Molecular Electrochemistry of Coordination Compounds …
413
for a given class of compounds have been established. This methodology, however,
does not account for a couple of factors that may also affect the redox potential.
One of these factors is the ion energy lowering caused by ionic strength, induced
by base electrolyte (also called supporting electrolyte; it is used in large excess to
ensure diffusion-controlled transport of electroactive species, thus making migration
in the electric field negligible). As the following example shows, this stabilisation
energy is important for non-polar solvents. In these solvents, it comes to the association of the base electrolyte and this should be taken into account, when calculating
the ionic strength. For n-Bu 4 NPF 6 in dichloromethane, the fraction of free ions in
0.1 M solution in CH 2 Cl 2 is equal to only 7.7% [14]. The resulting ionic strength
would cause Fc
+ /Fc potential shift to a value lower by 53 mV, as estimated from
the extended Debye–Hückel law. For more polar solvents, the shift would be much
smaller, even, if assumed that there would not come to any association in these polar
solvents. Another approach is measuring the variation in the dielectric constant (relative permittivity) value effected by the presence of electrolytes. Using dyes that
show known pronounced solvatochromism, it is possible to determine the dielectric
constant based on their fluorescence measurements. It appeared that for n-Bu 4 NBF 4
solutions in dichloromethane, its dielectric constant increases substantially from 8.93
for neat solvent to 14.1 for 0.1 M concentration, and to 18.0 for 0.2 M solution [15],
which lowers the oxidation potential of ferrocene accordingly.
Another effect that leads to redox potential shift owing to stabilisation of ions
is ion-pairing. For highly charged species, like [Ru(bpy) 3 ]
2+/3+ , ion-pairing effect
becomes significant attaining almost 500 mV in CH 2 Cl 2 , however, this effect is
drastically reduced to below 50 mV in polar MeCN, as shown by Batista group
[16] by including explicit BF 4
− counterions in modelling. Ion-pairing effect may be
relatively small for molecules like ferrocene in 1,2-dichloroethane containing 0.1 M
hexafluorophosphate or perchlorate, in which case the measured ion-pairing effect is
equal to ca. 30 mV [17], which means indistinguishable from the simple ionic strength
effect described above. However, in the case of ferrocenes with electron-donating or
withdrawing substituents, interacting more specifically with counter ions, this may
bring about the shift of almost 200 mV [17].
It is still a common practice of quoting the redox potentials of processes proceeding in non-aqueous systems as they were measured, usually versus aqueous reference
electrodes, like SCE or Ag/AgCl electrodes. This practice should be discouraged,
as the potential measured in this way contains liquid junction potential occurring at
the boundary between the non-aqueous phase in the cell and the aqueous solution
in the reference electrode. This potential of the magnitude from tens to even a few
hundred mV depends not only on the solvents, but is a function of the types and
concentrations of electrolytes in both phases [18, 19]. For instance, liquid junction
potential between 0.1 M solution of tetraethylammonium picrate in DMF and a saturated aqueous KCl solution in SCE is 172 mV, and when the solvent is acetonitrile, it
is equal to 93 mV [18]. In most cases, the higher potential is at the organic phase side,
making the potentials measured versus SCE higher by these values. Despite some
newly published procedures that enable the experimental determination of the Gibbs
energy of transfer of single ions between solvents [20], it is by far more convenient to
413
for a given class of compounds have been established. This methodology, however,
does not account for a couple of factors that may also affect the redox potential.
One of these factors is the ion energy lowering caused by ionic strength, induced
by base electrolyte (also called supporting electrolyte; it is used in large excess to
ensure diffusion-controlled transport of electroactive species, thus making migration
in the electric field negligible). As the following example shows, this stabilisation
energy is important for non-polar solvents. In these solvents, it comes to the association of the base electrolyte and this should be taken into account, when calculating
the ionic strength. For n-Bu 4 NPF 6 in dichloromethane, the fraction of free ions in
0.1 M solution in CH 2 Cl 2 is equal to only 7.7% [14]. The resulting ionic strength
would cause Fc
+ /Fc potential shift to a value lower by 53 mV, as estimated from
the extended Debye–Hückel law. For more polar solvents, the shift would be much
smaller, even, if assumed that there would not come to any association in these polar
solvents. Another approach is measuring the variation in the dielectric constant (relative permittivity) value effected by the presence of electrolytes. Using dyes that
show known pronounced solvatochromism, it is possible to determine the dielectric
constant based on their fluorescence measurements. It appeared that for n-Bu 4 NBF 4
solutions in dichloromethane, its dielectric constant increases substantially from 8.93
for neat solvent to 14.1 for 0.1 M concentration, and to 18.0 for 0.2 M solution [15],
which lowers the oxidation potential of ferrocene accordingly.
Another effect that leads to redox potential shift owing to stabilisation of ions
is ion-pairing. For highly charged species, like [Ru(bpy) 3 ]
2+/3+ , ion-pairing effect
becomes significant attaining almost 500 mV in CH 2 Cl 2 , however, this effect is
drastically reduced to below 50 mV in polar MeCN, as shown by Batista group
[16] by including explicit BF 4
− counterions in modelling. Ion-pairing effect may be
relatively small for molecules like ferrocene in 1,2-dichloroethane containing 0.1 M
hexafluorophosphate or perchlorate, in which case the measured ion-pairing effect is
equal to ca. 30 mV [17], which means indistinguishable from the simple ionic strength
effect described above. However, in the case of ferrocenes with electron-donating or
withdrawing substituents, interacting more specifically with counter ions, this may
bring about the shift of almost 200 mV [17].
It is still a common practice of quoting the redox potentials of processes proceeding in non-aqueous systems as they were measured, usually versus aqueous reference
electrodes, like SCE or Ag/AgCl electrodes. This practice should be discouraged,
as the potential measured in this way contains liquid junction potential occurring at
the boundary between the non-aqueous phase in the cell and the aqueous solution
in the reference electrode. This potential of the magnitude from tens to even a few
hundred mV depends not only on the solvents, but is a function of the types and
concentrations of electrolytes in both phases [18, 19]. For instance, liquid junction
potential between 0.1 M solution of tetraethylammonium picrate in DMF and a saturated aqueous KCl solution in SCE is 172 mV, and when the solvent is acetonitrile, it
is equal to 93 mV [18]. In most cases, the higher potential is at the organic phase side,
making the potentials measured versus SCE higher by these values. Despite some
newly published procedures that enable the experimental determination of the Gibbs
energy of transfer of single ions between solvents [20], it is by far more convenient to