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BNL-112077-2016-JA
Identification of Ion-Pair Structures in Solution by Vibrational Stark Effects
J. Hack
Submitted to Journal of Physical Chemistry B
February 2016
Chemistry Department
Brookhaven National Laboratory
U.S. Department of Energy [DOE Office of Science]
Notice: This manuscript has been authored by employees of Brookhaven Science Associates, LLC under Contract No. DE- SC0012704 with the U.S. Department of Energy. The publisher by accepting the manuscript for publication acknowledges that the United States Government retains a non-exclusive, paid-up, irrevocable, world-wide license to publish or reproduce the published form of this manuscript, or allow others to do so, for United States Government purposes.
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1
Identification of Ion-Pair Structures in Solution by
Vibrational Stark Effects
John Hack, †‡ David C. Grills,† John R. Miller,† and Tomoyasu Mani†*
†Chemistry Department, Brookhaven National Laboratory, Upton NY 11973-5000, United States
‡Chemical Engineering Department, University of Virginia, 102 Engineers’ Way, PO Box
400741, Charlottesville, VA 22904-4741, United States
ABSTRACT:
Ion pairing is a fundamental consideration in many areas of chemistry and has implications in a
wide range of sciences and technologies that include batteries and organic photovoltaics. Ions in
solution are known to inhabit multiple possible states, including free ions (FI), contact ion pairs
(CIP), and solvent-separated ion pairs (SSIP). However, in solutions of organic radicals and non-
metal electrolytes, it is often difficult to distinguish between these states. In the first part of this
work, we report evidence for the formation of SSIPs in low-polarity solvents and distinct
measurements of CIP, SSIP, and FI, by using the ν(C≡N) infrared (IR) band of a nitrile-
substituted fluorene radical anion. Use of time-resolved IR detection following pulse radiolysis
allowed us to unambiguously assign the peak of the FI. In the presence of non-metal electrolytes,
two distinct red-shifted peaks were observed and assigned to the CIP and SSIP. The assignments
2
are interpreted in the framework of the vibrational Stark effect (VSE) and are supported by 1) the
solvent-dependence of ion-pair populations, 2) the observation of a cryptand-separated sodium
ion pair that mimics the formation of SSIPs, and 3) electronic structure calculations. In the
second part of this work, we show that a blue-shift of the ν(C≡N) IR band due to the VSE can be
induced in a nitrile-substituted fluorene anion by covalently tethering it to a metal-chelating
ligand that forms an intramolecular ion pair upon reduction and complexation with sodium ion.
This adds support to the conclusion that the shift in IR absorptions by ion pairing originates from
the VSE. These results combined show that we can identify ion-pair structures by using the VSE,
including the existence of SSIPs in a low-polarity solvent.
1. INTRODUCTION
Ion pairing between oppositely charged ions is fundamental to many areas of chemistry
and applications including batteries and organic photovoltaics.1-3 It is widely accepted that ion
pairs exist in solvents having dielectric constants De <30.3 Ion pairing affects both the energetics
and dynamics of ions. Ion pairings of electrolytes and those between organic anion radicals and
metal cations are well studied, but studies on ion pairing between organic radicals and non-metal
electrolytes are scarce despite their significant implications in the control of the energetics of
these molecules, including redox potentials measured by means of electrochemistry. Ion pairing
is particularly important in solvents of low dielectric constant (De <11, here we use the definition
by Van Der Hoeven and Lyklema4 based on Fuoss’s theory5) where Coulomb interactions are
less screened by the solvent.
The exact form of ion pairs in solution is a topic of debate,3 and is often difficult to
identify experimentally. In general, we classify ion pairs into three populations: contact ion pairs
3
(CIPs) or tight ion pairs, solvent-separated ion pairs (SSIPs) or loose ion pairs, and solvent-
shared ion pairs (SIPs) (Figure 1). We can consider SIP as one case of SSIP with one solvent
molecule between the two ions (therefore shared). For simplicity and due to the currently
achievable experimental resolution, we do not attempt to make any distinction between SIPs and
SSIPs, and call them SSIPs collectively, assuming that there is a single solvent molecule between
ions.
The formation of a SSIP can have significant implications for the kinetics of charge
separation and thermodynamic properties, including electrical conductivity, of a solution6-7 and
photo-induced charge transfer processes in condensed phases.8-9 The concept of a SSIP was first
introduced to explain unusual conductivity and kinetic data by Fuoss, Grunwald, and Winstein in
three independent studies.6-7, 10 The proposed SSIP species rests in a relative potential energy
minimum, separated from the CIP by a solvation energy barrier.
Figure 1. Pictorial representation of free ions (FI), a solvent-separated ion pair (SSIP) or solvent-
shared ion pair (SIP), and a contact ion pair (CIP). Void circle = solvent molecule.
