbation to
the electronic structure of the copper ion itself. To rationalize
the more positive E1/2 in the presence of [Na(15-crown-
5)]OTf, we hypothesize that in DFB solvent the Lewis basic
F
DOI: 10.1021/acs.inorgchem.8b01529
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Inorganic Chemistry
Article
4a and 5a, titration experiments revealed a 1:1 reaction
stoichiometry (Figures S20 and S21), confirming that
oxidation of the copper(II) hydroxide precursor proceeds
cleanly via a one-electron process. The absorption bands for
the oxidized products decay fully within 2−3 h at −25 °C and
immediately upon warming. Additionally, the rates of decay at
lower temperatures are enhanced upon the addition of external
substrates containing C−H bonds that can be activated (see
below). These observations are consistent with prior character-
ization of analogous [CuOH]2+ complexes.3,4
Scheme 4. Equilibrium Measured to Determine the ΔBDE
between [Cu(OH2)]2+ Cores
BDE (kcal mol−1) = 1.37 pKa + 23.06 E° + CH
(1)
Thus, [K(18-crown-6)]2[SO3LCuOH] (8a) was generated
by reaction of [K(18-crown-6)]3[SO3LCuOH] (4a) with
In line with the electrochemical measurements discussed
above, the chemical oxidation of [K(18-crown-
6)]3[SO3LCuOH] (4a) also was found to be sensitive to the
nature of the countercation. In the presence of [Na(15-crown-
5)]OTf (100 equiv), the reaction of 4a with [Fc]BArF4 did not
proceed to give 8a as shown above. Conversion to 8a′ instead
[Fc]BArF (1 equiv) in DFB at −25 °C and then treated
4
with varying amounts of the parent aqua complex pyrLCu-
(OH2) (3a, 0.25−2.50 equiv).2 Monitoring this reaction by
UV−vis spectroscopy showed decay of 8a and concomitant
conversion of aqua complex 3a to pyrLCuOH (2a) until
equilibrium was reached, as evinced by cessation of change in
the UV−vis spectrum (Figure S24). Deconvolution of the
equilibrium spectrum by multicomponent fitting enabled
required the use of the stronger oxidant [AcFc]BArF (E1/2
≈
4
0.27 V vs Fc+/Fc),24 with the product exhibiting a UV−vis
spectrum with slightly different features from those of 8a
(Figure 9, Figure S22). Interestingly, the analogous spectrum
can be generated by treating a solution of 8a (generated using
determination of the equilibrium constant Keq = 4.2
1.8
(illustrative deconvolution traces shown in Figure S25),
corresponding to a ΔBDE of ∼0.8 kcal mol−1, such that the
O−H BDE for [K(18-crown-6)]2[SO3LCu(OH2)] (10a) is
∼0.8 kcal mol−1 stronger than that for 3a (90 3 kcal mol−1;
Table 3).3
[AcFc]BArF ) with an excess (>5.0 equiv) of [Na(15-crown-
4
5)]OTf (Figure S23). This behavior is consistent with the
observed shift in E1/2 in the presence of [Na(15-crown-5)]OTf
(Figure 8, bottom). We speculate that the subtle difference
between the observed ligand-to-metal charge transfer (LMCT)
features in the presence or absence of [Na(15-crown-5)]OTf is
a consequence of different interactions between the pendant
sulfonates and the Na+/K+ counterions, which result in subtle
electronic perturbations to the highest occupied molecular
orbital (HOMO) known to have significant ligand π
character.2,4 The larger red shifts of the UV−vis features for
[M(18-crown-6)]2[SO3LCuOH] (8a, M = Na or K) relative to
Table 3. BDE Values for [CuOH2]2+ Complexes
supporting ligand
SO3L4−
solvent
BDE (kcal/mol)
ref
a
DFB
H2O
DFB
H2O
THF
DFB
THF
∼91
this work
∼87
∼91.5
this work
this work
a
NMe3L
∼91
this work
pyrL2−
NO2L2−
pipL2−
90
91
88
3
3
3
4
4
4
those of [NMe3LCuOH](BArF )2 (9a) also are consistent with
4
the expected effects on the ligand-centered HOMO (more
+
electron-withdrawing -NMe3 substituent increases energy of
a
2
2
Approximate value, see text.
ligand π → Cu dx −y LMCT transition).
