On the nature of solvent effects on redox properties

Article abstract of DOI:10.1021/jp031268q

The one-electron reduction potentials of six radical cations, four cations, and four neutral radicals in tetrahydrofuran, dichloromethane, dimethyl sulfoxide, N-methyl-2-pyrrolidinone, N,N-dimethylformamide, acetonitrile, methanol, ethanol, 2-propanol, acetone, formamide, and 1,1,1,3,3,3-hexafluoropropan-2-ol were measured by cyclic voltammetry. For 10 of the redox couples, the redox process was reversible in all solvents. The results were used to evaluate solvent effects using Kamlet-Taft relationship. The entropy contribution to some of the observed solvent effects were examined by measuring the redox potentials as a function of temperature. The absolute value of the entropy increased with increasing hydrogen bond donor ability of the solvent. The variation in entropy indicated that specific solvation was very important when considering solvent effects on redox characteristics.

Full text of DOI:10.1021/jp031268q

J. Phys. Chem. A 2004, 108, 4805-4811
4805
On the Nature of Solvent Effects on Redox Properties
Heidi Svith,† Henrik Jensen,† Johan Almstedt,‡ Paula Andersson,‡ Thomas Lundba1ck,§
Kim Daasbjerg,*,† and Mats Jonsson*,‡,
Department of Chemistry, Aarhus UniVersity, DK-8000 Århus C, Denmark, Department of Chemistry,
Nuclear Chemistry, Royal Institute of Technology, SE-100 44 Stockholm, Sweden, and
BioVitrum AB, SE-112 76 Stockholm, Sweden
ReceiVed: NoVember 24, 2003; In Final Form: March 2, 2004
The one-electron reduction potentials of six radical cations, four cations, and four neutral radicals in
tetrahydrofuran, dichloromethane, dimethyl sulfoxide, N-methyl-2-pyrrolidinone, N,N-dimethylformamide,
acetonitrile, methanol, ethanol, 2-propanol, acetone, formamide, and 1,1,1,3,3,3-hexafluoropropan-2-ol have
been measured by cyclic voltammetry. For 10 of the redox couples, the redox process was reversible in all
solvents. These results have been used to evaluate solvent effects by means of the Kamlet-Taft relationship.
The relative importance of the solvent parameters R, â, π*, and δH is 54.9, 9.6, 15.5 and 20.0%, respectively,
for the radical cation displaying the strongest solvent dependence. In addition, we have studied the entropy
contribution to some of the observed solvent effects by measuring the redox potentials as a function of
temperature. The absolute value of the entropy appears to increase with increasing hydrogen bond donor
ability of the solvent. The variation in entropy indicates that specific solvation is of main importance when
considering solvent effects on redox properties.
Introduction
In practice, solvent effects on redox properties are usually
quantified by one-electron reduction potentials measured against
Solvent effects on reaction kinetics and mechanisms have
been a subject of interest for a number of years.1 More recently,
studies of solvent effects on redox properties of radicals and
radical ions have occurred in the literature.2,3 Quantitative
descriptions as well as fundamental understanding of the nature
of solvent effects on redox properties are very important for
scientists trying to understand the chemistry of more complex
systems where direct measurements are difficult or even
impossible. In addition, improved understanding of solvent
effects is useful when making quantitative comparison of redox
data from different solvents. Solvent effects on one-electron
reduction potentials (eq 1) are a measure of the solvent
dependence on the difference in free energy of solvation for a
given redox couple O/R (eq 2).4
a reference redox couple for which the solvent sensitivity is
assumed to be very small. Ferrocene is one possible candidate.
Interestingly, a recent study on the redox properties of substi-
tuted ferrocenes indicates that the solvent effects are substantial.7
However, the potentials were determined directly against the
reference electrode without taking differences in liquid junction
potential into account. Other studies display a weaker solvent
dependence on the redox properties of ferrocene.8
Properties in solution, e.g., solubility, rates of reactions and
free energy, and enthalpy of equilibria, can often be described
by so-called linear free energy relationships (LFER) or linear
solvation energy relationships (LSER).9 One of the most
successful relationships has been found to be the Kamlet-Taft
expression (eq 3), where XYZ is the property of interest, XYZ0,
a, b, s, and h are solvent independent coefficients characteristic
of the process, R is the hydrogen bond donor (HBD) ability of
the solvent, â is the hydrogen bond acceptor (HBA) or electron
pair donor ability to form a coordinative bond, π* is its
dipolarity/polarizability parameter, and δH is the Hildebrand
solubility parameter which is a measure of the solvent-solvent
interactions that are interrupted in creating a cavity for the
solute.9,10
O + e- h R
(1)
(2)
∆G° (R) - ∆G° (O)
solv
solv
IP ≈ C + E° +
F
This can be understood from the above approximate relation
between E° and the corresponding gas-phase ionization potential,
IP, where the constant C is the absolute potential of the reference
electrode in a given solvent (e.g., 4.44 ((0.02) eV5 for the
hydrogen electrode in water), ∆G°solv (R) and ∆Gs°olv (O) are the
free energies of solvation of species R and O, respectively, and
F is the Faraday constant.6 It should be noted that the ionization
potential is the enthalpy of ionization at 0 K; thus, the ionization
entropy and temperature correction are neglected in eq 2.