Various studies have been performed to identify and characterize SSIPs, but many are limited to
pairing with metal cations.3 Identifying and distinguishing SSIPs from CIPs and free ions (FIs) in
solution, especially of organic radicals and non-metal electrolytes, is not a trivial task. Until
recently, spectroscopic techniques such as electron paramagnetic resonance, visible, and
vibrational spectroscopy have been able to observe distinct peaks assigned to CIPs and SSIPs,
- -
+
-
+
FI CIPSSIP (or SIP)
+
4
but unambiguous measurements of FIs in neat solvents, especially in media of low polarity, were
difficult.2, 11-13 This technical limitation left some ambiguity as to the assignment of SSIPs as
distinct thermodynamic states with unique properties and behaviors from FIs. In previous work,
we reported a distinct IR absorption band of the free anion of F1CN, a nitrile-substituted fluorene
(Chart 1), in moderately non-polar neat tetrahydrofuran (THF, De = 7.52),14 using nanosecond
time-resolved infrared (TRIR) detection15 following pulse radiolysis at the Laser-Electron
Accelerator Facility (LEAF) at Brookhaven National Laboratory16. The ν(C≡N) IR peak for the
free ion of F1CN-• (νmax = 2110 cm-1) was red-shifted in comparison to the peak in the neutral
form (νmax = 2224 cm-1). Two further red-shifted peaks were observed upon ion pairing with
tetrabutylammonium ion (TBA+), which we assigned to the CIP (νmax = 2094 cm-1) and the SSIP
(νmax = 2108 cm-1) (Figure 2). We also determined the rate constants for the formation of ion
pairs by monitoring the growth of the ion pair’s IR band.14
Chart 1. Structures of the Nitrile-Substituted Fluorenes Studied.
N
O
OO
OO
N
F1CN
15-crown-5-Ph-F1CN
5
Figure 2. IR spectra of the ν(C≡N) stretch of F1CN anion in THF. The FI (νmax = 2110 cm-1) was
obtained 50 ns after pulse radiolysis. A further red-shift occurs when F1CN is reduced
electrochemically in the presence of 0.1 M TBA+PF6- or TBA+BF4
-. The ion pair spectrum fits to
two Voigt functions, assigned to the CIP (νmax = 2094 cm-1) and SSIP (νmax = 2108 cm-1). The
data shown in Figure 2 were taken from ref.14
The observed red-shifts in ν(C≡N) IR frequency of the F1CN anion with ion pairing can
be interpreted in the framework of the vibrational Stark effect (VSE).14 This interpretation for
ion pairing between F1CN-• and positive charge was supported by electronic structure
calculations.14 The VSE relates shifts in the vibrational energy to the dot product of the
difference dipole of the vibration (between ground and excited states) and an exerted external
electric field. Vibrational shifts of nitrile and carbonyl probes in neutral molecules have been
calibrated to a known external field in frozen glass and in solution.17-18 These neutral probes can
be used to obtain a quantitative measure of the local electric field in proteins, other
macromolecular complexes, or environments, pioneered by Boxer and colleagues.18-19 Neutral
molecules (benzonitrile)20 and charged species (e.g. thiocyanate)21-22 can be used to investigate
the properties of ionic liquids. In the case of ion pairing between F1CN-• and TBA+, the inert
non-metal ion TBA+ induces an external electric field on the C≡N dipole (the difference dipole
1.0
0.8
0.6
0.4
0.2
0.0N
orm
aliz
ed A
bsor
ban
ce2160214021202100208020602040
Wavenumber (cm-1
)
FIIon Pairs
CIP
SSIP
6
moment points N→C),23 causing a change in the energy of its vibrational transition (Supporting
Information Figure S1). The observed red-shift in the ion pairs of F1CN-• and TBA+ implies that
a cation is oriented on the nitrile side of F1CN-• in the ion pairs. We assumed that the ion pair
spectrum, shown in Figure 2, includes both CIP (F1CN-•, TBA+) and SSIP (F1CN-•||TBA+) based
on the goodness of fits, but more support is required for a clearer confirmation of the existence of
a SSIP as one of the distinct ion-pair structures and as a thermodynamically stable species.
Within the framework of the VSE, we also hypothesize that we can induce a blue-shift in ν(C≡N)
IR frequency of the nitrile-substituted fluorenes when a cation is positioned on the other side of
the nitrile of anions, away from the CN. Such a blue-shift was previously observed in the cases
of simple salts such as Na+SCN-,24 Li+SCN-,25 Li+SeCN-,26 and inorganic cyanide compounds,27
but the relationships between the VSE and ion-pair structures are often obscured because of the
lack of a priori structural information in such intermolecular complexes in solution.
In the first part of the current paper, we present evidence that supports the existence of
the SSIP in the ion pairs of F1CN anion and TBA+ by 1) the solvent-dependence of the SSIP/CIP
ratio in dimethylformamide (DMF)/THF mixtures, 2) the agreement of peak position in
cryptand-separated sodium ion pair that mimics the formation of a SSIP, and 3) electronic
structure calculations. In the second part of the paper, we report a blue shifted ν(C≡N) IR
vibration due to the VSE in the nitrile-functionalized fluorene radical anion by using a newly
synthesized fluorene derivative (Chart 1) whose intramolecular ion-pair structure is known a
priori. These works combined highlight the strength of the VSE in the studies of inter- and
intramolecular ion pairs. We will discuss ion-pair structures in media of low polarity and their
implications in the energetics of ions in such environments.
7
2. EXPERIMENTAL METHODS
2.1. General. All reagents and solvents were used as received from standard sources. Silica gel
(pore size 60 Å, 70-230 mesh, Sigma-Aldrich) was used for column chromatography. FTIR
spectra were obtained using a Thermo Nicolet Nexus 670 FTIR spectrometer. UV-Vis-NIR
spectra were obtained using a Cary 5 spectrophotometer (Varian) or Cary 5000
spectrophotometer (Agilent). 1H and 13C NMR spectra were obtained with a Bruker Avance III
spectrometer operating at 400.16 and 100.62 MHz, respectively. Samples were prepared under
an atmosphere of argon for measurements other than NMR spectroscopy.
2.2. Synthesis. Syntheses and characterizations of F1CN,14 BrF1CN,28 and 4-Bromobenzo-15-
crown-529-30 are reported elsewhere.