Determination of O−H BDEs. As in the previous work,4
we sought to understand the influence of the ligand-based
electronic perturbations on the HAT reactivity of the new
[CuOH]2+ complexes. A measure of the thermodynamic
driving force for abstraction of H atoms by the complexes with
[CuOH]2+ cores is the bond dissociation enthalpy of the O−H
bond in the product [Cu(OH2)]2+ unit.3 In previous work, two
methods for determining this value for the products 3a−c
(Figure 1) of the HAT reactions of 2a−c were used.3,4 One
method involved use of a thermodynamic square scheme,
whereby values for the [CuOH]2+/1+ reduction potential and
the pKa for the [Cu(OH2)]2+ complex were inserted into eq 1
(where CH is associated with the H+/H· standard reduction
potential in the solvent used, which was THF). A second
method (“cross-HAT”) involved mixing of a [CuOH]2+
complex supported by one ligand (L) with the [Cu(OH2)]2+
complex of another (L′) and measuring the Keq for the
resulting equilibrium expressed by Scheme 4. The difference in
the BDE values (ΔBDE) between the two [Cu(OH2)]2+ cores
may then be approximated by −RTlnKeq, under the
assumption that ΔSrxn is negligible and ΔHrxn is independent
of temperature. We considered both methods in seeking to
determine the BDE values for the [Cu(OH2)]2+ cores
supported by NMe3L and SO3L4− and to compare them to the
values previously determined for 3a−c.
Unfortunately, the cross-HAT protocol proved unsuitable
for determining the O−H BDE of [NMe3LCu(OH2)](BArF )2
4
(6a), the product of HAT by [NMe3LCuOH](BArF )2 (9a),
4
owing to difficulties in reproducibly determining the point of
equilibrium. However, addition of an excess of [pyrLCu(OH2)]
(3a, 2.5 equiv) to a solution of 9a (generated by oxidation of
5a with 1 equiv of [AcFc]BArF in DFB solvent at −25 °C)
4
resulted in the complete conversion of 3a to pyrLCuOH (2a)
within ∼2 min (Figure S26). Assuming ∼95% conversion of 3a
of 2a, we estimate the Keq for this reaction to be ∼11, which
gives a ΔBDE between 3a and 6a of ∼1.5 kcal mol−1, such that
the O−H bond of the latter is stronger than that of the former.
Thus, the O−H BDE for complex 6a is estimated to be ∼91.5
kcal mol−1 (Table 3).
To further validate the results from the cross-HAT
experiments, we took advantage of the water solubility of
[K(18-crown-6)]3[SO3LCuOH] (4a) and [NMe3LCuOH]ClO4
(5b) to more directly evaluate the BDE of the O−H bonds in
[K(18-crown-6)]2[SO3LCu(OH2)] (10a) and [NMe3LCu-
(OH2)](ClO4)2 (6b) using eq 1, where CH = 55.8 kcal
mol−1.25 The pKa of 10a was determined through UV−vis
titrations of 4a dissolved in water, by varying the pH between
4.75 and 11.35, where the pH was adjusted by using either
KOH or HClO4 and measured using a pH probe (Figure S27,
G
DOI: 10.1021/acs.inorgchem.8b01529
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Inorganic Chemistry
Article
Table 4. Thermodynamic and Kinetic Data for HAT Reactions of [CuOH]2+ Complexes with 9,10-Dihydroanthracene (DHA)
and 1,4-Cyclohexadiene (CHD) at −25 °C in DFB
a
a
compound
ligand
ΔBDEDHA (kcal mol−1
)
kDHA (M−1 s−1
)
kH/kD (DHA)
kCHD (M−1 s−1
)
ΔBDECHD (kcal mol−1
)
8a
SO3L4−
SO3L4−
NMe3L
pyrL2−
NO2L2−
pipL2−
14.7
14.7
14.7
13.7
14.7
11.7
400(22)
500(30)
2.8(7)
12
70(8)
90(9)
28(2)
15
15
15
14
15
12
b
8a′
9a
−
18
c
d
e
f
2a
50(8)
27
38(2)
d
e
2b
346(50)
10
−
−
d
e
2c
25(4)
17
a
Calculated as the difference between the O−H BDE of the product [Cu(OH2)]2+ complex and the BDE of the substrate C−H bond (taken from
b
c
ref 30). 100 eq. of [Na(15-crown-5)]OTf added. For this value kD is taken as the average of the kD values measured at 243 and 253 K (see Figure
d
e
S33). From ref 4 unless otherwise stated. Extrapolated values, determined from activation parameters reported in ref 4 measured in CH2Cl2.
