Countercations and Solvent Influence CO2 Reduction to Oxalate by Chalcogen-Bridged Tricopper Cyclophanates

Article abstract of DOI:10.1021/jacs.8b02508

One-electron reduction of Cu₃EL (L^3- = tris(β-diketiminate)cyclophane, and E = S, Se) affords [Cu₃EL]^-, which reacts with CO₂ to yield exclusively C₂O₄^2- (95% yield, TON = 24) and regenerate Cu₃EL. Stopped-flow UV/visible data support an A→B mechanism under pseudo-first-order conditions (k_{obs, 298K} = 115(2) s^-1), which is 10⁶ larger than those for reported copper complexes. The k_obs values are dependent on the countercation and solvent (e.g., k_obs is greater for [K(18-crown-6)]⁺ vs (Ph₃P)₂N⁺, and there is a 20-fold decrease in k_obs in THF vs DMF). Our results suggest a mechanism in which cations and solvent influence the stability of the transition state.

Full text of DOI:10.1021/jacs.8b02508

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Cite This: J. Am. Chem. Soc. XXXX, XXX, XXX−XXX
Countercations and Solvent Influence CO2 Reduction to Oxalate by
Chalcogen-Bridged Tricopper Cyclophanates
Brian J. Cook,†,‡ Gianna N. Di Francesco,† Khalil A. Abboud,† and Leslie J. Murray*,†,‡
†Center for Catalysis and ‡Florida Center for Heterocyclic Compounds, Department of Chemistry, University of Florida, Gainesville,
Florida 32611-7200, United States
S
* Supporting Information
ABSTRACT: One-electron reduction of Cu3EL (L3−
=
tris(β-diketiminate)cyclophane, and E = S, Se) affords
[Cu3EL]−, which reacts with CO2 to yield exclusively
2−
C2O4 (95% yield, TON = 24) and regenerate Cu3EL.
Stopped-flow UV/visible data support an A→B mecha-
nism under pseudo-first-order conditions (kobs,
=
298K
115(2) s−1), which is 106 larger than those for reported
copper complexes. The kobs values are dependent on the
countercation and solvent (e.g., kobs is greater for [K(18-
crown-6)]+ vs (Ph3P)2N+, and there is a 20-fold decrease
in kobs in THF vs DMF). Our results suggest a mechanism
in which cations and solvent influence the stability of the
transition state.
here is an urgent need to utilize or trap CO2 to mitigate
1
Tclimate change. Although biological systems effectively
fix CO2 in photosynthesis,2 similar chemical approaches remain
limited and typically rely on activated nucleophiles or
electrophiles.3 Conversion of CO2 into reduced C1 products
(e.g., CO or HCOO−) electrochemically or with dihydrogen
has been well explored.4 In contrast, homogeneous catalysts
that facilitate C−C bond formation from CO2 are scarce. CO2
reduction products with Cn, where n ≥ 2, are observed in
heterogeneous systems but are usually minor components.5
The simplest C−C coupled product from CO2 reduction is
oxalate, formally arising from one-electron reduction of two
CO2 molecules, followed by radical coupling.6 Few homoge-
neous systems catalyze CO2 reduction to oxalate, or CO2-
Figure 1. Reported catalysts for CO2 reduction to C2O42−. All except
Maverick’s are catalytic either chemically or under an applied potential.
sulfido)- and (μ3-selenido)tricopper complexes and the effect of
the cation and solvent on CO2RedOx rate.
RedOx (Figure 1).7 Jager’s electrocatalysts are proposed to
̈
liberate CO2•− from a transient NiI−CO2, with radical coupling
leading to oxalate.7b,8 Tanaka’s [M3S2]2+ catalysts (M = Co, Rh,
Ir) span a range of onset potentials with strong selectivity for
oxalate formation but are hampered by slow reaction kinetics.7d
In Bouwman’s electrocatalytic system, a dicopper(I) complex
reacts with CO2 in air to yield a bis[(μ-oxalato)dicopper(II)]
product and requires Li+ to induce oxalate dissociation under
catalytic turnover.7a Relatedly, Maverick’s dicopper complex is
capable of CO2RedOx, albeit under stoichiometric conditions
or as a batch-process catalyst.9 Oxalate dissociation kinetics
remain a challenge for CO2RedOx, and particularly for the
reported copper systems. The effect of cations and other
reaction parameters on CO2RedOx are underexplored for
homogeneous systems; recent reports have interrogated solvent
and cation effects on CO2 reduction to formate and CO,
respectively.10 Herein, we report CO2RedOx by anionic (μ3-
Cu3(μ3-S)L (1) and Cu3(μ3-Se)L (2) can be readily reduced
with 1.0 equiv of [K(18-crown-6)][C10H8] in THF to [K(18-
crown-6)][1] and [K(18-crown-6)][2], as expected from the
respective cyclic voltammograms (Figure S105).11−13 Notably,
[K(18-crown-6)][1] and [K(18-crown-6)][2] are unreactive
toward up to 5 equiv of H2O, but react upon exposure to O2.
