Hirotaka
Inoue
a,
Hongyao
Zhou
*a,
Hideo
Ando
b,
Sakuya
Nakagawa
b and
Teppei
Yamada
*a
aDepartment of Chemistry, Graduate School of Science, The University of Tokyo, 7-3-1 Hongo, Bunkyo-ku, Tokyo, 113-0033, Japan. E-mail: teppei@chem.s.u-tokyo.ac.jp; hozhou@chem.s.u-tokyo.ac.jp
bFaculty of Science, Yamagata University, 1-4-12 Kojirakawa-machi, Yamagata, 990-8560, Japan
First published on 27th November 2023
A thermocell is an emerging alternative to thermoelectric devices and exhibits a high Seebeck coefficient (Se) due to the large change of solvation entropy associated with redox reactions. Here, the Se of p-chloranil radicals/dianions (CA˙−/2−) in acetonitrile was drastically increased from −1.3 to −2.6 mV K−1 by the addition of ethanol, and the increment surpassed the estimation of the classical Born model with continuum solvent media. UV-vis spectroscopy and electrochemical measurements at various mixing ratios of acetonitrile to ethanol revealed that the strong hydrogen bonding between ethanol and oxygen atoms of CA2− forms a 4:1 solvent–ion pair, while the ethanol molecules binding to CA2− dissociate upon its oxidation to CA˙−. The local solvation structures of CA2− are in good agreement with density functional theory. This order–disorder transition of the local solvation structure around the CA˙−/2− ions produces a large entropy change and results in a large Se value. The tailored solvation structure of redox ions by hydrogen bonding is a versatile method applicable to a variety of redox pairs and solvents, contributing to the development of electrolyte engineering for thermocells.
The improvement of Se is always desirable, especially for the utilization of a small temperature difference. Increasing the Se value of thermocells is a major challenge in improving the output power and thermal efficiency. The Se value is proportional to the entropy change induced by the redox reaction (ΔSrc) and can be expressed with the number of electrons mediating the redox reaction (n) and Faraday constant (F):
(1) |
The value of ΔSrc is induced by the reorganization entropy of the solvation structure around the redox species in the charge-transfer process.2,14 Weaver and Hupp evaluated the ΔSrc of transition metal complexes in various solvents.14 Based on the experimental results and the Born model, the authors found that the ΔSrc values increase with a larger charge density of the redox pair and smaller acceptor numbers of the solvents. Their theory agreed well with the general observations that a ferri/ferrocyanide redox pair (Fe(CN)63−/4−)15,16—with its extremely large charge density—shows a large ΔSrc value (−135 J K−1 mol−1) and a large Se value (−1.4 mV K−1).17
For improving the ΔSrc values, solvent mixing to decrease the acceptor numbers of the electrolyte was studied previously.18,19 Pringle and coworkers showed that an addition of 50 vol% dimethyl sulfoxide (DMSO) in an ionic liquid increased the Se value of the cobalt(II/III) tris(bipyridyl) redox pair from +1.64 to +1.97 mV K−1.19 In contrast, mixing of DMSO in water decreased the Se of Fe(CN)63−/4−, because of the weakened interaction between water and Fe(CN)63−/4−.18 The two opposite effects of mixing DMSO in the ionic liquid or in water indicate that the influence of solvent mixing on the Se value is unpredictable solely by the simple addition rule of the physical properties of the bulk solvents based on the Born model of solvation.
Recently, molecular dynamics (MD) simulation provided detailed information on the local solvation structures of redox-active ions.20,21 Sun and coworkers improved the Se of an aqueous ferrous/ferric (Fe2+/Fe3+) thermocell in a hydrogel by mixing acetone,21 and Yang and coworkers investigated the microscopic solvation structure around Fe2+/Fe3+ ions.20 The authors revealed that the insertion of one acetone molecule into the solvation shell of Fe2+ results in its structural distortion. The authors claimed that the order–disorder transition of the solvation structure upon the redox reaction of Fe3+ to Fe2+ increases the ΔSrc and the Se. Furthermore, the hydration structure of the Fe2+/3+ redox pair can be tailored by changing pH and counter anions to increase the ΔSrc and Se.22,23 Tailored local solvation around the redox molecules or ions is the key to increasing the Se; however, in a water–organic mixed solvent, the strong hydrogen bonding between the solvent water molecules themselves in the organic media often limits the designability of the solvation structure.
