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Control of Fe3+ coordination by excess Cl inalcohol solutions†
Yunika Nomura,a Dai Inoueb and Yutaka Moritomo*abcd

We spectroscopically investigated coordination state of Fe3+ in methanol (MeOH) and ethanol (EtOH)
solutions against Cl concentration ([Cl]). In both the system, we observed characteristic absorption
Received 8th March 2022
Accepted 10th June 2022
DOI: 10.1039/d2ra01522f
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bands due to the FeCl4 complex at high-[Cl] region. In the MeOH system, the proportion (r) of [FeCl4]
exhibits a stationary value of 0.2–0.3 in the intermediate region of 10 mM < [Cl] < 50 mM, which isinterpretted in terms of [FeClnL6n]3n (n ¼ 1 and 2). In the EtOH system, r steeply increases from 0.1 at[Cl] ¼ 1.5 mM to 0.7 at [Cl] ¼ 3.5 mM, indicating direct transformation from [FeL6]3+ to [FeCl4]. We
further found that the coordination change signiﬁcantly decreases the redox potential of Fe2+/Fe3+.

Introduction
The coordination state around the redox pair in solution has
a great inuence on the redox potential (V) as well as its temperature coeﬃcient (SEC) because V is equivalent to DG/e, where DG
and e are the variation in the Gibbs free energy associated with
reduction reaction and elementary charge (>0). In a technologicalpoint of view, the electrochemical parameters can be used forenergy harvesting device, such as liquid thermoelectric cell(LTE).1–6 In this sense, it is scientically and technologicallyimportant to deeply comprehend and control the coordination
state of redox pairs in solution. The Fe ion in solution is usuallyoctahedrally coordinated by six solvent molecules (L) forming the
FeL6 complex. Inada et al.7 reported that Fe2+ is coordinated by
six L in aqueous, methanol (MeOH), ethanol (EtOH), dimethyl
sulfoxide (DMSO) solutions. If the aqueous solution contains Cl,
however, it is reported that Fe3+ takes various coordination state,8,9such as [FeCln(H2O)6n]n3 and [FeCl4], reecting a strong
interaction between Fe3+ and Cl. The Fe coordination in aqueous
solution containing Cl is still controversial.
Recently, there has been an interest in the electrochemistry ofredox pairs not only in aqueous solutions but also in organic
solutions. Especially, Inoue et al.10 reported that SEC of Fe2+/Fe3+
in several organic solvents are much higher than SEC in aqueous
solution. For example, SEC (¼ 3.6 mV K10) of Fe2+/Fe3+ in acetone
is much larger than that (¼ 1.5 mV K6) in aqueous solution. InaSchool of Science and Engineering, University of Tsukuba, Tsukuba 305-8571, Japan.
E-mail: moritomo.yutaka.gf@u.tsukuba.ac.jpbGraduate School of Pure and Applied Sciences, University of Tsukuba, Tsukuba 3058571, Japanc

Faculty of Pure and Applied Sciences, University of Tsukuba, Tsukuba 305-8571, Japand

Tsukuba Research Center for Energy Materials Science (TREMS), University ofTsukuba, Tsukuba 305-8571, Japan† Electronic
supplementary
information
https://doi.org/10.1039/d2ra01522f

17932 | RSC Adv., 2022, 12, 17932–17936(ESI)

available.

Seeaddition, Wake et al.11 reported that the dimensionless performance index (ZT) of LTE is much enhanced if aqueous solution of
Fe2+/Fe3+ is replaced with acetone solution of Fe2+/Fe3+. In thissituation, a comprehension of the coordination state of Fe ions in
the organic solution is desired. In general, the formation of FeCl4
in solution is governed by the equilibrium constant (K); K ¼[FeCl4]/[Fe3+][Cl]4 for Fe3+ + 4Cl 4 FeCl4. Then, systematicinvestigation against Cl concentration ([Cl]) at a xed Fe3+
concentration is eﬀective for comprehension and control of thecomplex formation. We emphasized that the ultraviolet-visible(UV-vis) absorption spectroscopy is a sensitive probe forcomplex formation, because Fe3+ complex exhibits characteristic
absorption bands in this region.12 In addition, the spectroscopy issensitive even in a dilute Fe3+ solution of sub mM and is suitablefor investigation of organic solution.
In this work, we spectroscopically investigated coordination
state of Fe3+ against [Cl] in the MeOH and EtOH solutions. In
MeOH solution, with increases in [Cl], octahedral [FeL6]3+
gradually transforms to tetrahedral [FeCl4] via mixed [FeClnL6n]3n state. In EtOH solution, [FeL6]3+ directly transforms to[FeCl4] with increase in [Cl]. The variation of the Fe coordination signicantly decreases V of Fe2+/Fe3+.