Although we were unable to obtain suitable single crystals of
[K(18-crown-6)][Cu3EL], we could readily crystallize reduc-
tion products of 1 and 2 using KC10H8. [K(THF)3][1] and
[K(THF)3][2] are isostructural, with [K(THF)3]+ interacting
in an η5 fashion with one β-diketiminate arm of a pseudo-D3h
[Cu3EL]− ion (Figures 2 and S104). Structures of 1− and 2−
are also comparable to the parent neutral molecules with minor
Received: March 6, 2018
© XXXX American Chemical Society
A
DOI: 10.1021/jacs.8b02508
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Communication
(E1/2 = −1.94 V vs Fc/Fc+), FeCp(C6Me6) (E1/2 = −2.07 V),
or KC8 (E1/2 < −3.7 V)14 in the presence or absence of 2 mol%
Cu3EL. For the CoCp*2 reaction, [CoCp*2][Cu3EL] and
CoCp*2 are recovered; inclusion of stoichiometric KPF6
relative to CoCp*2 in the reaction results in consumption of
the reductant and formation of K2C2O4. Similarly, catalysis is
observed for FeCp(C6Me6) only in the presence of KPF6 and
2−
for KC8 (Figures S58−S60). The C2O4 yield exceeds 95%
(TON = 24) with respect to reductant using 2 as a catalyst, but
is lower for 1 (54%).25 Minimal oxalate is observed in the
absence of 1 or 2 (Figures S59 and S60), independent of
reductant. These results highlight the catalytic potential of 1
and 2 for CO2RedOx and reinforce the observed cation
influence.
To quantify the cation effect, we monitored single-turnover
reactions of 1− and 2− with CO2 by stopped-flow UV/visible
spectrophotometry (Figure 3). DMF was our initial solvent
choice because of concerns with oxalate precipitation (Figure
S55).24 Absorption data of the reaction of [K(18-crown-6)][1]
with CO2 (168 equiv) in DMF at −55 °C evidence formation
of 1; thus, an A→B mechanism was used to model all data sets,
as no intermediates were detected. The estimated pseudo-first-
Figure 2. ORTEP representation (50% probability) of [K(THF)3][2].
H-atoms and guest solvent are omitted for clarity. Cu, Se, C, N, and O
atoms are depicted as lime green, orange, gray, blue, and red ellipsoids,
respectively.
changes to the bond lengths.12,13 The structural changes
observed upon reduction agree with the Cu−E π* and Cu−N
σ* nature of the a2″ LUMO of Cu3EL reported elsewhere.12,13
Exposure of thawing THF slurries of solid [K(18-crown-
6)][Cu3EL] to ∼1 atm CO2 results in dissolution of the -ate
complexes to afford Cu3EL in solution and a gray precipitate.
This reaction takes ∼30 min at −80 °C, with near-quantitative
1
regeneration of Cu3EL (>99% by mass), as confirmed by H
NMR spectra of product mixtures (Figures S37−S40).
Resonances for free crown ether in these reactions are apparent
in the spectra, suggesting dissociation of K+. The following data
recorded on this precipitate support formation of an oxalate salt
from CO2: (i) IR absorption bands range from 1615 to 1680
cm−1 (Figure S62); (ii) a single resonance is seen at ∼160 ppm
in 13C NMR spectra (Figures S42 and S55); and (iii) the IR
absorptions shift and the 13C NMR resonance is enhanced in
products from reaction using 13CO2.7a,d,8,9 We conclude that
[K(18-crown-6)][Cu3EL] effects C−C coupling of CO2 to
yield K2C2O4 and regenerate Cu3EL. Importantly, the isolated
yield for K2C2O4 is 95%, exhibiting the high selectivity of the
Cu3EL complexes to effect CO2RedOx.