2.2.1. 9,9-dihexyl-7-(4,4,5,5-tetramethyl-1,3,2-dioxaborolan-2-yl)-9H-fluorene-2-carbonitrile
(BpinF1CN) A dry flask with an argon atmosphere was charged with BrF1CN (0.61 g, 1.39
mmol), Bis(pinacolato)diboron (0.45 g, 1.74 mmol), Pd(dppf)Cl2 (6.0 mol%), and KOAc (0.273
g, 2.78 mmol). After addition of 5 mL of DMF, the reaction mixture was stirred at 80 oC for 5 h,
after which it was cooled down to room temperature and poured into 50 mL of H2O. The mixture
was extracted with DCM (50 mL). The organic layer was collected and washed with brine three
times, and then dried over MgSO4. After filtration and removal of the solvents in vacuo, the
crude mixture was purified by column chromatography on silica gel (hexane:DCM = 7:3) to
afford the title compound as a pale yellow solid (0.40 g, 59%): 1H NMR (CD2Cl2, 400 MHz)
7.81 (m, 4H), 7.65 (m, 2H), 2.03 (m, 4H), 1.36 (s, 12H), 1.10-0.97 (m, 12H), 0.76 (t, J = 7.1 Hz,
6H), 0.51 (m, 4H); 13C NMR (CD2Cl2, 100 MHz) 52.4, 151.0, 145.9, 142.4, 134.2, 131.6,
127.2, 121.0, 120.4, 120.1, 110.8, 84.4, 56.0, 40.4, 31.9, 29.9, 25.1, 24.1, 22.9, 14.1.
8
2.2.2. 9,9-dihexyl-7-(2,3,5,6,8,9,11,12-octahydrobenzo[b][1,4,7,10,13]pentaoxacyclopentadecin-
15-yl)-9H-fluorene-2-carbonitrile (15-crown-5-Ph-F1CN). The compound BpinF1CN (0.200 g,
412 μmol), 4-bromobenzo-15-crown-5 (0.158 g, 453 μmol), Pd(OAc)2 (11 mg, 49 μmol), SPhos
(41 mg, 99 μmol), and K3PO4 (0.179 g, 824 μmol) were combined in a 10 mL round-bottom
flask. Toluene (2 mL) and H2O (0.1 mL) were added and the resulting mixture was
deoxygenated by a couple of evacuation-refill (Ar) cycles. The mixture was stirred at 80 oC for 2
h and then cooled down and the solvents were removed in vacuo. The crude mixture was purified
by column chromatography on silica gel (5% THF in DCM) to afford the title compound as a
yellow solid (0.183 g, 71%): 1H NMR (CD2Cl2, 400 MHz) 7.81-7.79 (m, 2H), 7.66-7.64 (m,
2H), 7.60 (dd, J = 7.9, 1.5 Hz, 1H), 7.57 (s, 1 7.23 (dd, J = , 8.3, 2.1 Hz, 1H), 7.19 (d, J = 2.1
Hz, 1H), 6.98 (d, J = 8.2 Hz, 1H) 4.22-4.15 (m, 4H), 3.91-3.87 (m, 4H), 3.74-3.67 (m, 10 H),
2.06-2.01 (m, 4H), 1.83-1.80 (m, 2H), 1.13-1.01 (m, 12H), 0.76 (t, J = 7.1 Hz, 6H), 0.61 (m, 4H);
13C NMR (CD2Cl2, 100 MHz)
, 120.3, 114.3, 113.6, 110.1, 71.19. 71.16, 70.7, 60.6, 69.9, 69.8, 69.6, 69.2, 40.5, 31.9,
29.9, 24.1, 22.9, 14.1.
2.3. Spectroelectrochemistry. FTIR spectroelectrochemical measurements were performed in a
spectroelectrochemical transmission cell GS20900 (Specac Limited), coupled with a CHI 600E
Potentiostat (CH Instruments, Inc.). All samples contained 0.1 M TBA+PF6-.
2.4. FTIR of Reduced Species. Chemical titrations were performed with sodium biphenyl
(Na+BP-) or sodium chunks, either in the absence or presence of cryptand-222 (C222), in a
glovebox. The concentration of C222 was at least one equivalent of neutral molecules before
9
reduction unless otherwise noted. The solutions were transferred into a demountable liquid flow
cell DLC-S25 (Harrick Scientific) equipped with 2 mm thick CaF2 windows and a 950 μm thick
spacer.
2.5. Pulse Radiolysis with Time-resolved Infrared (TRIR) Detection. Pulse radiolysis
experiments were conducted in LEAF at Brookhaven National Laboratory. A detailed
description of the experimental setup is given elsewhere.15 An airtight IR solution flow cell was
used (1.1 mm pathlength), equipped with 0.35 mm thick CaF2 windows. Continuous wave
external-cavity quantum cascade lasers (models 21047-MHF and 21043-MHF, Daylight
Solutions, Inc.) were used as the IR probe source. Samples (5-20 mM) were dissolved in THF or
DMF and purged with argon.
2.6. Quantum Calculations. Computations were performed with either Gaussian0931 or Q-
Chem 4.232 using density functional theory (DFT). All calculations on anions were spin-
unrestricted. All hexyl groups at the 9-position of fluorenes were replaced by ethyl groups. The
geometry optimizations were performed without symmetry constraints. Calculations used the 6-
31G(d) basis set and the LC-ωPBE (ω = 0.1 bohr-1)28, 30, 33 or B3LYP34-35 functional. A PCM
solvation model36-37 (Gaussian) or C-PCM solvation model38-40 (Q-Chem) for THF was used.