ExperimentUV-vis absorption
The UV-vis absorption spectra of Fe3+ solutions were investigated with a spectrometer (V750, Jasco) at room temperature.
Absorption spectra were obtained by dividing the transmission
intensity spectra (I) of the solution by that (I0) without cell. The
molar absorption coeﬃcient (3) was dened by ln(I/I0)/cd,
where c (¼ 0.5 mM) and d (¼ 1 cm) are the molar concentrationof Fe3+ and thickness of the optical cell, respectively.
The solvents investigated were water (H2O), MeOH (FUJIFILM Wako corp.), EtOH (FUJIFILM Wako corp.). The solvents© 2022 The Author(s). Published by the Royal Society of ChemistryView Article Online

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are purchased and used as received. The solutions contain
0.5 mM Fe3+. The [Cl] value was controlled by dissolving
FeCl3$6H2O (FUJIFILM Wako corp.), Fe(ClO4)3$6.9H2O (FUJIFILM Wako corp.), and NaCl (FUJIFILM Wako corp.) in appropriate molar ratios. At several [Cl] in MeOH and EtOH
solutions, we conrmed that the crystal water in solution does
not aﬀect the spectra. The intensity and shape of the spectra ofanhydrous FeCl3 (FUJIFILM Wako corp.) solutions are essentially the same as that of the corresponding hydrous FeCl3
solution. In aqueous solution, LiCl (FUJIFILM Wako corp.) wasused instead of NaCl, because LiCl shows higher solubility inwater. We further investigated the solutions containing 0.5 mM
Fe(ClO4)3$6.9H2O (FUJIFILM Wako corp.). The ClO4 concentration ([ClO4]) was controlled by dissolving extra NaClO4(FUJIFILM Wako corp.).
Variation of redox potential V
We systematically investigated variation in V of Fe2+/Fe3+
against [Cl]. We dene DV as Vsample  Vref, where Vsample andVref are V of the sample and reference cells, respectively. The
reference and sample cells were beaker cells which were connected by a salt bridge. The salt bridge was made as follows.
NaClO4 (10 g per 100 mL) and agar (4 g per 100 mL) were added
to water. Then, the solution was heated, dissolved, poured intoa U-shaped tube, and cooled to harden. A Pt electrode wasinserted into each cell.
The electrolyte in the reference cell was MeOH (or EtOH)
solution containing 0.5 mM Fe(ClO4)2$6H2O (FUJIFILM Wako
corp.) and 0.5 mM Fe(ClO4)3$6.9H2O. The electrolytes in thesample cell were [Cl] controlled MeOH (or EtOH) solution
containing 0.5 mM Fe2+ and 0.5 mM Fe3+. The [Cl] value wascontrolled by dissolving FeCl2$4H2O, FeCl3$6H2O, Fe(ClO4)2$6H2O, Fe(ClO4)3$6.9H2O, and NaCl in appropriate molar ratios.
DV between the cells were carefully investigated with conrming
stabilization of the voltage. At several [Cl] in MeOH and EtOH
solutions, we conrmed that the crystal water in solution does
not aﬀect the redox potential. The redox potentials of anhydrous
FeCl2 (FUJIFILM Wako corp.)/FeCl3 (nacalai tesque) solutions are
the same as that of the corresponding hydrous FeCl2/FeCl3
solutions within the experimental error (<6 mV).

Results and discussion
Overall feature of spectra
Fig. 1(a) shows molar absorption coeﬃcient (3) spectra ofFe(ClO4)3 dissolved in H2O, MeOH, and EtOH. Thick and thincurves represent for the spectra without and with excess ClO4,
respectively. In all solutions, the spectra with excess ClO4 are
the same as those without excess ClO4. This observation
indicates that the Fe3+ complex is stable even with excess ClO4.
In MeOH and EtOH solutions, the spectra exhibit two absorption bands at 360 nm and 260 nm. We ascribed the spectralfeature to formation of [FeL6]3+. We calculated the absorption
spectra of [Fe(MeOH)6]3+ cluster with Gaussian 16W program13(Fig. S1†). The calculated spectrum shows two-band structure at
190 nm and 370 nm due to the ligand to metal charge transfer© 2022 The Author(s). Published by the Royal Society of ChemistryFig. 1 Molar absorption coeﬃcient (3) spectra of (a) Fe(ClO4)3 and (b)FeCl3 dissolved in H2O, methanol (MeOH) and ethanol (EtOH). Thincurves (a) and (b) are the spectra with excess ClO4 and Cl, respectively. Filled triangles in (b) represent the absorption bands due to[FeCl4].