However, these corrections are assumed to be fairly small.
XYZ ) XYZ0 + aR + bâ + sπ* + hδH
(3)
For some processes, any of the coefficients XYZ0, a, b, s, and/
or h may be negligibly small, so that the corresponding terms
do not play a role in the characterization of the solvent effects
for these processes. This approach has been criticized for not
separating specific and nonspecific effects.11 Alternative ap-
proaches which separate specific and nonspecific effects have
also been elaborated, e.g., by Koppel and Palm1,11,12 and more
recently by Drago and co-workers.13,14
† Aarhus University.
‡ Royal Institute of Technology.
§ Biovitrum AB, SE-112 76 Stockholm, Sweden
E-mail: matsj@nuchem.kth.se. Telephone: +46 8790 9123. Fax: +46
8790 8772.
The Kamlet-Taft expression has been found to describe
solvent effects on one-electron reduction potentials of dications,
10.1021/jp031268q CCC: $27.50 © 2004 American Chemical Society
Published on Web 05/05/2004

4806 J. Phys. Chem. A, Vol. 108, No. 21, 2004
Svith et al.
TABLE 1: Reduction Potentials (V vs Fc+/Fc) for Substances 1-14 Measured by Ordinary Cyclic Voltammetry in Different
Solvents Containing 0.1 M Bu4NBF4 Unless Otherwise Noteda
solvent
THF
1b
2c
3b
4c
5b
6b
7c
8c
9c
10c
11c
12c
13c
14c
0.549 0.236 0.933 -0.241 0.385 0.978 -1.190 -1.095 -1.403 -0.601 -2.044 -1.748 -1.904 -1.655
CH2Cl2 0.581d 0.317 0.904e -0.278 0.338 0.983 -1.147 -1.021 -1.430 -0.653 -1.989 -1.693 -1.948 -1.677
DMSO 0.633 0.253 0.886 -0.275 0.376 0.880 -1.193 -1.091 -1.424 -0.652 -1.947 -1.720 -1.819 -1.568
NMP
DMF
MeCN
MeOH
EtOH
0.605 0.222 0.883 -0.278 0.399 0.963 -1.190 -1.087 -1.428 -0.645 -2.039 -1.793 -1.868 -1.622
0.573 0.232 0.843 -0.277 0.371 0.872 -1.194 -1.081 -1.433 -0.650 -2.016 -1.760 -1.871 -1.623
0.624 0.314 0.915 -0.279 0.367 1.005 -1.171 -1.051 -1.438 -0.648 -1.984 -1.697 -1.883 -1.623
0.495 0.438 0.893 -0.279 0.353 0.939 -1.144 -0.992 -1.440 -0.661 -1.784f -1.184g -1.691 -1.556
0.484 0.439 0.959 -0.275 0.379 0.963 -1.156 -0.997 -1.439 -0.648 -1.899h -1.354h -1.922i
2-PrOH 0.500 0.388 1.105 -0.239 0.424 0.999 -1.179 -1.015 -1.423 -0.645
acetone 0.599 0.258 0.913 -0.290 0.355 0.999 -1.183 -1.079 -1.432 -0.642 -2.061 -1.718 -1.922 -1.668
FA 0.640 0.409 0.927 -0.266 0.403 0.913 -1.135 -0.969 -1.421 -0.651
-1.494
a The following abbreviations have been used: THF ) tetrahydrofuran (with 0.3 M Bu4NBF4), DMSO ) dimethyl sulfoxide, NMP ) N-methyl-
2-pyrrolidinone, DMF ) N,N-dimethylformamide, MeCN ) acetonitrile, MeOH ) methanol, EtOH ) ethanol, 2-PrOH ) 2-propanol (with 0.2 M
Bu4NBF4), and FA ) formamide (with 0.2 M Bu4NBF4). The concentration of Bu4NBF4 in CH2Cl2 and acetone was 0.2 M. b Peak oxidation
potentials measured at 0.5 V s-1 with a standard deviation of (5 mV. c Standard potentials determined with an uncertainty of (3 mV unless
otherwise noted. d The standard potential was determined to be 0.549 V vs Fc+/Fc. e The standard potential was determined to be 0.866 V vs
Fc+/Fc. f Determined by fast cyclic voltammetry with the switch potential set at the background reduction; uncertainty ) 37 mV. g Determined by
normal cyclic voltammetry with the switch potential set at the background reduction. h Determined by fast cyclic voltammetry with an uncertainty
of 10 mV. i Determined by fast cyclic voltammetry with an uncertainty of 19 mV.
