Control of Fe3+ coordination by excess Cl− in alcohol solutions

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 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 is interpretted in terms of [FeClnL6−n]3−n (n = 1 and 2). In the EtOH system, r steeply increases from 0.1 at [Cl−] = 1.5 mM to 0.7 at [Cl−] = 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 (S EC ) because V is equivalent to ÀDG/e, where DG and e are the variation in the Gibbs free energy associated with reduction reaction and elementary charge (>0). In a technological point of view, the electrochemical parameters can be used for energy harvesting device, such as liquid thermoelectric cell (LTE). [1][2][3][4][5][6] In this sense, it is scientically and technologically important to deeply comprehend and control the coordination state of redox pairs in solution. The Fe ion in solution is usually octahedrally coordinated by six solvent molecules (L) forming the FeL 6 complex. Inada et al. 7 reported that Fe 2+ is coordinated by six L in aqueous, methanol (MeOH), ethanol (EtOH), dimethyl sulfoxide (DMSO) solutions. If the aqueous solution contains Cl À , however, it is reported that Fe 3+ takes various coordination state, 8,9 such as [FeCl n (H 2 O) 6Àn ] nÀ3 and [FeCl 4 ] À , reecting a strong interaction between Fe 3+ and Cl À . The Fe coordination in aqueous solution containing Cl À is still controversial.
Recently, there has been an interest in the electrochemistry of redox pairs not only in aqueous solutions but also in organic solutions. Especially, Inoue et al. 10 reported that S EC of Fe 2+ /Fe 3+ in several organic solvents are much higher than S EC in aqueous solution. For example, S EC (¼ 3.6 mV K À10 ) of Fe 2+ /Fe 3+ in acetone is much larger than that (¼ 1. 5  . Then, systematic investigation against Cl À concentration ([Cl À ]) at a xed Fe 3+ concentration is effective for comprehension and control of the complex formation. We emphasized that the ultraviolet-visible (UV-vis) absorption spectroscopy is a sensitive probe for complex formation, because Fe 3+ complex exhibits characteristic absorption bands in this region. 12 In addition, the spectroscopy is sensitive even in a dilute Fe 3+ solution of sub mM and is suitable for investigation of organic solution. In

UV-vis absorption
The UV-vis absorption spectra of Fe 3+ 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 (I 0 ) without cell. The molar absorption coefficient (3) was dened by Àln(I/I 0 )/cd, where c (¼ 0.5 mM) and d (¼ 1 cm) are the molar concentration of Fe 3+ and thickness of the optical cell, respectively.
The Variation of redox potential V

We systematically investigated variation in
where V sample and V ref 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. NaClO 4 (10 g per 100 mL) and agar (4 g per 100 mL) were added to water. Then, the solution was heated, dissolved, poured into a U-shaped tube, and cooled to harden. A Pt electrode was inserted into each cell.

Results and discussion
Overall feature of spectra indicates that the Fe 3+ complex is stable even with excess ClO 4 À . In MeOH and EtOH solutions, the spectra exhibit two absorption bands at 360 nm and 260 nm. We ascribed the spectral feature to formation of [FeL 6 ] 3+ . We calculated the absorption spectra of [Fe(MeOH) 6 ] 3+ cluster with Gaussian 16W program 13 (Fig. S1 †). The calculated spectrum shows two-band structure at 190 nm and 370 nm due to the ligand to metal charge transfer (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)(H 2 O) 5 ] 2+ . 14,15 Fig. 1(b) shows the 3 spectra of FeCl 3 dissolved in H 2 O, MeOH, and EtOH. Thick curves represent for the spectra without excess Cl À . The spectrum of the MeOH and EtOH solutions without excess Cl À still exhibit two-band structure at 250 nm and 370 nm. The spectral weight of the 260 nm band is higher than that of the Fe(ClO 4 ) 3 solutions [(a)]. We ascribed the spectral feature to formation of [FeCl n L 6Àn ] 3Àn (n ¼ 1 and 2), because the oscillator strength (f) of an electron transfer from Cl to Fe 3+ is larger than that from O to Fe 3+ . We calculated the absorption spectra of [  14,15 Surprisingly, addition of excess Cl À completely changes the spectra. Thin curves in Fig. 1(b) represent for the spectra with excess Cl À . In the MeOH and EtOH solutions, excess Cl À 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 Fe 3+ is not coordinated by L, but Cl À . In Fig. 2 3 Concerning to the lower-lying two absorption bands, the spectral proles are essentially the same. This clearly

Spectral change against [Cl À ]
Now, let us investigate in detail how the 3 spectrum changes with increase in [Cl À ]. Fig. 3(a) shows the 3 spectra of the MeOH solution containing 0.5 mM Fe 3+ against [Cl À ]. At [Cl À ] ¼ 0.0 mM, the spectrum exhibits two broad absorption bands at 360 nm and 260 nm, which are due to an electron transfer from L to Fe 3+ within the FeL 6 complex. At [Cl À ] ¼ 1.5 mM, the spectra still show two-band structure, but the spectral weight of the higher energy band is much higher than those at [Cl À ] ¼ 0.0 mM. As discussed in the previous subsection, the spectral change is interpreted in terms of the formation of [FeCl n -L 6Àn ] 3Àn (n ¼ 1 and 2). At [Cl À ] 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 À ] increases. At [Cl À ] ¼ 71.5 mM, the spectrum shows characteristic three band structure due to FeCl 4 À (Fig. 2) Fig. 3(b)]. At [Cl À ] ¼ 0.0 and 0.9 mM, the spectra exhibit two-band structure. At [Cl À ] ¼ 2.5 mM, an additional absorption band appears as indicated by an open triangle. Its spectral weight steeply increases as [Cl À ] increases. At [Cl À ] ¼ 4.5 mM, the spectrum shows characteristic three band structure due to FeCl 4 (Fig. 2). The spectra remain unchanged in the [Cl À ] region above 4.5 mM, indicating that all Fe 3+ forms the FeCl 4 complex. We found three isosbestic points at 225 nm, 260 nm, and 385 nm above [Cl À ] ¼ 0.

FeCl 4 formation against [Cl À ]
The solution system investigated contains multiple complex species, such as, [FeL 6 ] 3+ , [FeCl n L 6Àn ] 3Àn (n ¼ 1 and 2), [FeCl 4 ] À . It is difficult to unambiguously decompose the spectrum into the respective components because there is no quantitative information on the spectra due to [FeCl n L 6Àn ] 3Àn . Fortunately, we know the spectrum of [Fe III Cl 4 ] À . In addition, the FeCl 4 band at 318 nm is well separated from the absorption bands due to other complexes. In the following, we focused our attention on the FeCl 4 formation against [Cl À ]. We evaluated the intensities and peak positions (l p ) of the FeCl 4 band by least-squares tting with three Gaussian functions ( Fig. S3 and S4 †). The intensities were normalized by the value at [Cl À ] ¼ 101.5 mM (7.5 mM) for the MeOH (EtOH) solutions, where all Fe 3+ is considered to form the FeCl 4 complex. Then, the proportion (r) of the FeCl 4 complex is the same value as the normalized intensity (I).   The difference in the complex formation between the MeOH and EtOH solutions is probably reects the difference in solubility of Cl À . The solubility (¼ 6 mM) of NaCl in EtOH is much smaller than that (¼ 100 mM) in MeOH.  Fig. 1(b)

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