(LMCT) transition and qualitatively reproduces the observed
spectra [Fig. 1(a)]. In aqueous solution, traces of absorption
bands are discernible at 300 nm, which is ascribed to[Fe(OH)(H2O)5]2+.14,15
Fig. 1(b) shows the 3 spectra of FeCl3 dissolved in H2O,
MeOH, and EtOH. Thick curves represent for the spectra
without excess Cl. The spectrum of the MeOH and EtOH
solutions without excess Cl still exhibit two-band structure at
250 nm and 370 nm. The spectral weight of the 260 nm band ishigher than that of the Fe(ClO4)3 solutions [(a)]. We ascribed thespectral feature to formation of [FeClnL6n]3n (n ¼ 1 and 2),
because the oscillator strength (f) of an electron transfer from Cl
to Fe3+ is larger than that from O to Fe3+. We calculated theabsorption spectra of [FeClMeOH5]2+, trans [FeCl2MeOH4]+, andcis [FeCl2MeOH4]+ clusters with Gaussian 16W program13(Fig. S1†). The total oscillator strength (ftot) of transitions
forming the higher energy band is larger in the Cl-substitutedclusters than ftot (¼ 0.42) in [FeMeOH6]3+; ftot is 0.56 in[FeClMeOH5]2+, 0.73 in trans [FeCl2MeOH4]+, and 0.74 in cis[FeCl2MeOH4]+.The spectral shape of the FeCl3 aqueous solution without excess Cl is the same as that of the Fe(ClO4)3
aqueous solution, indicating that the Fe coordination stateremains [Fe(OH)(H2O)5]2+.14,15
Surprisingly, addition of excess Cl completely changes the
spectra. Thin curves in Fig. 1(b) represent for the spectra withexcess Cl. In the MeOH and EtOH solutions, excess Cl causes
sharp absorption bands at 362, 318, and 242 nm, as indicated bylled triangles. We emphasize that the spectral prole of the
MeOH solution is the same as that of the EtOH solution. The
sameness of the two spectra indicates that Fe3+ is not coordinated by L, but Cl. In Fig. 2, we compared the spectra of the
MeOH solution containing 0.5 mM Fe3+ with excess at [Cl] ¼
101.5 mM with that of nitromethane solution of [N(C2H5)4][FeCl4].3 Concerning to the lower-lying two absorption bands,
the spectral proles are essentially the same. This clearly

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Fig. 2 Molar absorption coeﬃcient (3) spectra of MeOH solution containing 0.5 mM Fe3+ at [Cl] ¼ 101.5 mM. Red symbols show the 3
spectra of nitromethane solution of [N(C2H5)4][FeCl4] (cited from ref. 8).

indicates formation of [FeCl4] in the MeOH and EtOH solutions with excess Cl. On the other hands, the spectrum of
aqueous solution at [Cl]/[Fe3+] ¼ 30 000 exhibits two-bandstructure. The spectral prole is consistent with the spectra at[Cl] ¼ 15 M reported by Liu et al.,12 who ascribed the spectralfeature to [FeCl3(H2O)3] and/or [FeCl4]. We emphasize that the
260 nm band is located at higher wavelength side than the
242 nm band due to [FeCl4], as indicated by broken straightlines. This clearly indicates that [FeCl4] is not the dominant
species even at the highest [Cl].
Spectral change against [Cl]Now, let us investigate in detail how the 3 spectrum changes
with increase in [Cl]. Fig. 3(a) shows the 3 spectra of the MeOH
solution containing 0.5 mM Fe3+ against [Cl]. At [Cl] ¼
0.0 mM, the spectrum exhibits two broad absorption bands at
360 nm and 260 nm, which are due to an electron transferfrom L to Fe3+ within the FeL6 complex. At [Cl] ¼ 1.5 mM, thespectra still show two-band structure, but the spectral weight of
the higher energy band is much higher than those at [Cl] ¼
0.0 mM. As discussed in the previous subsection, the spectralchange is interpreted in terms of the formation of [FeClnL6n]3n (n ¼ 1 and 2). At [Cl] ¼ 11.5 mM, trace of an additional
absorption band is discernible, as indicated by an open
triangle. Its spectral weight gradually increases as [Cl]
increases. At [Cl] ¼ 71.5 mM, the spectrum shows characteristic three band structure due to FeCl4 (Fig. 2). The spectra
remain unchanged in the [Cl] region above 71.5 mM. Thismeans that all Fe3+ form the [FeCl4] complex, because theequilibrium does not move with increase in [Cl]. Hereaer, we
will call the absorption band at 318 nm as “FeCl4 band”. A
detailed investigation revealed three isosbestic points at
225 nm, 255 nm, and 385 nm above [Cl] ¼ 11.5 mM[Fig. S2(a)†]. The isosbestic point reects the reaction between
octahedral ([FeClnL6n]3n) to tetrahedral ([FeCl4]) complexes
with increase in [Cl].
A similar [Cl]-dependent spectral change is observed in the
EtOH solution [Fig. 3(b)]. At [Cl] ¼ 0.0 and 0.9 mM, the spectra
exhibit two-band structure. At [Cl] ¼ 2.5 mM, an additional
absorption band appears as indicated by an open triangle. Its