TABLE 2: Kamlet-Taft Solvent Parameters10,23
CHART 1
solvent
R
â
π*
δH
THF
DMSO
NMP
0
0
0
0
0.19
0.98
0.86
0.76
0.08
0.71
1.96
0.55
0.76
0.77
0.69
0.40
0.66
0.75
0.84
0.43
0.48
0
0.58
1
9.1
12
0.92
0.88
0.75
0.60
0.54
0.48
0.71
0.97
0.65
11.3
12.1
11.9
14.5
12.7
11.5
9.9
DMF
MeCN
MeOH
EtOH
2-PrOH
acetone
FA
19.2
9.72
HFP
it is plausible that different families of compounds may follow
slightly different trends due to differences in solvation of the
neutral species. Solvents with more extreme properties (e.g.,
hexafluoro-2-propanol, which has very high hydrogen bond
donor ability) can be expected to deviate from Kamlet-Taft
expressions due to significantly stronger solvation of the neutral
species.
To fully understand solvent effects on redox properties, the
nature of the solvent effects, i.e., the enthalpic and entropic
contributions to the variation in solvation free energies between
different solvents must be known. This has to our knowledge
never been studied explicitly, although Svaan and Parker have
studied the entropy contribution to redox potentials in a few
aprotic solvents.20 In many cases, where free energies expressed
by redox potentials have been used to calculate enthalpies (e.g.,
bond dissociation enthalpies) via thermochemical cycles, the
entropy contribution to the redox potential has been assumed
to be negligible.22 Yet, other studies have shown that the entropy
changes for electron-transfer processes are directly related to
the solvation energy changes.15-21
In this work, we have studied the solvent effects on the one-
electron reduction potentials of the six radical cations, four
cations, and four neutral radicals depicted in Chart 1. In addition,
we have made an attempt to elucidate the entropy contribution
to some of the observed solvent effects by measuring the redox
potentials as a function of temperature.
radical cations, neutral molecules, and radical anions fairly well.3
However, previous studies have included relatively few sub-
stances, and it has therefore been difficult to draw reliable
general conclusions from the trends observed. Furthermore,
some of the trends are based on irreversible peak potentials and
may therefore not describe solvent effects on the thermodynami-
cal potentials accurately. The magnitude of the solvent effects
has been found to depend roughly on the charge localization
reflected by the gas-phase ionization potential or electron
affinity.3 In addition, a series of studies by Svaan and Parker
have shown that the entropy as well as the enthalpy contribution
to the redox potential closely follow the charge localization.15-21
This indicates that the relatively weak solvation of neutral
molecules and radicals can be regarded as solvent independent
and the main contribution to the solvent effect originates from
differences in ion solvation. As the observed trend is quite rough,
Results and Discussion
The measured potentials and the Kamlet-Taft parameters for
the solvents used in this study are collected in Tables 1 and 2,
respectively.

Solvent Effects on Redox Properties
J. Phys. Chem. A, Vol. 108, No. 21, 2004 4807
TABLE 3: Kamlet-Taft Coefficients for Substances 2, 4, and 7-10
substance
XYZ0
a (R)
b (â)
s (π*)
h (δH)
R2
Fa
2
4
7
8
9
10
0.37 ( 0.07
-0.20 ( 0.06
-1.19 ( 0.01
-1.13 ( 0.02
-1.45 ( 0.03
-0.62 ( 0.03
0.49 ( 0.02
0.12 ( 0.02
0.05 ( 0.02
0.07 ( 0.03
-0.09 ( 0.04
-0.11 ( 0.04
-0.22 ( 0.06
-0.06 ( 0.05
-0.05 ( 0.02
-0.02 ( 0.02
0.04 ( 0.03
0.47 ( 0.08
0.16 ( 0.07
0.02 ( 0.03
-0.04 ( 0.05
-0.14 ( 0.07
-0.20 ( 0.07
-0.04 ( 0.005
-0.02 ( 0.004
0.001 ( 0.003
0.008 ( 0.004
0.01 ( 0.005
0.01 ( 0.005
0.99
0.92
0.97
0.99
0.53
0.76
247
17
39
89
1.4
0.03 ( 0.03
4
a The F statistic, or the F-observed value can be used to determine whether the observed relationship between the dependent and independent
variables occurs by chance.
