Iodine Mediator with Anomalously High Redox Potential and Its

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C: Energy Conversion and Storage; Energy and Charge Transport

Iodine Mediator with Anomalously High Redox Potential and Its Application to a Catalytic Cycle for Lithium Carbonate Decomposition toward Future Lithium Reproduction Tohru Shiga, Hiroki Kondo, Yuichi Kato, and Yoko Hase J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b09967 • Publication Date (Web): 28 Jan 2019 Downloaded from http://pubs.acs.org on February 3, 2019

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The Journal of Physical Chemistry

Iodine Mediator with Anomalously High Redox Potential and Its Application to a Catalytic Cycle for Lithium Carbonate Decomposition toward Future Lithium Reproduction Tohru Shiga*, Hiroki Kondo, Yuichi Kato, and Yoko Hase Toyota Central Research & Development Laboratories Inc. Yokomichi, Nagakute-city, Aichi-ken, 480-1192 Japan

ABSTRACT. Lithium carbonate (Li2CO3) is observed in conventional lithium-ion batteries in which cell performance has been reduced. In this paper, a catalyst to electrochemically decompose Li2CO3 was investigated to develop a Li reproduction technology. It has been accepted that iodine with a redox potential below 3.5 V vs. Li+/Li cannot decompose Li2CO3, because the decomposition potential of Li2CO3 is 3.82 V vs. Li+/Li. Here, we observed that the redox potential of iodine in trimethyl phosphate (TMP) and dimethyl sulfoxide (DMSO) solutions increased up to 4.0 V, and thus Li2CO3 could be chemically decomposed by iodine in these solvents. This unexpectedly high redox potential may be caused by deviation of the electric charges on two iodine atoms due to a 1:1 complex formation of iodine with the solvent. Reaction pathways are discussed with respect to the chemical decomposition of Li2CO3 by iodine. We combined the chemical decomposition behavior with the redox couple of iodine species. A double-chamber type cell with

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carbon paper as cathode in catholyte that dissolved iodine and suspended Li2CO3 powders was constructed to substantiate a catalytic cycle for Li2CO3 decomposition. The cell was demonstrated to be rechargeable due to the decomposition of Li2CO3.

1. Introduction Lithium-ion batteries (LiBs) are among the most important energy storage devices for future vehicle applications. Recent declarations for the production of electric vehicles will dramatically increase the demand for LiBs. Lithium (Li) is a relatively rare element with a Clarke number of 0.006, although it is found in many rocks and some brines. The rapid lithium demand for battery applications will lead to a scarcity in lithium resources in the near future with concerns for the environmental destruction caused by lithium mining. Recovering Li from seawater is an attractive alternative, because seawater contains about 230 billion tons of Li.1,

2

However, large-scale

equipment for recovering Li from dilute solution (0.17 mgL-1) would be prohibitively expensive. One strategy to overcome this problem is Li reproduction technology. Lithium compounds such as lithium carbonate (Li2CO3), lithium fluoride (LiF), and alkyl lithium are observed in LiBs in which performance has been reduced.3-5 These compounds are generated from side reactions between Li ions and electrolytes. The reproduction from these deactivated lithium compounds will open a door toward Li reproduction. In this paper, we propose a catalyst that employs a redox mediator to decrease the charge voltage for the decomposition of Li2CO3. The most well-known redox mediator is a soluble redox catalyst such as benzoquinone and 2,2,6,6-tetramethylpiperidine 1-oxyl (TEMPO).6-8 We have examined the catalytic capability of soluble TEMPO molecules9 and acrylate polymers with TEMPO units as a pendant group10,11 as redox mediators to decompose peroxides and metallic


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oxides. Zhang et al demonstrated that Li2CO3 was decomposed by cobalt phthalocyanine as a mobile catalyst at 4.31 V.12 Halogens such as bromine and chlorine are also used as soluble redox mediators. Liang and Lu proposed the use of a bromine redox mediator to realize rechargeable batteries.13 Whereas the redox potential of bromine (2Br3- → 3Br2 + 2e-) is 4.02 V vs. Li+/Li, the decomposition potential of Li2CO3 calculated from the Nernst equation based on the reaction Li2CO3 → 2Li+ + 2e- + CO2 + 0.5O2 is 3.82 V vs. Li+/Li.14 Thus, Br2 can oxidize Li2CO3; 3Br2 + Li2CO3 → 2Br3- + 2Li+ + CO2 + 0.5O2. In the same manner, the oxidation of Li2CO3 by chlorine with a redox potential of 4.39 V vs. Li+/Li is observed (Cl2 + Li2CO3 → 2LiCl + CO2 + 0.5O2). However, bromine and chlorine are highly toxic chemicals and must be handled with care. On the other hand, iodine is a solid at room temperature and less reactive compared to Br2 and Cl2. The redox couple of iodine (I2 →2I-) works near 3.5 V vs. Li+/Li, which is lower than the decomposition potential of Li2CO3. Therefore, it has been considered that I2 cannot oxidize Li2CO3. We have previously investigated the interactions between magnesium oxide (MgO) powders and iodine molecules.15 The MgO powders were decomposed in some iodine solutions, such as in dimethyl sulfoxide (DMSO). The results suggested that iodine has potential in a catalytic cycle as a typical redox mediator. Iodine molecules are complexed with the solvent in a molar ratio of 1:1 (1:1 complex formation)16-19 so that the oxidizing power of iodine may change depending on the extent of the complexion of iodine and solvent. Here, we investigate the decomposition of Li2CO3 by some iodine-solvent complexes and construct a catalytic cycle using this soluble iodine redox mediator to electrochemically decompose Li2CO3 at room temperatures. 2. Experimental



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Materials: Trimethyl phosphate (TMP), propylene carbonate (PC), and dimethyl sulfoxide (DMSO)

were

available

from

Kishida

Chemicals

(battery

grade).