17934 | RSC Adv., 2022, 12, 17932–17936

Fig. 3 Molar absorption coeﬃcient (3) spectra of (a) MeOH and (b)EtOH solution containing 0.5 mM Fe3+ against Cl concentration([Cl]). Open triangles indicate FeCl4 band (see text).

spectral weight steeply increases as [Cl] increases. At [Cl] ¼
4.5 mM, the spectrum shows characteristic three band structure
due to FeCl4 (Fig. 2). The spectra remain unchanged in the [Cl]
region above 4.5 mM, indicating that all Fe3+ forms the FeCl4
complex. We found three isosbestic points at 225 nm, 260 nm,
and 385 nm above [Cl] ¼ 0.9 mM [Fig. S2(b)†]. The isosbestic
point reects the reaction between octahedral ([FeClnL6n]3n)
to tetrahedral ([FeCl4]) complexes with increase in [Cl].
FeCl4 formation against [Cl]
The solution system investigated contains multiple complex
species, such as, [FeL6]3+, [FeClnL6n]3n (n ¼ 1 and 2), [FeCl4].
It is diﬃcult to unambiguously decompose the spectrum into
the respective components because there is no quantitative
information on the spectra due to [FeClnL6n]3n. Fortunately,
we know the spectrum of [FeIIICl4]. In addition, the FeCl4 bandat 318 nm is well separated from the absorption bands due toother complexes. In the following, we focused our attention on
the FeCl4 formation against [Cl]. We evaluated the intensities
and peak positions (lp) of the FeCl4 band by least-squares ttingwith three Gaussian functions (Fig. S3 and S4†). The intensities
were normalized by the value at [Cl] ¼ 101.5 mM (7.5 mM) for
the MeOH (EtOH) solutions, where all Fe3+ is considered toform the FeCl4 complex. Then, the proportion (r) of the FeCl4
complex is the same value as the normalized intensity (I).
Fig. 4 shows I (upper panel) and lp (middle panel) against[Cl] in (a) MeOH and (b) EtOH solutions. For convenience of
explanation, we dene regions I, II, and III as regions where r (¼I) < 0.2, 0.2 < r < 0.7, and r > 0.7, respectively. The region I and III
are dominated by the [FeL6]3+ and [FeCl4] complexes,© 2022 The Author(s). Published by the Royal Society of ChemistryView Article Online

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Fig. 4 (a) Normalized intensity (I; upper panel) and wavelength (lp;
middle panel) at the peak of the FeCl4 band, and variation (DV; bottom
panel) in V of Fe2+/Fe3+ in MeOH solution against [Cl]. The intensity
was normalized by the value at [Cl] ¼ 101.5 mM. (b) I (upper panel) andlp (middle penal), and DV (bottom panel) in EtOH solution against [Cl].
The intensity was normalized by the value at [Cl] ¼ 7.5 mM. For
convenience of explanation, regions I, II, and III are deﬁned (see text).