f
From ref 31.
left).26 The intermediate equilibrium spectra were deconvo-
luted into the proportions of 10a and 4a via multicomponent
fitting, using the two-component spectra measured at pH 4.75
(10a) and pH 11.35 (4a).27 With these proportions and the
measured pH of the solution at the equilibrium points, the pKa
of 10a was found to be 10.1 0.2. In a similar fashion, the
corresponding titration of 5b in water also was followed by
UV−vis spectroscopy and deconvoluted to obtain the pKa of
organic product of the HAT reaction. The decays of the
LMCT band were modeled with a single exponential function
to obtain pseudo-first-order rate constants (kobs), and plots of
these values versus substrate concentration (Figures S30−S33)
yielded the second-order HAT rate constants (k). For
comparison, we tabulate these rate constants and the
corresponding kinetic isotope effect (KIE) values determined
by using DHA-d4 in Table 4 with those reported previously for
the other [CuOH]2+ complexes 2a−c.
[
NMe3LCu(OH2)](ClO4)2 (6b) as 8.0
0.2 (Figure S27,
right). Alternatively, for 5b a standard pH titration curve was
generated by measuring the pH with a pH probe (Figure S28),
and a pKa value of ∼8.3 was obtained for 6b, which is in
reasonable agreement with the result obtained from the UV−
vis titration.
Determination of the E1/2 values in water for the one-
electron redox potentials corresponding to the 8a/4a and 9b/
5b redox couples was achieved through cyclic voltammetry
experiments performed at high pH to ensure full conversion to
the hydroxide form during the experiment. This proved
difficult in the case of the 8a/4a couple, owing to the
significant overlap between the oxidative feature and the
apparent onset of catalytic water oxidation under basic
conditions.28 Nevertheless, an approximate E1/2 = 0.75 V (vs
normal hydrogen electrode (NHE)) was determined (Figure
S29, left) at pH = 11. In the case of the 9b/5b redox couple, an
Comparison of the second-order rate constants shows
[K(18-crown-6)]2[SO3LCuOH] (8a) to react most rapidly
with DHA, with a rate constant similar to that measured for 2b
supported by NO2L2− and in line with the strong O−H BDE
measured for 8a.4 Notably, the second-order rate constant
measured for reaction with [NMe3LCuOH](BArF )2 (9a) is
4
∼150 times less than that for 8a, despite having a similar
thermodynamic driving force. It is also noteworthy that both
8a and 9a exhibit large kinetic isotope effects of 12 and 18 (at
−25 °C; Figures S32 and S33). These large values are
consistent with rate-determining C−H bond cleavage and are
similar to those reported previously for the analogous
[CuOH]2+ complexes 2a−c.2,4
In assessing the possible reasons for the significantly slower
reaction of 9a with DHA, we noted that the 2+ charge of the
complex is balanced by the two large BArF − counterions,
4
E
1/2 = 1.03 V (vs NHE) was measured at pH = 14 (Figure S29,
which might crowd the [CuOH]2+ core and hinder reactivity
with a sterically encumbered substrate such as DHA. To probe
this hypothesis, we measured the second-order rate constant
for reactions of 8a and 9a with 1,4-cyclohexadiene (CHD), a
smaller hydrogen atom donor with a comparable C−H BDE to
right). Thus, with these values, eq 1 gave O−H BDE values of
∼87 and ∼91 kcal mol−1 for 10a and 6a/b, respectively (Table
3). While the BDE value for 10a is slightly lower than the BDE
value of ∼91 kcal mol−1 determined from the nonaqueous
cross-HAT measurements, it is well-within the range of O−H
BDE values that have been seen thus far for [Cu(OH2)]2+
complexes.3,4,29 Importantly, the trend in O−H BDE values
measured in both organic and aqueous media matches (6a/b >
10a).