n
Solvent Effects on Electrode Potentials for Compounds
(R - Rj)2
1)10. First of all it should be noted that the electrode pro-
cess studied for the neutral compounds 1-6 is the oxidation to
∑
i
i)1
a′ ) |a|
b′ ) |b|
s′ ) |s|
(4)
(5)
(6)
(7)
the pertinent radical cations while it for the cations 7-10 it is
the reduction to the pertinent radicals. In the case of 2, 4, and
7-10 it was possible to attain reversibility in the cyclic
voltammetry experiments for all solvents leading to the direct
measurement of standard potentials. The solvent 1,1,1,3,3,3-
hexafluoropropan-2-ol (HFP) was included in a few measure-
ments where the standard oxidation potentials for 1 and 4 were
determined to be 0.695 and -0.008 V vs Fc+/Fc (abbreviation
for ferrocenium/ferrocene), respectively, and the peak oxidation
potential for 2 and 5 to be ca. 1.3 and 0.62 V vs Fc+/Fc,
respectively. In addition, the standard potential for 3 in
HFP has been determined to be 1.48 V vs Fc+/Fc by Eberson
et al.24,25
n
(E 0 - Eh0)2
∑
i
x
i)1
n
(â - âh)2
∑
i
i)1
n
(E 0 - Eh0)2
∑
i
x
i)1
n
(π* - πj*)2
∑
i
i)1
As can be seen in Table 1, the one-electron oxidation po-
tential for 1, 2, 3, 5 and 6 vary significantly with solvent.
Moreover, when including HFP we see that even the rela-
tively solvent insensitive radical cation of 4 is strongly af-
fected. The quite dramatic effect of HFP can only be attri-
buted to its extremely high hydrogen bond donor ability
(Table 2). In this context, however, it should be emphasized
that the potentials listed for 1, 3, 5, and 6 are irrever-
sible, meaning that they are affected to an unknown extent by
the presence of homogeneous kinetics and/or charge-
transfer kinetics. Still, the results seem to be in accordance with
the expectation that the radical cation of 4 as the most
delocalized system is the one least influenced by solvent. As
to the cations the standard potentials for 7-10 appear to be
more or less solvent independent although measurements were
not carried out in HFP in these cases. Thus, we can conclude
that the solvation of these cations is relatively weak or at least
similar to that of ferrocenium. This is attributed to the fact that
the positive charge is shielded by the methyl and/or phenyl
groups.
n
(E 0 - Eh0)2
∑
i
x
i)1
n
(δ - δh )2
∑
Hi
H
i)1
h′ ) |h|
n
(E 0 - Eh0)2
∑
i
x
i)1
The relative importance of a given parameter can be calculated
from the beta coefficients using eq 8.
a′
aj )
(8)
a′ + b′ + s′ + h′
This equation gives the relative importance of R, the hydrogen
bond donor ability of the solvent. For compound 2 (displaying
the strongest solvent dependence) the relative importances of
R, â, π*, and δH are 54.9, 9.6, 15.5, and 20.0%, respectively,
when HFP is included and 52.9, 8.9, 22.2, and 16.0%,
respectively, when HFP is excluded. Hence, the most important
solvent parameter is R while â appears to be of minor
importance.
The redox data for compounds 2, 4, and 7-10 were all
analyzed in terms of the Kamlet-Taft expression given in eq
3. The resulting coefficients are given in Table 3.
This trend is similar to that previously found for the reduction
potentials of neutral species but it clearly deviates from that
found for the oxidation of a compound such as 2 (DABCO).3
In this context, it is important to note that the previous set of
oxidation potentials of 2 was based on irreversible measurements
in a limited set of solvents.3 However, the Kamlet-Taft
relationship based on the reversible potentials obtained herein
becomes similar to the previous one, if the solvent effects are
only quantified in the solvent set used in ref 3. Thus, we can
conclude that the main reason for the discrepancy observed for
2 is the different sets of solvents considered. Admittedly, the
The relative importance of the different solvent properties
can be evaluated from so-called beta coefficients derived
according to eq 4-7 where a′, b′, s′, and h′ are the partial
regression coefficients (or “â coefficients”),|a|, |b|, |s|, and |h|
are the absolute values of the regression coefficients, Ri, âi, π*i,
and δH are the Kamlet-Taft parameters of a given solvent (i),
Rj, âh, πji*, and δhH are the average values of these quantities in a
given set of solvents, Ei0 is the potential measured in a given
solvent and Eh° is the average value of the potentials in a given
set of solvents.26

4808 J. Phys. Chem. A, Vol. 108, No. 21, 2004
Svith et al.