3. RESULTS AND DISCUSSION
3.1. SSIP vs. CIP. 3.1.1. Ion Pair Populations in DMF/THF Mixtures In the first part of this
work, we examine the existence of a SSIP in the ion pairs of F1CN-• and TBA+ in solvents of low
polarity. In THF solution containing 0.1 M TBA+PF6-, which is a typical total concentration of
electrolytes in electrochemical measurements, the results indicated that >95% of the ions exist as
ion pairs based on fits such as those shown in Figure S2. We determined that the spectrum of the
10
SSIP is distinct from that of the FI. To further support the assignment of the SSIP peak (Figure
2), IR spectra in the region of the ν(C≡N) stretch were measured spectroelectrochemically in
mixtures of DMF (De = 38.25) and THF containing 0.1 M TBA+PF6- (Figure 3). The solvent
properties of the mixtures are reported in Supporting Information Table S1. The spectra obtained
in mixtures up to 20% DMF by volume fit to two Voigt functions centered at 2094 cm-1,
assigned to the CIP, and 2108 cm-1, assigned to the SSIP, with a SSIP/CIP ratio that increases
with DMF composition. The characteristics of their spectra are listed in Table 1. The two Voigt
functions were constrained to the same position, full width at half maximum (FWHM), and
shape as the functions used to fit the spectrum of ion pairs in pure THF solvent. Broader
linewidths observed for the CIP and the SSIP bands in the presence of TBA+ compared to that
for the FI indicate that multiple conformations exist in both the CIP and the SSIP. However, as
mentioned above, in this work we do not attempt to distinguish these conformers. The fittings of
the spectra at each concentration are reported in Supporting Information Figure S2.
Figure 3. Spectroelectrochemical measurements of the ν(C≡N) stretch of F1CN-• in mixtures of
DMF and THF containing 0.1 M TBA+PF6-. The spectra obtained in mixtures up to 20% DMF
CIP SSIP
Abs
orb
an
ce (
a.u
.)
21402120210020802060
Wavenumbers (cm-1
)
100%
70%
50%30%20%
15%10%
5%
0%
% DMF
11
by volume fit well to two Voigt functions centered at 2094 cm-1, assigned to the CIP, and 2108
cm-1, assigned to the SSIP, with a SSIP/CIP ratio that increases with DMF composition.
Beyond 20% DMF, the peaks shift noticeably and no longer fit cleanly to the same two
Voigt functions. As DMF becomes a larger fraction of the bulk solvent, several confounding
effects may occur. As the dielectric constant of the bulk solvent increases the ν(C≡N) peak
maximum can shift to the red probably due to solvent-induced VSE,17-18 as seen for F1CN-• free
ions, νmax = 2101 cm-1 in DMF, vs. νmax = 2110 cm-1 in THF (the 9 cm-1 red-shift can be seen in
Supporting Information Figure S3). As the dielectric constant of the solvent increases, the
Coulomb attractions between oppositely charged species are more screened. The free ion
becomes more stabilized, and may begin to contribute to the spectra. The observed change in FI
from THF to DMF may be explained by a simple Onsager-like model of solvation as in the case
of neutral aromatic nitriles, but more local solute-solvent interactions41 are likely to play an
important role in anions. Interestingly, the spectrum reported in 100% DMF solvent containing
0.1 M electrolyte is centered, except for a slight electrolyte-induced red-shift, at almost the same
position, 2101 cm-1, as the free ion peak in neat DMF, observed by TRIR detection following
pulse radiolysis (Supporting Information Figure S4a). We can attribute this red-shift to the
presence of ion pairs. We can fit the spectra to two functions by fixing one function to the
spectrum of the free ion in DMF while adjusting the other, assumed to arise from ion pairs
(Supporting Information Figure S4b). Here, we did not make any distinctions among ion pairs,
and we find adequate, though not excellent fits. Assuming that the ratio of the areas under the
curves reflects the population ratio, we can estimate the association constant for ion pairing
between F1CN-• and TBA+ in DMF to be KA ~ 3 M-1, which is about four orders of magnitude
smaller than that in THF (KA = 9.2 × 104 M-1).14 The value KA ~ 3 M-1 is one order of magnitude
12
smaller than that for TBA+PF6- in acetonitrile reported by Geiger and coworkers (KA = 35.7 M-
1),42-43, which seems reasonable given that KA(F1CN-•,TBA+) is about four times smaller than
KA(TBA+,PF6
-) in THF (see Supporting Information Table S2). An alternative is to interpret a
slight broadening in the redder side in the spectroelectrochemical peak in terms of changes in the
dielectric constants due to the presence of 0.1 M electrolyte. The first interpretation which
identifies a reasonable KA(F1CN-•,TBA+) seems preferable. In both cases, the observation
supports the picture that a radical anion is mainly “free” in a polar solvent in the presence of
electrolytes, while some may be stabilized by ion pairings. This is different from the picture we
previously put forward that ions are mainly stabilized by a single ion pair in non-polar
environments.14
Table 1. Observed ν(C≡N) (cm-1) of Anion Species in THFa
Molecules FIb
Ion Pairs
TBA+ Na+
CIPc SSIPc Na+d Na(C222)+d
F1CN 2110
(12.1)
2094
(21.5)
2108
(18.8)
2094
(22.8)
2108e
(15.9)
2106f
(20.0)
15-crown-5-Ph-F1CN 2124
(18.4)
2106
(23.6)
2122
(26.3)
2136
(23.1)
2122
(24.0)
aThe numbers in parentheses are the fwhm values determined by fitting with a Voigt function. Uncertanities are ± 1 cm-1. bObtained by pulse radiolysis. cObtained by spectroelectrochemistry. dObtained by chemical reductions. eFitted to two Voigt functions. The reported value in the Table is assigned to the species that “mimics” the SSIP. The other peak was observed at 2097 cm-1. fFitted to one Voigt function.