that of DHA (76.0
1.2 and 76.3, respectively), which
typically undergoes HAT reactions at similar rates.30,31 Indeed,
for 9a the reactions with CHD were found to be ∼10 times
faster than for DHA, with k = 28(2) M−1 s−1 (Figure S34,
right). This value is quite similar to what was previously
measured for 2a (k = 38 M−1 s−1) under analogous
conditions.31 The value determined for 8a is also similar
(70(8) M−1 s−1; Figure S33), and it is smaller than the rate
constant for its reaction with DHA. Taken together, the
evidence supports the notion that the slow rate of reaction of
9a derives from steric hindrance by the bulky counterions.32
Inspired by the observations that the nature of the
counterion had a considerable effect on the redox potential
for the [CuOH]2+/[CuOH]+ couple for 8a, we investigated
the influence of counterions on its HAT reactivity. The
reaction between DHA and 8a was repeated in the presence of
100 equiv of [Na(15-crown-5)]OTf and yielded a slightly
larger second-order rate constant (k) of 500(30) M−1 s−1
HAT Reaction Kinetics. To probe how the distal-charged
substituents alter the HAT reaction kinetics of [K(18-crown-
6)]2[SO3LCuOH] (8a) and [NMe3LCuOH](BArF )2 (9a), we
4
measured the rates for reactions of these complexes with 9,10-
dihydroanthracene (DHA) and/or DHA-d4 in DFB solvent at
−25 °C. In the experiments, solutions of 4a or 5a in the
presence of excess substrate were treated with the appropriate
oxidant, which immediately yielded the diagnostic charge-
transfer features associated with the formation of the
corresponding [CuOH]2+ species. This charge-transfer feature
decayed over time, concomitant with the appearance of sharp
absorption features in the 360−380 nm range, indicative of the
formation of anthracene (Figures S30 and S31), the expected
H
DOI: 10.1021/acs.inorgchem.8b01529
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Inorganic Chemistry
Article
(Figure 10, Table 4). Concerned that changes in the ionic
DISCUSSION
■
strength might affect the reaction rate,33 we performed a
Using two new derivatives of the pyridine-dicarboxamide
ligand L2− having charged −SO3− or −NMe3+ substituents, we
prepared and characterized the Cu(II) hydroxide ([CuOH]+)
complexes [M]3[SO3LCuOH] (4a: M = [K(18-crown-6)]+, 4b:
M = [K(crypt-222)]+) and [NMe3LCuOH]X (5a: X = BArF −
,
4
−
5b: X = ClO4 ). A key question we wished to address was how
the charged substituents influenced the properties of the
[CuOH]+ cores and whether the nature of the electronic
perturbation (resonance vs field effect) was an important
factor. The X-ray structures showed little variation in the
Cu(II) coordination geometries as a function of para
substituent, but the presence of significant interactions
between the −SO3 groups and the K+ ions in the solid-state
−
structures of 4a and 4b is an important feature in these
complexes that has implications for their properties in solution.
The tendency of the −SO3− units to act as ligands was further
illustrated in the X-ray structure of [(Et4N)2(SO3LCu)]2 (7),
which was isolated in attempts to prepare a [Cu(OH2)]2+
complex. UV−Vis and EPR spectra of the [CuOH]+ complexes
4 and 5 were closely analogous, indicative of similar electronic
structures for the tetragonal Cu(II) compounds despite their
different overall charges. In contrast, considerable differences
in the redox potentials for the [CuOH]2+/[CuOH]+ couples
were observed, which varied as a function of the supporting
−
ligand substituents and, for the −SO3 system, the nature of
the countercation (Table 2). A plot of E1/2 value as a function
of Hammett parameter (σp) for the complexes supported by
the pyridine-containing ligands shows a reasonable correlation
(slope = +0.26) for the complexes bearing the pyrL2−, NO2L2−,
and NMe3L ligands, but those featuring the SO3L2− ligand with
either K+ or Na+ counterions fall off the line (Figure 11a). We
attribute this observation to the use of a single σp value for the
−SO3 group (reported from measurements in water)20 that
−
does not take into consideration the effects of interactions with
countercations in aprotic organic solvents. That is, differences
in the interactions with Na+ and K+ evidently induce disparities
in the E1/2 values by modulating the electron-withdrawing
−
capabilities of the −SO3 group.