TABLE 4: Ionization Potentials and Solvation Free Energies (Relative to Fc+/Fc) for Substances 2-5 and Fc
a
∆∆Gs°olv (kJ mol-1
)
substance
IP (eV)26
THF
CH2Cl2
DMSO
NMP
DMF
MeCN
MeOH
EtOH
2-PrOH
acetone
FA
2
3
4
5
7.321
7.56
6.1
7.12
6.71
36.2
-8.0
-35.6
2.4
28.4
-5.2
-32.0
6.9
34.5
-3.5
-32.2
3.3
37.5
-3.2
-32.0
1.1
36.6
0.7
28.7
-6.3
-31.9
4.1
16.6
-10.5
-32.3
3.0
16.6
-10.5
-32.3
3.0
21.5
-24.6
-35.8
-1.4
0
34.1
-6.1
-30.9
5.3
19.5
-7.4
-33.2
0.7
-32.1
3.8
Fc
0
0
0
0
0
0
0
0
0
0
∆∆
∆
)
Gs°olv(R) - ∆Gso°lv(O).
a
G°
solv
are known are 2-6 and therefore we can only base the analysis
on these substances. However, the 4-nitroaniline radical cation
6 can undergo deprotonation and therefore it is excluded from
the analysis. Using eq 2, we can calculate the difference in free
energy of solvation, ∆∆G° , between the neutral species and
solv
the corresponding cation. To avoid errors due to uncertainties
in the absolute potential for the reference electrode in the
different solvents, we have calculated ∆∆G° for substances
solv
2-5 relative to ∆∆G° for Fc+/Fc. These values are given in
solv
Table 4 along with the corresponding gas-phase ionization
potentials.
When plotting ∆∆G° against the ionization potential we
solv
can make the general observation that ∆∆G° indeed in-
solv
creases with increasing ionization potential (i.e., charge localiza-
tion) as shown for acetonitrile in Figure 2; similar correlations
can be obtained for the other solvents used in this study.
However, the compound 3 (i.e., 1,4-dimethoxybenzene) deviates
significantly from the general trend. This deviation is much
larger than the uncertainty introduced in the electrochemical
determination due to the use of irreversible peak potentials.
Moreover, we see no similar deviation for compound 5, the
measurements of which also were based on irreversible poten-
tials. The deviation indicates that the suggested proportionality
Figure 1. Experimentally determined oxidation potentials for DABCO
(2) plotted against the corresponding values estimated using the
Kamlet-Taft relationship given in Table 3: (a) DMF, (b) THF; (c)
NMP, (d) DMSO, (e) MeCN, (f) acetone, (g) FA, (h) 2-PrOH, (i) EtOH,
(j) MeOH, and (k) HFP.
strong dependence on the hydrogen bond donor ability originates
from the high potentials determined in HFP, but even when the
Kamlet-Taft relationship for DABCO is determined without
including the value in HFP, the strong dependence on R is
present. Furthermore, the relative importance of R does not
change significantly when excluding the value in HFP. It is quite
plausible that bases like DABCO may be protonated in HFP.
Measurements of the conductivity of HFP and MeCN containing
DABCO show that DABCO increases the conductivity of HFP
significantly while the conductivity of MeCN remains virtually
unchanged. This indicates that DABCO is protonated to some
extent in HFP. However, this does not appear to cause any
significant deviation from the Kamlet-Taft correlation. In
Figure 1, the experimentally determined oxidation potentials of
2 are plotted against the corresponding values estimated using
the Kamlet-Taft relationship. As can be seen, the Kamlet-
Taft relationship successfully describes the solvent dependence
on the one-electron oxidation potential of 2.
The reason for mainly focusing on the oxidation of 2 in the
above discussion is that this radical cation shows the most
pronounced solvent dependence and that the potentials are
reversible. The potentials for 4 (TMPD) as well as 7-10 are
also reversible but the solvent dependence is relatively weak
with the notable exception of the value for 4 determined in HFP.
We also measured the peak oxidation potentials for 2 in the
solvents given in Table 2, and it was found that the irreversible
potentials essentially showed the same solvent dependence as
the reversible potentials although to a slightly larger extent. The
irreversible potentials determined for 1, 3, and 5 could thus
describe the solvent dependence for these substances at least
qualitatively correctly.
between the ionization potential and ∆∆G° cannot be taken
solv
for granted and that the presence of different heteroatoms may
cause deviations from the general trend.
The suggested relationship between the solvent effect on E°
and ∆∆Gs°olv cannot be verified from the present data, although
it is obvious that the most strongly solvated cation 2 also
displays the strongest solvent sensitivity (judging from the K-T
relationships). It should be noted that compound 3, which is
considerably less solvated, also displays a significant solvent
sensitivity. The difference between the highest and the lowest
potential is in fact larger for compound 3 than for compound
2. Again, this could suggest that the solvent dependence is
strongly governed by the nature of the heteroatom. However, it
should be noted that the potentials for compound 3 are
irreversible and that the irreversible potentials for 2 were shown
to be more solvent sensitive than the corresponding reversible
potentials.