13
In order to calculate association constants for the CIP and SSIPs, the SSIP/CIP ratios
obtained in mixtures up to 20% DMF were fit to the model represented in Scheme 1. This model
allows for the formation of a CIP, a SSIP separated by one THF molecule, and a SSIP separated
by one DMF molecule. We abbreviate these as SSIPTHF and SSIPDMF. The SSIP peak in Figure 3
is taken to be a combination of contributions from both SSIP species. The model makes the
following assumptions. First, we assume that there is only one solvent molecule separating ions
in the SSIP species, though it is possible that there are multiple.3 Second, we assume that in
mixtures up to 20% DMF we can neglect formation of free ions and assume that DMF does not
preferentially solvate either type of ion pair. Third, we assume that the ratio of the extinction
coefficients of ν(C≡N) absorption in CIP and SSIP is unity. The second assumption is supported
by the absence of a clear peak corresponding to that of the FI in DMF. The last assumption is
supported by DFT calculations, which show that the extinction coefficients of ν(C≡N) absorption
of anions differ by only 1-1.5 times over a distance of 4-5 Å between F1CN-• and a positive
charge.14 The maximum difference in the distance between SSIP and CIP was reported to likely
be ~4-5 Å.8-9 This difference is partially confirmed by computations (Supporting Information
Figure S5). Importantly, using a ratio of 1.5 does not yield significantly different results from
those reported (Supporting Information Table S3). A similar approach estimating extinction
coefficients of IR absorbance using the outputs from electronic structure calculations was
recently adopted by Ludwig and coworkers for far-IR absorption bands.12-13
Scheme 1. Ion Pairing Equilibria in DMF/THF Mixtures
TBA+ + PF6- ⇌ (TBA+, PF6
-) (1)
F1CN-• + TBA+ ⇌ CIP (F1CN-•, TBA+) (2)
14
F1CN-• + THF + TBA+ ⇌ SSIPTHF (F1CN-•, THF, TBA+) (3)
F1CN-• + DMF + TBA+ ⇌ SSIPDMF (F1CN-•, DMF, TBA+) (4)
Calculated association constants for the formation of CIP and SSIPs are reported in Table
2. The association constants KA for TBA+PF6- were taken to be mole-fraction weighted linear
combinations of association constants determined by Geiger and coworkers42 in THF (3.73 x 105
M-1), and DMF (35.7 M-1),43 and not adjusted. The association constant for the formation of
SSIPDMF, K4, is about an order of magnitude larger than for SSIPTHF, K3. This result matches the
expectation that DMF, a more polar solvent than THF, has a higher association constant for SSIP
formation, and explains the observed growth in the combined SSIP peak with increasing DMF
composition. Standard free energy changes (ΔG0) are reported in Table 2 as well. K3 was
multiplied by the concentration of THF so that the energy of the SSIPTHF could be compared to
the energy of the CIP. It was found that the CIP (ΔG0 = -300 ± 5 meV) is stabilized by 30 meV
compared to the SSIPTHF (ΔG0 = -270 ± 10 meV). This value is not much larger than the average
thermal energy, kBT at room temperature.
Table 2. Ion Pairing Equilibria for F1CN-• with TBA+
Equationa KAb ΔG0 (meV)b
2 K2 = (9.8 ± 1.8) × 104 M-1 -300 ± 5
3 K3 = (2.9 ± 1.1) × 103 M-2 -200 ± 10
3 K3 [THF] = (3.4 ± 1.3) × 104 M-1 -270 ± 10c
4 K4 = (2.6 ± 1.1) × 104 M-2 -260 ± 10
a Equations as defined in Scheme 1. b The values are the averages from five different DMF/THF ratios. c100% THF ([THF] = 12.3 M).
15
3.1.2. Macrocycle Polyether-Separated Intermolecular Ion Pair.
The FTIR spectrum of F1CN-• in THF obtained by chemical reduction of F1CN by
sodium biphenyl is shown in Figure 4. The ν(C≡N) absorption band of the CIP (F1CN-•, Na+)
appears at 2094 cm-1, within one wavenumber of that of the CIP with TBA+ in THF (Figure 2
and Supporting Information Figure S6). Here we assign the peak to CIP based on the previous
studies by Szwarc and coworkers showing that Na+ preferentially forms a CIP with anions in
THF except for bulky anions (e.g. tetraphenylboride).1, 44 This agreement, despite the large
difference in ionic radii of Na+ and TBA+ (crystal ionic radii for Na+ and TBA+ are 0.95 and 4.13
Å, respectively),1, 45-46 may support the theoretical model presented by Fry47-48 that the alkyl
chains in TBA+ contribute negligible steric hindrance to ion pairing with anions. This model is
similar in nature to the concept of penetrated ion pairs.49-51
Figure 4. FTIR spectra obtained for F1CN anions in THF. The red spectrum, peaking at 2094
cm-1, labeled F1CN-•, Na+, was produced by a chemical reduction with sodium biphenyl. The
green spectrum, labeled F1CN-•||Na(C222)+,was produced by reduction with excess sodium
biphenyl in the presence of cryptand-222 (C222). A main peak is observed at 2108 cm-1 with an
16
additional shoulder at 2097 cm-1. C222 encapsulates the Na+ ion, creating a “separated” ion pair,
which mimics the SSIP.52-54 The FI spectrum, labeled F1CN-•, (νmax = 2110 cm-1) was recorded
with TRIR detection obtained 50 ns after pulse radiolysis in neat THF. The balls in the top of the
figure represent Na+.