As noted previously4 and reproduced in Figure 11b, the Az
value (which reflects spin localization on the Cu nucleus)
varies linearly with E1/2 value for the complexes supported by
pipL2−, pyrL2−, and NO2L2− (black circles). The data for the new
ligands with −NMe3+ and −SO3− substituents do not fit to this
line, however. As noted above, the disparity for the SO3L4−
systems with Na+ and K+ countercations is likely due to ionic
interactions that modulate the electron-withdrawing properties
Figure 10. (top) Kinetics plots for the reaction between 8a and DHA
in the presence of 100 equiv of [Na(15-crown-5)]OTf (blue), 100
equiv of Bu4NPF6 (red), or no added electrolyte (black) in DFB at
−25 °C. (bottom) Kinetics plots for the reaction between 9a and
DHA in the absence (black) and presence (red) of 100 equiv of
Bu4NBArF in DFB at −25 °C.
4
−
of the −SO3 moiety. A different explanation is needed to
rationalize the significant difference in Az values for the NMe3L
system compared to that comprising NO2L2− that has similarly
electron-withdrawing −NO2 groups (with a similar σp value).
We postulate different sensitivity of Az to resonance and field
effects that contribute to σp, as it is known that the −NO2 and
−NMe3+ groups differ with respect to these effects.35 Thus, we
hypothesize that electron withdrawal via resonance is a
significant contributor to a decreased Az value and that the
lack of resonance contributions for the −NMe3+ groups results
in the higher Az for the complex supported by NMe3L.
control reaction of 8a with DHA in the presence of 100 equiv
of [Bu4N]PF6 (Figure 10). The observed rate constant in this
case was found to be 300(12) M−1 s−1, slightly lower than in
the absence of added electrolyte, and consistent with the
presence of ion pairing in 8a.34 Moreover, analogous reactions
with 9a performed in the presence of excess [Bu4N]BArF
yielded no difference in the rate of HAT. Taken together, these
data provide firm evidence for an intrinsic difference in the
influence of the [Na(15-crown-5)]+ and [K(18-crown-6)]+
counterions on the HAT properties of the [CuOH]2+ core,
presumably resulting from divergent interactions between the
Na+ and K+ ions and the sulfonate groups.
−
With the goal of evaluating how the charged −SO3 and
+
−NMe3 substituents influence the properties and reactivities
of the [CuOH]2+ core, the products of the one-electron
chemical oxidation of [K(18-crown-6)]3[SO3LCuOH] (4a) and
I
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With the aim of understanding the HAT reactivity of the
new [CuOH]2+ complexes 8a and 9a, the BDE of the O−H
bond in the product aqua complexes was evaluated (Table 3).
The O−H BDEs for [K(18-crown-6)]2[SO3LCu(OH2)] (10a)
and [NMe3LCu(OH2)](BArF )2 (6a) measured in DFB are
4
slightly higher than in the parent complex pyrLCu(OH2) (3a)
by ∼0.8 and 1.5 kcal mol−1, respectively. In addition, for the
first time, thermodynamic parameters also were obtained in
water, which was enabled by the aqueous solubility of the
charged complexes. Thus, values of pKa ≈ 8.0 for [NMe3LCu-
(OH2)](ClO4)2 (6b) and E1/2 = 1.03 V (vs NHE) for the
[
NMe3LCuOH](ClO4)2 (9b)/[NMe3LCuOH]ClO4 (5b) couple
yielded an O−H BDE of ∼91 kcal mol−1 for 6. A higher pKa of
10.1 0.2 for [K(18-crown-6)]2[SO3LCu(OH2)] (10a) and a
lower E1/2 = 0.75 V (vs NHE) for the [K(18-crown-
6)]2[SO3LCuOH] (8a)/[K(18-crown-6)]3[SO3LCuOH] (4a)
couple resulted in a lower BDE of ∼87 kcal mol−1 for 10a.