In this work, we focused on designing the local solvation structure of an organic redox molecule in an organic–organic solvent mixture. We used a p-chloranil radical anion and dianion (CA˙−/CA2−, Fig. 1a) as the redox pair, and ethanol (EtOH) as the modifier of their local solvation structure in acetonitrile media. Previously, a positive shift of the electrochemical potential of chloranil was observed when methanol was added in an acetonitrile electrolyte.24 In this study, acetonitrile was selected as the solvent medium because of its high dielectric constant, and ethanol was selected as the additive because of its higher boiling point compared to methanol. CA2− creates an ordered solvation structure with EtOH via hydrogen bonding between the negatively charged oxygen atoms of CA2− and the hydroxy group of EtOH (Fig. 1b). When CA2− is oxidized to CA˙−, the solvation structure of EtOH is broken because of the decreasing charge density of the CA˙− species. UV-vis spectroscopy and density functional theory (DFT) calculations confirmed that the hydroxy group of EtOH is strongly attracted to the electronegative oxygen atoms of CA2−. In addition, electrochemical measurements showed that EtOH and CA2− form a 4:1 solvent–ion pair in the acetonitrile media. The order–disorder transition of the EtOH–CA2− pair increased the ΔSrc by −160 J K−1 mol−1 and resulted in the highest Se of −2.6 mV K−1 in the organic–organic mixed solvent system. Our results give a new concept of tailoring the local solvation structure around organic redox molecules, which is potentially applicable to many other combinations of benzoquinone derivatives and hydrogen bonding donors.
The normalized power output, P/(ΔT)2, of the CA˙−/2− cell was measured with various concentrations of EtOH (Fig. 1d). The result was compared with that of an aqueous electrolyte containing the same concentration of Fe(CN)63−/4− (0.5 mM each).25,26 The Se value of Fe(CN)63−/4− was −1.6 mV K−1 (Fig. S3, ESI†) and agreed with the literature value.27 The P/(ΔT)2 of the organic CA˙−/2− thermocell surpassed the value of the aqueous Fe(CN)63−/4− thermocell when 1.3 M EtOH was added to acetonitrile electrolyte because of the increased Se value (see Fig. S4, ESI,† for the I–V curve). One drawback of the organic CA˙−/2− thermocell is, however, a much lower solubility of CA˙−/2− in acetonitrile compared to the solubility of Fe(CN)63−/4− in water. The output power of the CA˙−/2− thermocell could be improved by using quinone derivatives as the redox molecule with a high solubility in organic electrolytes.
Continuous power generation of the CA˙−/2− thermocell with EtOH (1.3 M) was confirmed at a constant voltage load (Fig. 1e). Stable power output for more than three hours was observed without any precipitation. Previously, an addition of 20 wt% methanol (MeOH) in water was reported to increase the Se of the Fe(CN)63−/4− thermocell to −2.9 mV K−1.28 The low solubility of Fe(CN)63−/4− in MeOH caused precipitation,18 which increased the Se values by the concentration gradient of Fe(CN)63−/4− between the hot and cold electrodes,29–31 while this precipitation can prevent the diffusion of the redox species and deteriorate the output power over time.32,33 In this study, no precipitation is caused by the mixing of EtOH, clearly discerning the influence of precipitation and the solvation effect.