As part of our prior report of 1, we obtained preliminary data
indicating no reactivity of [CoCp*2][1] toward CO2, strongly
contrasting the observations for the [K(18-crown-6)]+ and
[K(THF)3]+ congeners.12 We, therefore, synthesized com-
pounds of the general formula Q[Cu3EL], where Q+ = Ph4P+,
PPN+ (PPN+ = bis(triphenylphosphine)iminium), Bu4N+,
[FeCp(C6Me6)]+, [K(2.2.2-cryptand)]+, [Na(15-crown-5)]+,
and [Na(THF)n]+, and evaluated reactions of these compounds
with CO2. We followed reaction progress qualitatively by the
diagnostic color change associated with regeneration of Cu3EL,
and analyzed product mixtures by 1H NMR spectroscopy
(Table S1). A cation dependence is evident from these data;
complexes with alkali countercations react at significantly lower
temperatures than the organic cation congeners. Encapsulation
of K+ within 2.2.2-cryptand agrees with this trend. Additionally,
2− is more reactive toward CO2 than 1−, consistent with 2−
having a ∼3.5 kcal/mol greater driving force given that E1/2(2)
− E1/2(1) = −0.15 V (Figure S105).12,13 As for reactions with
[K(18-crown-6)][Cu3EL], free crown ether or cryptand is
Figure 3. Stopped-flow UV/visible spectra in the 300−700 nm range
recorded at −55 °C in DMF for reaction of 8.4 mM CO2 with 0.05
mM [K(18-crown-6)][1] (top) and 0.12 mM [K(18-crown-6)][2]
(bottom). Isosbestic points are observed at 406 and 608 nm (top) and
at 408, 567, and 667 nm (bottom). Red trace represents the first time
point at 2.5 ms, and blue trace represents the spectrum at 1500 ms.
Relative to unreacted [K(18-crown-6)][2], A645 decreases by ∼0.40
within the instrument dead time.
1
observed in H NMR spectra of products from reaction of,
respectively, [Na(15-crown-5)]+ or [K(2.2.2-cryptand)]+ salts
with CO2 (Figures S46 and S47).
Given that 1 and 2 are regenerated, we postulated that these
complexes could function as homogeneous catalysts, and
evaluated the reaction products of CO2 with either CoCp*2
B
DOI: 10.1021/jacs.8b02508
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Journal of the American Chemical Society
Communication
Table 1. Values for kobs and Calculated Activation Parameters for CO2RedOx by 1−
kobs (s−1
)
c
countercation
[K(18-crown-6)]+
−55 °C
−15 °C
25 °C
ΔH⧧ (kcal/mol)
ΔS⧧ (cal/mol·K)
ΔG⧧ (kcal/mol)
298
a
11(2)
6.3(5)
0.39(4)
49(1)
38(5)
115(2)
141(9)
11(1)
2.9(1)
4.0(3)
−39(1)
−35(4)
−37(5)
−44(16)
14.6(4)
15(1)
16(2)
15(5)
a
[PPN]+
b
[PPN]+
0.93(3)
5.0(3)
a
d
d
[K(THF)3]+
14(1)
>231
>231
1.85(0.24)
a
b
c
d
Solvent = DMF. Solvent = THF. 25 °C values extrapolated from Eyring−Polanyi analysis. kobs estimated based on 3 ms instrument dead time
order rate constant for reaction of [K(18-crown-6)][2] with
CO2 (kobs > 231 s−1) is significantly greater than that for [K(18-
crown-6)][1] (kobs = 11(2) s−1), as 50+% of 2− reacts within
the 3 ms instrument dead time (Figure S65). All subsequent
kinetic analyses were performed with 1− salts, as the reaction
rate for 2− precludes an accurate determination of kobs. Kinetic
data collected at λmax of 1− (680 nm) and of 1 (805 nm) at
various temperatures were reliably fit to a single exponential,
affording statistically indistinguishable kobs values (Figure S66−
S102). Calculated activation parameters evidence a low ΔH⧧
Scheme 1. Proposed Mechanism of CO2 Activation and
Influence of Cation in Stabilizing Intermediates
and a large ΔS⧧ (Table 1) to afford ΔG⧧ values comparable
298
to experimentally and computationally determined values for
other CO2RedOx reactions.15 Surprisingly, kobs is either inverse
first or second order in [CO2] (Table S4 and Figure S103), and
implicates a more complex mechanism than generation and
homocoupling of freely diffusing CO2•−. Similar experiments
performed using THF as the solvent instead of DMF revealed a
significant solvent effect (Table 1 and Figures S92−S96). The
solvent trend is consistent with the observed cation effect,
providing two variables to tune the CO2RedOx rate in this
system.
Bouwman’s and Maverick’s compounds have slow rate
constants for oxalate formation (4.76 × 10−4 and 5.28 × 10−6
s−1, respectively), and both require weak reductants to effect
CO2RedOx.7a,9,15b In addition, product inhibition is likely a
limiting step under turnover conditions for these systems.