Solvent Effects on Electrode Potentials for Compounds
11)14. In these cases the electrode process studied is the
reduction of the radicals to the pertinent anions. Table 5
comprises the Kamlet-Taft coefficients obtained for 11-14.
As can be seen, the quality of the Kamlet-Taft relationships
is relatively low and there is no significant dependence of E°
on the Kamlet-Taft parameters within experimental error. As
previously shown for reduction of neutral species, the hydrogen
bond donor ability, R, and to some extent also the dipolarity/
polarizability, π*, of the solvent are of major importance. The
small coefficients and the poor correlations found for substances
11-14 can probably be attributed to the structure of these
compounds. They contain bulky substituents that could shield
the charge and thereby weaken the solvation of the ion. This
Qualitatively, it would be expected that the solvent effect
increases with increasing charge localization of the correspond-
ing cation. The only substances for which the ionization
potentials (assumed to reflect the degree of charge localization)

Solvent Effects on Redox Properties
J. Phys. Chem. A, Vol. 108, No. 21, 2004 4809
Figure 2. Solvation free energies (∆∆G° ) for substances 2-5 and Fc relative to Fc+/Fc plotted against the corresponding ionization potentials.
solv
TABLE 5: Kamlet-Taft Coefficients for Substances 11-14
substance
XYZ0
a (R)
b (â)
s (π*)
h (δH)
R2
F
11
12
13
14
-2.2 ( 0.4
-2.3 ( 0.2
-2.3 ( 0.3
-1.8 ( 0.1
0.2 ( 0.3
0.1 ( 0.2
-0.06 ( 0.36
0.06 ( 0.14
-0.07 ( 0.25
0.02 ( 0.15
-0.02 ( 0.26
0.07 ( 0.10
0.17 ( 0.57
-0.02 ( 0.36
-0.05 ( 0.62
0.08 ( 0.22
0.009 ( 0.06
0.03 ( 0.04
0.04 ( 0.06
0.01 ( 0.02
0.90
0.90
0.54
0.85
4.3
6.8
0.9
2.8
would also make the redox properties less solvent sensitive,
which is exactly what we observe.
As discussed for cations, general equations describing solvent
Interestingly, the absolute value of the entropy appears to
increase with increasing hydrogen bond donor ability of the
solvent going from dimethyl sulfoxide and acetonitrile to
methanol. For a series of structurally similar substances in a
given solvent, Svaan and Parker have shown that the enthalpic
contribution to the potential is proportional to the entropic
contribution.16 Thus, strong solvation of an ion is accompanied
by a large decrease in entropy as also seen herein for the
delocalized ions of nitrobenzene 4-nitrobenzonitrile, 2,5-di-
methylbenzoquinone, and 3,5-di-tert-butyl-1,2-benzoquinone.
This is also what we would expect from Born’s model for ion
solvation. In Born’s model for solvation of small ions the free
energy of solvation as well as the entropy of solvation are
proportional to z2/r where z is the charge and r is the ionic radius.
It should be noted that this purely electrostatic model does not
describe solvent effects on one-electron reduction potentials of
organic substances successfully. Furthermore, the entropy for
compound 3 is relatively small, which is well in line with the
low free energy of solvation observed in this work.
When analyzing the data in Table 6, it is obvious that the
difference in entropy between two solvents is larger than the
apparent difference in solvation free energy for a given solute.
Hence, the change in entropy is to some extent compensated
by a change in enthalpy. Consequently, the main contribution
to the observed solvent effects is a change in the relative
importance of the entropic and the enthalpic contribution to the
free energy of solvation. This, in turn, is strongly governed by
the hydrogen bond donating ability of the solvent. The variation
in entropy indicates that specific solvation is of main importance
when considering solvent effects on redox properties.
effects on redox properties based on ∆∆G° or EA only, are
solv
not to be expected for the reduction of neutral species either.
Equations describing the variation in the Kamlet-Taft coef-
ficients as a function of ∆∆G° or EA has been derived for
solv
series of structurally similar substances, e.g., nitrobenzenes.3
Differences in Solvation between Anions and Cations. The
standard potentials for substances 7-14 constitute an interesting
set of data since we here have the possibility of elucidating the
solvent effects on both the reduction and the oxidation potentials
for four different neutral radicals. In addition, this set of data,
in principle, also allows the evaluation of the difference in
solvation between cations and anions with identical structure
as a function of solvent and completely independent of the
radicals. From the above treatment of reduction potentials for
cations and neutral species, we can see that the main contribution
to the solvent effects on substances 7-14 originates from the
reduction of neutral species. It should be noted that the absence
of solvent effects on the reduction potentials of the cations does
not imply that the solvation energies are solvent independent.