λ calc (nm) | λ obs (nm) | Peak assignment | |
---|---|---|---|
a The λcalc values correspond to the calculated excitation energies in stick spectra (see the ESI for further details, including the molecular orbitals in Fig. S7 and S8). | |||
CA˙− | 219 | 214 | 2B3u ← 2B2g (π*(au) ← π(b1g)) |
274 | 321 | 2B3u ← 2B2g (π*(b3u) ← π*(b2g), π*(au) ← π(b1g)) | |
406 | 420 | 2Au ← 2B2g (π*(au) ← π*(b2g)) | |
428 | 448 | 2B3u ← 2B2g (π*(b2g) ← π(b3u)) | |
CA2− | 219, 222 | 223 | 1B1u ← 1Ag (π*(au) ← π(b1g), Rydberg(b3u) ← π*(b2g)) |
253 | 258 | 1B1u ← 1Ag (π*(b3u) ← π*(b2g)) | |
381 | 359 | 1B2u ← 1Ag (π*(au) ← π*(b2g)) |
The UV-vis spectra of the CA species with the addition of EtOH were measured to investigate the effect of EtOH on the electronic states of the CA species. The addition of EtOH caused negligibly small red shifts (by 2 nm at maximum) in the absorption peaks of CA (at 288 nm) and CA˙− (at 448 nm) (Fig. S6a and b, ESI†). In contrast, a significantly large blue shift (by 30 nm at maximum) was observed in the absorption peak of CA2− at 359 nm with increasing concentration of EtOH (Fig. 2a (top), Fig. 2b, and S6c, ESI†). The TD-DFT calculations show a similar blue shift (Fig. 2a, bottom) and indicate that the absorption peak at 359 nm (381 nm in theory) is assignable to the electron excitation from a π*(b2g) orbital to a π*(au) orbital. The π*(b2g) one is the highest occupied molecular orbital (HOMO), and the π*(au) one is the second lowest unoccupied molecular orbital (LUMO+1) (Fig. 2c). Solvent exchange from acetonitrile to EtOH within the polarizable continuum model (PCM) framework resulted in no peak shift of CA2− (Fig. S9b, ESI†). The change in the dielectric constant of the solvent causes a minimum effect on the electronic structures of the CA species, which highlights the importance of the hydrogen bonding around CA2−.
The calculated molecular geometry of the EtOH–CA2− solvent–ion pair shows that, because of a balance between the hydrogen bonding and the steric jamming among the EtOH molecules, four EtOH molecules are expected to bind with one CA2− ion at maximum (Fig. 2c; the detailed explanation presented in the ESI†). The energy gap between π*(b2g) and π*(au) (i.e., the HOMO–LUMO+1 gap), as well as the corresponding excitation energy, increases with increasing number of EtOH molecules binding with CA2− (Fig. 2c). The electropositive proton of the hydroxy group of EtOH approaches the two electronegative oxygen atoms of CA2−, particularly their out-of-plane 2p-type lobes of the π*(b2g) orbital (Fig. 2c), thereby decreasing the π*(b2g) energy level, increasing the HOMO–LUMO+1 gap and causing the blue shift.
The temperature dependence of E1/2 corresponds to the ΔSrc according to eqn (1), and changes in ΔSrc caused by the addition of EtOH can be ascribed to the effect of ΔSas. The SWV measurement at various temperatures shows that the addition of EtOH drastically enhances the negative shift of EII1/2 with increasing temperature, indicating that the coordination of EtOH to CA2− increases the ΔSrc (Fig. 3b and c). The increment of EI1/2 and EII1/2versus the EtOH concentration was fitted by using theoretical models based on the Nernst equation, and the binding constant, ΔHas and ΔSas of EtOH to CA˙−/2− were obtained from the fitting parameters (Fig. S18–S20, ESI,† the derivation of the theoretical models is shown in the ESI†).