In the presence of cryptand-222 (C222), a main peak is observed at 2108 cm-1 together
with a shoulder at 2097 cm-1, the former of which is within one wavenumber of the SSIP
observed with TBA+. C222, a chelating agent, encapsulates the Na+ ion, creating a “separated”
ion pair with F1CN-• as F1CN-•||Na(C222)+, which can mimic the SSIP (F1CN-•||TBA+). This
strategy prevents close contact between alkali metals and anions, as shown elegantly by Kochi
and coworkers.52-54 The notable agreement between the observed peak and the peak at 2108 cm-1
in the presence of TBA+ further supports the assignment of the SSIPTHF (F1CN-•, THF, TBA+)
observed in Figures 2 and 4. The shoulder observed at 2097 cm-1 (FWHM = 15.8 cm-1 from
multi-peak curve fitting) is relatively close to that of the CIP observed with TBA+ and Na+ (νmax
= 2094 cm-1). This may indicate either an incomplete encapsulation of Na+ by C222 or the
possibility that some Na(C222)+ can be positioned close enough to resemble the CIP. Changing
the concentration of C222 (from one equivalent of neutral F1CN to three equivalents) appears
not to change the ratio of the two peaks, which suggests the latter is the case. It is important to
note that the observed spectrum of “separated” ion pair with Na(C222)+ is clearly different from
that of the FI, showing that while the cryptands can reduce the effects of ion pairing they will not
create FIs in media of low polarity.
Combining the solvent dependence and chemical reduction data clearly shows that we
can distinguish the SSIP from the FI and in turn supports our initial assumption that anions in
THF containing 0.1 M TBA+PF6- primarily exist as ion pairs. While it is expected from the
17
thermodynamic equilibria (KA = 9.2 × 104 M-1 for ion pairs as a whole14) that 2% of F1CN-• will
be present as FIs in such a solution, our IR measurements cannot readily detect this small
population. Analysis indicates that it would likely be possible to detect if at least 5% FIs are
present.
3.1.3. Energetics of Ion Pairs. We now turn our attention to the effect of ion parings on
the energetics of the ions, obtaining estimates of the standard reduction potential (E0) in the
absence of electrolytes and in the presence of electrolytes (typically 0.1 M) determined by cyclic
voltammetry. The standard energetics of the SSIPTHF and CIP in THF (vs. Fc+/0) are shown in
Figure 5, including the previously determined energies of the FI and Ion Pairs.14 We further
compare the experimental values with estimates by electronic structure calculations. The typical
calculated structures of the SSIPTHF and the CIP are displayed in Supporting Information Figure
S5.
Figure 5. Standard energy diagram in THF (E0 vs. Fc+/0 in eV) of the Ion Pairs, CIP, SSIP, and
FI of F1CN-• in THF. CIP = (F1CN-•, TBA+). SSIP = (F1CN-•, THF, TBA+). FI = F1CN-•. The
energy of the Ion Pairs is taken from ref28 and the energy of the FI was determined using
FI-3.01
CIP-2.71
∆G0IP = -290 meV
SSIP-2.74
∆ ∆ G0
= -30 meV
Ion Pairs-2.72
18
equation 5 with the value of KA from ref14. Note in ref14 the reduction potential was measured
with 0.1 M TBA+BF4-, which is -2.74 V vs. Fc+/0. Although the difference is only 20 meV, for
clarity, we used the reduction potential determined under the same conditions as those used in
the current study (0.1 M TBA+PF6-). The energy of the FI shifts by 20 meV accordingly while KA
does not change. The energy spacing in the graph is arbitrary.
The SSIPTHF was calculated as a F1CN-•, TBA+ pair surrounded by four explicit THF molecules.
We found that the TBA+ ion is stabilized by approximately four THF molecules (Supporting
Information Figure S7) in a similar manner to DMF and DMSO.55 The CIP was calculated by
removing one THF molecule and allowing F1CN-• to move closer to the cation. The distances
between the nitrogen atoms in F1CN-• and TBA+ (dNN) are ~3.8 Å for the CIP and ~7.8 Å for the
SSIP. We did not fully optimize the structures, but selected the structures, obtained in the course
of the optimization processes, which are consistent with our IR observations and interpretation
by VSE; namely the configuration in which TBA+ is positioned on the same side as the nitrile of
F1CN-•. (See Supporting Information pp. 2-3 for the details). Our current result of intramolecular
ion pairs further supports that the IR shift is induced by the VSE (vide infra). The CIP was
calculated to be more stable than the SSIPTHF by a free energy difference of approximately 40
meV. This value is in a good agreement with the experimentally determined value (~30 meV).
However, we would like to note that the structures used here are among the many stable
conformations and more samplings of conformations of ion pairs would be required to further
validate these preliminary computational results.