Similar product inhibition is inferred for the bis[β-diket-
iminatonickel(I)] complex from Limberg and co-workers.8
Contrastingly, 1 and 2 require more reducing potentials (−0.89
and −1.04 V vs SCE, respectively) than either the Bouwman
(−0.21 V) or Maverick (−0.32 V) reports, correlating with
larger estimated pseudo-first-order rate constants for CO2-
RedOx at 298 K. Attempts to coordinate ligands to 1 or 2 have
been unsuccessful, implying minimal product inhibition here.
Consistently, the highest catalytic rate constant for a Ni(I)
our activation parameters are comparable to those reported for
O2 activation in the presence of cations and for CO2
capture.18b,c,20 Therefore, an analogous transient in which the
cation accelerates electron transfer to CO2 is likely present here
as well. In a similar vein, solvent is reported to tune product
distribution for stoichiometric CO2 reduction by an iron(I)
complex;21 solvent modulates reaction kinetics without
affecting product distribution for [Cu3EL]−.
•−
The order in CO2 is inconsistent with CO2 dissociation,
differing from Jager’s examples and suggesting a more complex
̈
mechanism (Scheme 1, vertical arrow).7b Tanaka proposed that
one CO2 molecule first binds to a μ3-sulfide in the [M3S2]2+
catalysts, and a second CO2 then coordinates to a metal center,
followed by intramolecular C−C coupling,7d which resembles
the mechanisms proposed for CO2 reduction by Mo(O)SCu
carbon monoxide dehydrogenase and CO2 reduction by a
nucleophilic metal nitride.22 However, the chalcogens in 1 and
2 are sterically inaccessible to CO2, implicating an alternate
structure for the initial CO2 adduct here. In contrast to other
multimetallic CO2RedOx catalysts, 1− and 2− are one-electron
reductants, necessitating a bis(Cu3EL) transient or inter-
complex electron transfer (Scheme 1). DFT calculations on
the Bouwman system propose a CO2-bridged dicopper species
prior to C−C bond formation.15b A [(Cu3EL)2(CO2)]2−
transient would be comparable, and we note that approach of
two trimetallic cyclophanates is sterically feasible.23 High [CO2]
could sequester [Cu3EL]− as the CO2 adduct, effectively
lowering [Cu3EL]− and slowing oxalate formation. Such a
kinetic condition necessarily accumulates [Cu3EL(CO2)]−,
which is inconsistent with our data. Reactions at significantly
higher [CO2] may allow observation of this adduct, and these
experiments are part of ongoing investigations. In the final step,
complex reported by Jager is ∼1 × 105 M−1 s−1, and the onset
̈
potential for that catalyst is −2.24 V vs SCE.7b Values for kobs
are larger for 2− than the analogous 1− complex, further
reinforcing this observation. This conclusion resembles prior
reported correlations between E1/2 and reaction rate for O2
reduction electrocatalysts16 and in H2 electrocatalysis.17 One
aspect of catalyst design for CO2RedOx is the balance between
reducing power of the reactive species and product inhibition.6b
In our proposed mechanism (Scheme 1), [Cu3EL]− reacts
with CO2, affording [Cu3EL(CO2)]−, which is short-lived based
on our kinetic data. Electron transfer from [Cu3EL]− to CO2 is
likely rate controlling and not C−C bond formation, as kobs
values determined from A680 and A805 are within error. The
cation dependence observed here evokes reports on the
influence of redox-innocent cations and proximal hydrogen-
bonding interactions in facilitating O2 and CO2 bond
scission.18,19 Indeed, the calculated ΔS⧧ values are consistent
with an ordered transition state in the rate-controlling step, and
C
DOI: 10.1021/jacs.8b02508
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Communication
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ASSOCIATED CONTENT
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■
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AUTHOR INFORMATION
Corresponding Author
*murray@chem.ufl.edu
ORCID
■
Leslie J. Murray: 0000-0002-1568-958X
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
L.J.M. acknowledges a University of Florida departmental
instrumentation award (NSF CHE-1048604), ACS Petroleum
Research Fund (ACS-PRF 52704-DNI3), and NSF CHE-
1464876. G.N.D. acknowledges a University of Florida, College
of Liberal Arts and Sciences Graduate Research Fellowship.
K.A.A. acknowledges University of Florida and NSF CHE-
0821346 for funding an X-ray equipment purchase.
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E
DOI: 10.1021/jacs.8b02508
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