The maximum coordination number of EtOH molecules to one CA2− dianion was set to 4:1 based on the calculated molecular geometry (Fig. 2c). To reduce the number of fitting parameters, only even numbers of EtOH molecules (2 or 4) were allowed to associate with one CA2− in this fitting, which is a reasonable approximation, considering the symmetric structure of CA2−. The fitting results of ΔHas and ΔSas are summarized in Table 2, and the selective association of EtOH to CA2− can be attributed to the larger ΔHas with CA2− than with CA˙−. The average number of EtOH molecules associated with one CA˙− or CA2− anion at each EtOH concentration was calculated from the binding constants, and the result shows that two EtOH molecules associate with one CA2− ion at 0.1 M EtOH, and four EtOH molecules above 1 M EtOH on average, which is larger than the association number of EtOH to CA˙− (Fig. 3d). This result is consistent with the DFT calculation—EtOH more strongly binds with CA2− than with CA˙− (Fig. S15 and S16, ESI†).
Number of EtOH | ΔHas (kJ mol−1) | ΔSas (J K−1 mol−1) | |
---|---|---|---|
CA˙− | i = 1 | −9.1 ± 0.2 | −39 ± 1 |
i = 2 | −12 ± 0 | −66 ± 0 | |
CA2− | j = 2 | −34 ± 1 | −59 ± 3 |
j = 4 | −77 ± 0 | −180 ± 1 |
The ΔSrc of the CA˙−/2− redox pair at various EtOH concentrations was evaluated by temperature-variable SWV measurement (Fig. S21, Table S2, ESI†) and plotted together with the theoretical curve (Fig. 3e; the derivation is shown in the ESI†). Fig. 3e can be divided into three regions: Region 1 shows only a minimum increase of ΔSrc, because of little or no association of EtOH with CA2− or CA˙− ions; Region 2 shows the maximum increase of ΔSrc, because the number of EtOH binding with CA2− significantly increases and forms a 4:1 solvent–ion pair at maximum, while the association number with CA˙− remains low; Region 3 shows decreasing ΔSrc values, mainly because the number of EtOH binding with CA˙− increases at such a high EtOH concentration (Fig. S22, ESI†). The above discussions point out that, in the mixed solvent system, the Born model—regarding the solvent as a continuum medium—fails to explain the molecular-level interaction between the redox ions and the solvent molecules. Rather, the molecular dynamics and the intermolecular forces at the microscopic level play a critical role in determining ΔSrc. The slight deviation of the experimental values of ΔSrc from the simulation curve is likely due to the gradual vaporization of EtOH from the electrolyte during the SWV measurement at high temperature (details are discussed in the ESI†).
Supramolecular thermocells developed previously by us exemplify that the molecular-level host–guest interaction can boost the Se value.30,47–49 Briefly, cyclodextrins (CDs) encapsulate the redox-active triiodide anion at the cold electrode and release the guest triiodide at the hot electrode. This host–guest interaction is driven by the hydrophobic interaction, and the increment of Se values induced by the addition of CDs is linearly proportional to the entropy change induced by the host–guest encapsulation.2
The CA-based thermocell developed here was inspired by the selective host–guest interaction used in a supramolecular thermocell system. The formation of the 4:1 EtOH–CA2− pair produces a ΔSas of −180 J K−1 mol−1, surpassing the largest entropy change obtainable from the host–guest encapsulation between the CD and triiodide (−150 J K−1 mol−1). The order–disorder transition of the solvent–ion pair was found to be a more effective strategy for enhancing the ΔSrc and Se.
The equivalent circuit diagram and the molecular pictures of the charge-transfer process of the CA˙−/2− redox pair on the electrode surface are shown in Fig. 4d. The EtOH molecules around CA2− form a strong hydrogen bonding, and the dissociation of EtOH molecules is the rate-determining step in the electron-transfer process from CA2− to the electrode. In contrast, the CA˙− radical anion and neutral CA0 species form weak or no hydrogen bonding with EtOH solvent molecules, and the electron-transfer process is unhindered between CA0 and CA˙−. A similar result was previously reported that the reorganization of solvent water molecules influences the kinetics of the Fe(CN)63−/4− redox reaction.50 The increasing Rct of the CA˙−/2− redox pair with increasing concentration of EtOH further supported our hypothesis that the large ΔSrc is caused by the tight hydrogen bonding between the solvent EtOH and CA2− dianion.
Footnote |
† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d3sc04955h |
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