The close agreement of the BDE values determined in organic
and aqueous media and with previous values determined in
organic solvent for complexes supported by pyrL2−, pipL2−, and
NO2L2− is notable.25 Importantly, the complex with higher E1/2
supported by NMe3L has a BDE larger than that supported by
SO3L4−, consistent with redox potential slightly overwhelming
the compensation of pKa in determining the BDE (eq 1).4
On the basis of the thermodynamics for HAT, we
anticipated the HAT rates for the [CuOH]2+ core supported
by NMe3L to be greater than those supported by SO3L4−.
However, the second-order rate constant for the reaction of
[
NMe3LCuOH](BArF )2 (9a) with DHA (DFB, −25 °C) was
4
∼ 1 6 0 t i m e s s m a l l e r t h a n t h a t o f
[K(18-crown-6)]2[SO3LCuOH] (8a) (Table 4). In addition,
the rate constant further increased for 8a in the presence of
excess [Na(15-crown-5)]OTf, an effect that was found via
control experiments to not be correlated with an increase in
ionic strength. We posit that the increased interaction of
[Na(15-crown-5)]+ with the anionic −SO3 groups, partic-
−
ularly in an aprotic solvent such as DFB, causes it to exist as a
“tighter ion pair”, which mitigates the electrostatic effects of
−
the anionic −SO3 moiety to a greater extent than in the
presence of [K(18-crown-6)]+. This effect is also manifested in
the change in the reduction potential of the [CuOH]2+/
[CuOH]+ couple by ∼180 mV, which effectively makes 8a′ a
stronger oxidant. This intrinsic increase in the oxidation
potential might translate into a higher thermodynamic driving
force for the HAT reaction, which would cause the observed
small increase in the rate of HAT from DHA (and CHD).
Using similar thermodynamic arguments, we expected the
Figure 11. (a) Plot of E1/2 (V vs Fc+/Fc measured in DFB) vs σp for
[CuOH]2+/[CuOH]+ couples supported by the indicated ligands.
The line is a fit to the black data points (slope = +0.26). (b) Plot of
the hyperfine splitting constant Az for the [CuOH]+ complexes
supported by the indicated ligands. The line is a fit to the black data
points, reproduced from ref 4.
rate constant for HAT from DHA by [NMe3LCuOH](BArF )2
4
[
NMe3LCuOH]BArF (5a) were studied. Consistent with the
4
(9a) to be similar or perhaps even greater than that of
electrochemistry results, [Fc]BArF oxidized 4a, but the
[
NO2LCuOH] (2b), as they have essentially identical
4
stronger oxidant [AcFc]BArF was needed to convert 5a.
ΔBDEDHA values (enthalpic driving force, Table 3), and 9a
has a slightly higher E1/2 value. Yet, the HAT rate constant for
9a is more than 100 times smaller than for 2b, and it is ∼20
times smaller than that for the parent system [pyrLCuOH]
(2a). The discrepancy is illustrated in a plot of log k vs O−H
BDE for the series of [CuOH]2+ complexes (Figure 12). We
postulate that the slow rate for 9a is due to enhanced steric
4
The oxidation products [K(18-crown-6)]2[SO3LCuOH] (8a)
and [NMe3LCuOH](BArF )2 (9a), respectively, were identified
4
on the basis of the similarity of their intense UV−vis features
2
2
to the signature ligand aryl-π to Cu dx −y LMCT features of
the [CuOH]2+ complexes supported by pyrL2−, pipL2−, and
NO2L2− defined previously.2,4,36 Also consistent with the one-
electron oxidation reactions to yield 8a and 9a were (a) the
results of titration experiments that indicated a 1:1
stoichiometry and (b) their second-order HAT reactions
with DHA or CHD to yield anthracene or benzene and the
respective [Cu(OH2)]2+ complexes, identified for the system
supported by NMe3L via independent synthesis.
hindrance by the associated BArF − counteranions. Consistent
4
with such a steric argument, the rate constant for HAT by 9a
from CHD, which is significantly smaller than DHA but with a
similar C−H BDE, is enhanced by ∼10-fold relative to that for
DHA. The likely presence of tunneling contributions to the
HAT reaction rate complicates the analysis, however, as noted
J
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Inorganic Chemistry
Article
secondary counterion interactions that can (1) lead to direct
thermodynamic consequences through perturbation of the
ligand electron donation/withdrawal and (2) affect kinetic
barriers, including by sterically crowding the active site of the
reactive [CuOH]2+ core. While this work corroborates
previous findings that the O−H BDE values of the product
aqua complexes are key predictors of HAT reactivity (keeping
in mind the caveat that tunneling contributions are
implicated), it also shows that the influence of counterions
in these systems can be significant and that the effects on the
reactivity of the [CuOH]2+ core are more nuanced than might
be suggested simply by analysis of substituent electron
donating/withdrawing effects.