RSC Advances(or EtOH) solution containing 0.5 mM Fe2+ and 0.5 mM Fe3+ at[Cl] ¼ 0.0 mM, where the FeL6 complex is dominated. We must
confess that we have no detailed information on the complex
state of Fe2+ against [Cl]. Like Fe3+, the complex state isconsidered to change from [FeL6]2+, [FeClnL6n]2n, to [FeCl4]2
with increase in [Cl]. In the following argument, we assume
that the redox reaction of the complex occurs without substitution of the coordinated molecules/ions.
In (a) MeOH system, DV is 0.16 V at the entrance of region
II. The variation in V between [FeL6]2+/[FeL6]3+ (region I) and[FeClnL6n]2n/[FeClnL6n]3n (region II) is well explained interms of the crystal eld splitting. We note that six d electrons in
the octahedral [FeL6]2+ and [FeClnL6n]2n takes high-spin (HS)conguration, similarly to the case of [Fe(H2O)6]2+.16,17 The
crystal eld from Cl is weaker than that from oxygen in L.
Then, V of [FeCl4]2/[FeCl4] is expected to be lower, reecting
narrower crystal eld splitting. In (b) EtOH system, DV is 0.24
V at the entrance of region III. Variation of V between [FeL6]2+/[FeL6]3+ (region I) and [FeCl4]2/[FeCl4] (region III) is also
explained in terms of the crystal led splitting. We note that six
d electrons take the HS conguration in tetrahedral [FeCl4]2,because the crystal eld splitting (¼ 10 Dq) in octahedral
complex is larger than that (¼ 4.45 Dq) in the tetrahedral
complex.18 V of [FeCl4]2/[FeCl4] is expected to be lower,reecting narrower ligand eld splitting.

Conclusions
respectively. In (a) MeOH system, r exhibits a stationary value of
0.2–0.3 in the region of 10 mM < [Cl] < 50 mM (region II). The
fact that r remains 0.2–0.3 indicates existence of a dominant
species other than [FeCl4] and [FeL6]3+, that is, [FeClnL6n]3n(n ¼ 1 and 2). In (b) EtOH, r steeply increases from 0.5 at [Cl] ¼
1 mM to 0.7 at [Cl] ¼ 2.5 mM. The steep increase in r indicatesthat FeL6 almost directly transfers to FeCl4 with increases in[Cl].
The diﬀerence in the complex formation between the MeOH
and EtOH solutions is probably reects the diﬀerence in solubility of Cl. The solubility (¼ 6 mM) of NaCl in EtOH is muchsmaller than that (¼ 100 mM) in MeOH. In EtOH system, Cl insolution is unstable and tend to form the [FeCl4] complex if
Fe3+ exits in solvent. The [Cl] value (¼ 2.0 mM) at the boundary
to region III is nearly the same value (¼ 2.0 mM) as required for
all Fe3+ to become FeCl4. In MeOH system, Cl in solution isstabler than in EtOH. This causes the stable [FeL6]3+ in small[Cl] region (<10 mM; region I) and [FeClnL6n]3n formation in
the intermediate [Cl] region (10 mM < [Cl] < 50 mM; region
II). In aqueous solution where Cl is much more stabler than inalcohol solutions, no trace of [FeCl4] is observed even at [Cl]¼ 15 M [Fig. 1(b)]. Thus, our investigation revealed the solvent
dependence in complex formation against [Cl].
Coordination eﬀect on V
Now, let us investigate the correlation between the Fe ion
coordination and V of Fe2+/Fe3+. The bottom panel of Fig. 4
shows DV (¼ Vsample  Vref) of Fe2+/Fe3+ in (a) MeOH and (b)EtOH solutions. The electrolyte in the reference cell was MeOH

© 2022 The Author(s). Published by the Royal Society of Chemistry

We spectroscopically investigated variation of the Fe ion coordination against [Cl] in MeOH and EtOH solutions. In MeOH
solution, with increases in [Cl], octahedral [FeL6]3+ graduallytransforms to tetrahedral [FeCl4] via mixed [FeClnL6n]3n (n ¼1 and 2). In EtOH solution, [FeL6]3+ directly transforms to[FeCl4] with increase in [Cl]. We found that the coordination
change signicantly decreases V of Fe2+/Fe3+ and interpreted theobservation in terms of the ligand eld splitting.

Author contributions
Y. N. performed spectroscopic and electrochemical experiment.
D. I. have supported the experiment and analysis. D. I further
performed quantum chemistry calculation. Y. M. made theexperimental plan and wrote the manuscript.

Conﬂicts of interest
There are no conicts to declare.

Acknowledgements
This work was supported by Murata Science Foundation, andjoint research with Taisei Rotec Corporation.

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