In a solution of single solvent where the concentration of solvent is constant, the
SSIP/CIP ratio should not change (eqs 2 and 3). We previously demonstrated that the reduction
potentials in the absence of electrolytes can be determined by measuring the association
19
constants for ion pairing.14 In the previous analysis,14 we did not distinguish SSIP/CIP
contributions, instead treating ion pairing as a mixture of the two. Here, we could determine the
respective standard energies of the CIP and SSIP within a couple of assumptions. The estimated
energy difference ~kBT makes it difficult to measure the contribution of individual species by
cyclic voltammetry measurements whose experimental uncertainties are typically about 10-20
mV.56 While it would be desirable to measure the difference by slowly scanning the potential, we
do not expect to alter the ratio of CIP and SSIP peaks by adjusting potentials if the two forms are
in equilibrium with each other for each ion pair. While the SSIP/CIP ratios are likely to depend
on the molecules of interest and electrolytes of choice, based on our current study, we can safely
say that both species are present in ion pairs between other molecules and electrolytes in a
similar environment. Therefore, the standard reduction potentials measured in the presence of
counter-ions M+ from electrolytes by electrochemical methods (E0([M+])) should be considered
to reflect the contributions from both species. In general, E0([M+]) is related to the standard
reduction potential in the absence of electrolytes (E([M+] = 0)) and to the association constants
(KA) of ion pairing between redox-active species and counter-ions from electrolytes through the
following equation:57-58
M M 0 ln 1 M (5)
where R is the gas constant, T is the temperature, F is the Faraday constant, and [M+] is the
concentration of free counter-ions M+ from the electrolyte. The energy diagram in Figure 5 and
our previous report14 used a standard condition for all solutes including the electrolytes, giving
E0([M+] = [M+]0 = 1M). We now consider the case where CIP and SSIP are present. Assuming
that only free ions contribute to the electron transfer reactions at the electrode, we obtain the
following equation:
20
M M 0 ln 1 M Solv M (6)
where KCIP and KSSIP are defined in Supporting Information pp. 4-5. This shows that the standard
energy of ion pairs E0([M+]), obtained by cyclic voltammetry, is affected by the additional
contribution of stabilization by SSIPs. In the current case, the standard energies of the CIP and
ion pairs obtained by cyclic voltammetry are within the experimental errors (± 10 meV) and we
are unfortunately not able to clearly observe the contributions from the SSIP in the standard
redox potentials measured electrochemically.
3.1.4. Differences in Standard and Experimental Conditions. As mentioned above, in our
previous report,14 we defined the reduction potential in the presence of electrolytes at the
standard condition, meaning that we have E0([M+]0), where usually [M+]0 = 1M and use E1/2
from the measurement of cyclic voltammetry as E0([M+]0) referenced to the ferrocene couple (vs.
Fc+/0).59 This convention allows us to use equation 5 in the simplest form, making it easy to
estimate and compare the redox potentials in the absence of electrolytes within the assumption
that the major stabilization of redox-active ions by electrolytes comes from single ion pairs.
However, the experimental concentration of free counter-ions from 0.1 M of typical electrolytes
is much lower in media of low polarity (e.g. [M+] = [TBA+] ~ 500 μM for TBA+PF6- in THF). In
this regard, it may be more appropriate to define the E1/2 measured by cyclic voltammetry as the
reduction potential in the presence of electrolytes at the specific experimental condition. Here the
concentration of free counter-ions in equation 5 is no longer [M+]0 and becomes the value at the
experimental condition in specific media with specific electrolytes. Based on this definition, we
obtained the contribution of ion pairs as ~100 meV for the condition of 0.1 M TBA+PF6- in THF.
The reduction potential of F1CN in the absence of electrolytes is then calculated to be -2.82 V vs.
Fc+/0 where the oxidation potential of ferrocene is defined under the same conditions. This value
21
is ~ 190 mV more positive than the energy estimated under the standard conditions (Figure 5).
We would like to emphasize that this difference is solely from the differences in the definition,
either at standard or specific experimental conditions for electrolytes, and it still holds that the
measurements of association constants between redox-active species and counter-ions from
electrolytes allow us to estimate the redox potentials in the absence of electrolytes, assuming that
only ion pairing shifts the potentials.
3.2. Blue-shift VSE. 3.2.1. Molecular Design. In the second part of the study, we examine
the possibility of inducing a blue-shift by VSE. Recall that a positively charged TBA+ sits on the
same side of the molecule as the nitrile group in F1CN-•, inducing a red-shift in ν(C≡N). In the
framework of the VSE a positive charge positioned on the opposite side, away from the nitrile
group, might produce a blue-shift in ν(C≡N). Such a blue-shift was previously observed in
simple inorganic slats.25-27 The blue-shift was interpreted based in the framework of the VSE:
The difference dipole moment of ν(C≡N) is aligned antiparallel to the external electric field
induced by the metal cation and cyanide-containing anions. However, the relationships between
the VSE and ion-pair structures are often obscured because of the lack of a priori structural
information in these complexes in solution.
Here, to obtain a clearer relationship between the VSE and ion-pair structures and to
further clarify the origin of IR shifts, we synthesized and studied a new molecular system whose
ion-pair structure is better defined through covalent bonding. The new molecule is named 15-
crown-5-Ph-F1CN where a benzo crown ether (15-crown-5-Ph) is tethered to the phenyl on the
opposite side, away from the nitrile group (Chart 1). The molecule can take either syn or anti
conformations (Supporting Information Figure S8). DFT calculations predict that the syn
conformer is only slightly (< 1 meV) more stable in the neutral molecule and that there is a ~ 150
22
meV torsional energy barrier between the syn and anti conformers. For the anion, the syn
conformers are again only slightly (~3-4 meV) more stable and the syn-anti torsional barrier is
higher, ~ 200-250 meV (Supporting Information Figure S9). Higher energy barriers are
consistent with a similar torsional motion previously reported in structurally related
oligofluorenes.60-61 We did not determine experimentally which conformer is dominant, but
based on the computational results, we depict the molecular structures as syn conformers here.