ASSOCIATED CONTENT
* Supporting Information
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/acs.inorg-
chem.8b01529.
■
S
Full descriptions of materials and methods of syntheses
and characterizations of studied compounds, determi-
nation of pKa values, full description of kinetics
measurements, preparation of aqueous EPR samples
(PDF)
Figure 12. Plot of log k vs O−H BDE for the series of copper(III)
hydroxide complexes studied to date.
by temperature dependencies of kinetic isotope effects and
compuations in previous work,3,4 as well as computational
analysis of the present work described separately.36
Accession Codes
CCDC 1845452−1845455 and 1845463−1845464 contain
the supplementary crystallographic data for this paper. These
data can be obtained free of charge via www.ccdc.cam.ac.uk/
data_request/cif, or by emailing data_request@ccdc.cam.ac.
uk, or by contacting The Cambridge Crystallographic Data
Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44
1223 336033.
CONCLUSIONS
■
In summary, two new copper(II) hydroxide ([CuOH]+)
+
complexes have been prepared, bearing charged para-NMe3
−
or -SO3 substituents on the flanking aryl rings of the
supporting bis(carboxamide)pyridine ligands. X-ray crystallo-
graphic and spectroscopic measurements of the [CuOH]+
complexes support similar geometries and electronic environ-
ments for their Cu(II) centers, as well as significant cation−
anion interactions in the secondary coordination sphere of the
SO3L4− supported compounds. The spectroscopic similarities,
however, contrast with the significant differences in electro-
chemical behavior, as indicated by distinct E1/2([CuOH]2+/
[CuOH]+) values for the NMe3L and SO3L4− supported
complexes. The former is oxidized at significantly higher
potentials when compared to the latter, by as much as 0.275 V.
The differential electronic effects of the cationic and anionic
substituents are also manifested in the UV−vis spectra of the
derived [CuOH]2+ species. Interestingly, the oxidation
AUTHOR INFORMATION
Corresponding Author
*E-mail: wbtolman@wustl.edu.
ORCID
Debanjan Dhar: 0000-0002-3027-0226
Gereon M. Yee: 0000-0001-7481-2018
William B. Tolman: 0000-0002-2243-6409
■
Author Contributions
†These authors contributed equally to the work.
Notes
The authors declare no competing financial interest.
−
potential of the −SO3 appended complexes is greatly
ACKNOWLEDGMENTS
enhanced (by ∼0.180 V) when measured in the presence of
excess Na(15-crown5)+, providing evidence for ion-pairing
interactions in solution. Perhaps most surprising is the fact
that, despite a nearly equal thermodynamic driving force for
the HAT reaction with DHA (i.e., almost identical BDE values
for the product O−H bond in the [Cu(OH2)]2+ complexes)
and greater oxidizing power, the [CuOH]2+ complex
supported by NMe3L reacts more than 150 times slower in
these reactions than the complex featuring the SO3L4− ligand.
We speculate that this difference is due to steric encumbrance
■
We thank the National Institutes of Health (R37GM47365 to
W.B.T.) and the Univ. of Minnesota (for the doctoral
dissertation fellowship to D.D.) for financial support of this
research, which was performed in the Dept. of Chemistry,
Univ. of Minnesota. The authors thank Dr. V. G. Young, Jr., for
assistance with X-ray crystallography. Some X-ray diffraction
experiments were performed using a crystal diffractometer
acquired through NSF-MRI Award No. CHE-1229400. We
thank B. Dereli, Dr. M. Momeni, and Prof. C. J. Cramer for
input and advice.
of substrate access by the bulky BArF − counterions in the
4
complex of NMe3L.36
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■
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