Formally, the only conclusion that can be drawn from this is
that solvation of these cations follows that of ferrocene. From
the two sets of potentials it is also possible to derive the
disproportionation equilibrium constants for the radicals, but
as it seems to present no obvious correlation with the solvent
parameters, we did not consider it further.
The Thermodynamic Nature of Solvent Effects. To assess
the entropy and enthalpy contribution to the observed solvent
effects on redox properties (relative to ferrocene) we measured
the standard potential for a number of substances in a few
solvents with different properties as a function of temperature
(∂G/∂T ) -S w ∆S° ) F dE°/dT).28 The temperature intervals
were 5-80 °C for MeCN, 20-60 °C for DMSO, and 5-60 °C
for MeOH. The resulting temperature effects on the reversible
potentials and the calculated entropies are given in Table 6.
As expected the entropy for reduction of neutral species is
negative while for the reduction of cations it is positive.15-21
Experimental Section
Reagents. The supporting electrolyte tetrabutylammonium
tetrafluoroborate (Bu4NBF4) was prepared by standard proce-
dures. Ferrocene (98%) was purchased at Aldrich and used as
received. Tetrahydrofuran (Lab-Scan) was distilled over sodium
and benzophenone under a nitrogen atmosphere. Dichlo-
romethane (Gropa) and acetonitrile (Lab-Scan) were both

4810 J. Phys. Chem. A, Vol. 108, No. 21, 2004
Svith et al.
TABLE 6: Slopes of ∆E°/∆T and Entropies ∆S° for the Reduction of Different Substrates in MeOH, MeCN, and DMSO
Relative to Fc+/Fc
MeOH
slope (mV K-1) ∆S° (J mol-1 K-1
(12.68)a
MeCN
slope (mV K-1) ∆S° (J mol-1 K-1
(48.18)a
DMSO
∆S° (J mol-1 K-1
)
substance
)
)
slope (mV K-1
)
Fc+
(52.38)a
-108 ( 7
(-56)b
nitrobenzene
-4.0 ( 0.2
-386 ( 19
(-373)b
-579 ( 96
(-566)b
-579 ( 193
(-566)b
-289 ( 96
(-276)b
48 ( 10
(61)b
-1.25 ( 0.3
-1.12 ( 0.3
-1.4 ( 0.2
-1.3 ( 0.1
0.1 ( 0.1
-120 ( 29
-1.12 ( 0.07
-1.4 ( 0.2
-0.3 ( 0.1
-0.5 ( 0.3
(-72)b
4-nitrobenzonitrile
2,5-dimethyl-benzoquinone
-6 ( 1
-108 ( 29
(-60)b
-138 ( 19
(-86)b
-6 ( 2
-133 ( 19
-29 ( 10
(23)b
(-85)b
3,5-di-tert-1,2-benzoquinone -3 ( 1
-124 ( 10
(-76)b
-48 ( 29
(4)b
3c
4
0.5 ( 0.1
10 ( 10
(58)b
-0.56 ( 0.04
0.9 ( 0.2
-54 ( 4
0.01 ( 0.03
1.1 ( 0.4
1 ( 3
(-41)b
(49)b
7
87 ( 19
(100)b
106 ( 39
0.4 ( 0.2
39 ( 19
(91)b
(154)b
8
0.8 ( 0.4
77 ( 38
0.5 ( 0.1
48 ( 10
(96)b
0.2 ( 0.1
19 ( 10
(90)b
(71)b
9
0.9 ( 0.3
87 ( 29
(100)b
0.1 ( 0.1
10 ( 10
0.2 ( 0.1
19 ( 10
(71)b
(58)b
10
11
12
13
14
0.8 ( 0.2
77 ( 19
0.2 ( 0.3
19 ( 29
(67)b
0.4 ( 0.1
39 ( 10
(90)b
(91)b
-1.1 ( 0.7
-2.6 ( 1.0
-0.4 ( 0.2
-1.3 ( 1.0
-106 ( 68
(-93)b
-0.9 ( 0.1
-1.3 ( 0.2
-0.4 ( 0.1
-0.5 ( 0.1
-87 ( 10
-0.5 ( 0.2
-0.6 ( 0.1
-0.6 ( 0.1
-0.3 ( 0.1
-48 ( 19
(4)b
(-39)b
-251 ( 96
(-238)b
-39 ( 19
(-26)b
-125 ( 19
(-77)b
-58 ( 10
(-6)b
-39 ( 10
-58 ( 10
(-6)b
(9)b
-125 ( 96
-48 ( 10
(0)b
-29 ( 10
(-112)b
(23)b
a From ref 8. b Absolute values calculated from the absolute entropy for Fc+/Fc given in ref 8. c On the basis of reversible potentials measured
at a sweep rate of 2 V s-1
.