Figure 6. IR spectra of F1CN-• anions in THF. The blue line represents the FTIR spectrum of the
species obtained by chemical reduction with excess sodium biphenyl, yielding the species
Na(15-crown-5)+-Ph-F1CN (νmax = 2136 cm-1). The FI (black, νmax = 2124 cm-1) and ion pair
(red) spectra were taken with TRIR detection obtained 60 and 100 ns after pulse radiolysis in
neat THF and in the presence of 0.1 M TBA+PF6-, respectively. The ball in the top of the figure
represents Na+.
3.2.2. Intramolecular Ion Pair. Crown ethers are known to encapsulate alkali metal ions,62-
63 and are widely used as sensors for metal ions in materials and biological sciences.64 They were
also employed in the studies of intermolecular ion pairs.52-54 When 15-crown-5-Ph-F1CN is
reduced with sodium or sodium biphenyl in the absence of any other chelating agents, Na+ is
23
expected to be captured by the crown ether within the molecule while the excess electron is
delocalized in the –conjugated part. This ensures that a positive charge is on the opposite side
of the molecule, away from the nitrile group. The complexation forms the neutral spin ½ species
Na(15-crown-5)+-Ph-F1CN-• that we can classify as an intramolecular ion pair (Figure 6 Top).
The nitrile mid-IR peak from the FI of the molecule, (15-crown-5)-Ph-F1CN-•, is shifted
to higher frequency by ~14 cm-1 (νmax = 2124 cm-1) compared to that of F1CN-• due to the further
delocalization of an electron in the –conjugated part of the molecule.28 Upon reduction and
complexation with Na+, the IR absorption from Na(15-crown-5)+-Ph-F1CN-• shifts further blue
by another ~12 cm-1 (νmax = 2136 cm-1). DFT calculations support our experimental observations
(Supporting Information Table S4). Computations predict that the anti conformer exhibits a
larger shift than the syn conformer. When electrochemically reduced in the presence of 0.1 M
TBA+PF6-, it forms ion pairs with TBA+, which cannot fit into the crown, and its IR absorption
band shifts to a lower frequency compared to that of the FI (νmax = 2110 cm-1) as expected. The
spectrum can be decomposed into CIP (15-crown-5)-Ph-F1CN-•, TBA+) and SSIP ((15-crown-
5)-Ph-F1CN-•, THF, TBA+) contributions in a similar manner to the case of F1CN-•. Their
spectroscopic characteristics are summarized in Table 1. In contrast to the IR spectra, the optical
spectra of anions of this molecule can distinguish between the FI and the intramolecular ion pair,
but not clearly between the FI and the intermolecular ion pairs (Supporting Information Figure
S10). We note that a slightly larger shift in the IR peak from the FI was observed for the CIP
with TBA+ (Δ ~-18 cm-1) than for the intramolecular ion pair with Na+ (Δ ~12 cm-1) in ((15-
crown-5)+-Ph-F1CN-•. This result further supports the VSE interpretation of the observed IR
peak shifts upon ion pairings and points to the possibility of calibrating these shifts to obtain
24
what has been called a Stark “tuning rate”,53 meaning the sensitivity of the vibrational frequency
to an external electric field, of ν(C≡N) in anions by careful molecular design.
4. CONCLUSIONS
Ion pairing can control the energetics and dynamics of ions in various environments. In
the first part of the study, we presented evidence for the formation of a SSIP between organic
radicals and non-metal electrolytes in moderately non-polar solvent THF, and distinct spectral
measurements of the CIP, SSIP, and FI using the ν(C≡N) IR absorption in a nitrile-substituted
fluorene. We support our assignment of the SSIP peak with 1) shifting ion pair populations in
DMF/THF mixtures, 2) agreement with the peak position in C222-separated sodium ion pair, and
3) electronic structure calculations. We report the association constants for SSIPs formed by
F1CN-• and TBA+ separated by either a THF or DMF molecule in DMF/THF mixtures of low
DMF composition. Our findings suggest that SSIPs can be formed and exist as
thermodynamically stable species in non-polar environments. In our system, CIPs are slightly
more stable than SSIPs. The existence of the SSIP implies that the redox potentials measured by
electrochemistry in such environments should be considered to reflect the contributions from
both species. In the second part of the study, we show that we can induce a blue-shift in the
ν(C≡N) IR band of a nitrile-substituted fluorene anion by covalently tethering it to a metal-
chelating ligand. This adds further support that the induced shift in IR absorptions by ion pairing
originates from the vibrational Stark effect. These results combined show the capability of
identifying ion-pair structures based on vibrational information.
25
ASSOCIATED CONTENT
Supporting Information. Computational details, detailed description of contribution of ion
pairings in redox potentials, Figures S1-S10, and Tables S1-4. This material is available free of
charge via the Internet at http://pubs.acs.org.
AUTHOR INFORMATION
Corresponding Author
ACKNOWLEDGMENT
This material, and use of the LEAF facility of the BNL Accelerator Center for Energy
Research and the computer Cluster at the Center for Functional Nanomaterials, is supported by
the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Division of
Chemical Sciences, Geosciences & Biosciences under contract no. DE-SC0012704. This project
was also supported in part by the U.S. Department of Energy, Office of Science, Office of
Workforce Development for Teachers and Scientists (WDTS) under the Science Undergraduate
Laboratory Internships Program (SULI). T.M. is grateful for the support by the Goldhaber
Distinguished Fellowship from Brookhaven Science Associates. The authors thank Bobby H.
Layne for technical assistance in LEAF experiments and Dr. Qin Wu for his assistance in
quantum calculations.
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