distilled over CaH2, followed in the case of acetonitrile by
successive distillation over P2O5 and K2CO3. The solvent
1-methylpyrrolidin-2-one (Merck-Schuchardt) was distilled
under a nitrogen atmosphere. Dimethyl sulfoxide (Aldrich, sure
seal bottle) was stored under nitrogen. All solvents including
N,N-dimethylformamide (Lab-scan), methanol (Merck), 99.9%
ethanol (DDSF), and 1,1,1,3,3,3-hexafluoropropan-2-ol (Lan-
caster) were dried over activated alumina (ICN Alumina-I-
Super) prior to use and handled with normal syringe techniques.
electrode applied was Ag/AgI, I- ) 0.1 M, but all potentials
measured were also referenced against the ferrocenium/ferrocene
(Fc+/Fc) redox couple. A platinum coil was used as counter
electrode. The ohmic drop was compensated with a positive
feedback system incorporated in the home-built potentiostat. In
fast cyclic voltammetry (500 V s-1 < ν < 3 kV s-1) the working
electrode was a gold ultra-microelectrode (Ø ) 25 µm).
Procedures. In normal cyclic voltammetry standard potentials
were measured as the midpoint between the cathodic and anodic
peak potentials against the Fc+/Fc redox couple for three
different sweep rates, typically 0.1, 1, and 10 V s-1. All systems
were well-behaved exhibiting a peak separation of less than 100
mV. The same procedure was employed in the fast cyclic
voltammetric experiments using sweep rates until 3 kV s-1 but
with a substantially larger uncertainty as the peak separation
amounted to as much as 100-400 mV. The measurements were
not influenced by the working electrode material. For instance,
the value of the standard potential of 8 in methanol was found
to be essentially the same at both glassy carbon (E° ) -1.000
V vs Fc+/Fc) and gold (E° ) -0.992 V vs Fc+/Fc). In THF,
the corresponding numbers were -1.099 and -1.095 V vs Fc+/
Fc, respectively. For compounds 1, 3 and 5 the voltammograms
appeared irreversible at all accessible sweep rates and accord-
ingly we have listed the peak potentials obtained at ν ) 0.5 V
s-1 in these cases. The standard potentials of the radical species
11-14 were measured from the second waves recorded in the
cyclic voltammograms of the corresponding cationic compounds
7-10. The reduction wave pertaining to 12 was irreversible in
methanol and ethanol unless the switch potential was adjusted
at the background reduction. Apparently, the use of such a large
overpotential increased the basicity of the solution at the vicinity
of the electrode surface thereby preventing the protonation of
the anion.
Tetramethoxydibenzofuran (1) was a generous gift from Dr.
Jonas Hellberg (Stockholm); the synthesis procedure is described
in ref 29. 1,4-Diazabicyclo[2.2.2]octane (DABCO) (2), 1,4-
dimethoxybenzene (3), N,N,N′,N′-tetramethylphenylene-1,4-
diamine (TMPD) (4), N,N-dimethylaniline (5), and 4-nitroaniline
(6) were of the purest grade available (Lancaster and Aldrich)
and used as supplied. The preparation of 4-methoxycarbonyl-
1-methylpyridinium perchlorate (7) is described in ref 30.
4-Benzoyl-1-methylpyridinium perchlorate (8) was synthesized
by reacting isonicotinophenone with dimethyl sulfate. The
compounds 2,4,6-triphenylthiopyrylium perchlorate (9) and
2,4,6-triphenyl-1-methylpyridinium perchlorate (10) were both
prepared from 2,4,6-triphenylpyrylium perchlorate as described
in ref 31. To obtain 9, 2,4,6-triphenylpyrylium perchlorate was
treated with sodium sulfide,32 while in the case of 10 it was
allowed to react with glycine.33 The neutral radicals 11-14 were
generated electrochemically during the cyclic voltammetric scan
from the pertinent cations 7-10, i.e., the redox features of 11-
14 could be characterized by studying the second redox waves
of 7-10. All other compounds were of commercial origin.
Instrumentation. In the normal cyclic voltammetric experi-
ments (sweep rate ν < 30 V s-1) a three-electrode setup was
employed with a glassy carbon (o.d. ) 1 mm) or a gold electrode
(o.d. ) 1 mm) serving as working electrode. The reference

Solvent Effects on Redox Properties
J. Phys. Chem. A, Vol. 108, No. 21, 2004 4811
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Acknowledgment. T.L. and M.J. thank the Carl Trygger
Foundation for financial support. K.D. thanks the Statens
Naturvidenskabelige Forskningsråd for financial support.
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