Lithium

bis(trifluoromethanesulfonyl)amide (LiTFSA, Kanto Chemicals) was used as a supporting salt. The salt was dried under vacuum at 150 °C for 6 hours before being mixed with the solvents. Iodine (purity 99.99%), iodine monobromide (IBr) and iodine monochloride (ICl) were obtained from Aldrich. Li2CO3 powder (purity 99.99%) was available from Kojundo Chemical Laboratory Co., Ltd. Measurement of redox potential: Electrochemical cells were fabricated using a modified beaker cell with a carbon cathode, a Li anode, and an electrolyte solution to measure redox potentials of iodine species in organic solvents. The electrolyte was 0.25 mol/L LiTFSA TMP solution that dissolved 10 mg of iodine per 15 mL of electrolyte. The carbon cathode was prepared by dry mixing of carbon black (Tokai Carbon, TB#5500, 95.0% by weight) with Teflon powder (Daikin, F-104, 5.0%) as a binder. The carbon cathode (weight, 10 mg; area, 0.49 cm2; thickness, 0.12 mm) was compressed on a 80-mesh Pt grid (Nilaco corporation). The discharging and charging were started to measure the redox potentials of iodine species with a Hokuto Denko potentiometer (HJ1001SM8A) using a current density of 0.04 mA/cm2 at 25 °C. The redox potentials of IBr and ICl were also measured in the similar way. Chemical decomposition test of Li2CO3 by iodine: In a 9 mL glass bottle, 35 mg of Li2CO3 powder was placed under argon, and then 4 mL of 2 mM iodine TMP solution was poured into the glass. The samples were stored at 60 °C for 72 hours. The tests were also carried out at 25 °C and 80 °C. After the test, the filtrate was identified by several spectroscopic methods. For reference, the decomposition test was also performed in cells using PC and DMSO solutions that dissolved iodine. The Li2CO3 decomposition test by iodine for gas chromatography-mass spectrometry


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(GC/MS) was performed as follows. 75 mg of Li2CO3 powder and 0.4 mL of I2-TMP solution (molar ratio; I2 : TMP = 1:4) were poured into a 5 mL multiple-dose container (in the Supporting Information, Figure S1). The container was stored at 60 °C for 72 hours. After the heating test, gases in the upper side of the container were analyzed. It was performed on a GC/MS spectrometer (Agilent Technologies, GSLG9810) to investigate the gas composition in the multi-dose containers after the decomposition test. The rate of argon flow was 10 mL/sec. Setup of catalytic cycle: A double-chamber type cell was setup to examine the presence or absence of repeatable decomposition of Li2CO3 by iodine (Figure S2). The cathode and the anode were carbon paper (Torey, TGP-H-060) and a lithium metal, respectively. The both chambers were separated by a Li-ion conductive glass-ceramic (LICGC, Li1+x+yAlx(TiGe)2-xSiyP3-yO12, OHARA Inc, thickness = 0.15 mm) to prevent shuttle behavior of iodine.20 The catholyte was an electrolyte of 1 mol/L LiTFSA-TMP solution that dissolved iodine. Li2CO3 powders were suspended in the catholyte to isolate chemical reaction from electrochemical behavior. The anolyte was a 1 mol/L LiTFSA-EC+DMC solution (EC: ethylene carbonate, DMC: dimethyl carbonate, volume ratio 1:1). This is referred to herein as the Catalytic Cycle Test. The amounts of the electrolyte at the both sides were 10 mL. The charging behavior of our cell was measured with an Asaka Charge/Discharge system (ACD-MO2) using a 0.05 mA (0.041 mA/cm2-cathode) discharge current at 25 °C. When the charge-voltage reached 4.4 V, the test was stopped. For reference, the Catalytic Cycle Test was also performed in electrochemical cell using DMSO or PC solution that dissolved iodine. The gases in the upper part of the cathode in Fig.S2 were measured by the GC/MS spectrometer. Analysis: Quantitative analysis of Li in the filtrate after the Li2CO3 chemical decomposition test was carried out by inductively coupled plasma atomic emission spectroscopy (ICP-AES) at Torey



Research Center. The residues after the Li2CO3 chemical decomposition test were examined by Raman analysis. Raman spectra were acquired with a JASCO laser Raman spectrophotometer (NRS 3300). The sample preparation was carried out in an argon-filled glove box as follows. The residues after the test were washed several times by dimethyl carbonate and dried under reduced pressure by using the side box of the glove box. Finally, the residues were set into a quartz cell (Figure S3) in the glove box. Laser irradiation to the chips was made through the argon-filled quartz cell. The wavenumber of the laser was 532 nm and the spot size was 30 µm in diameter. The degraded compounds were extracted by washing the residue into acetonitrile. The FTIR measurements were made under attenuated total reflection mode to identify them using a Nicolet FTIR spectrophotometer (Avatar360). The degraded compounds were also analyzed by NMR techniques (JEOL, JNM-ECA500). Computational details: The Gaussian program 09E01 was used to calculate the geometries of I2solvent 1:1 complexes. The geometries were fully optimized at the B3LYP/6-311++G(d,p) level.2123

3. Results and discussion 3.1 Redox potentials of iodine species Redox potentials of iodine species can be estimated from discharge/charge curves of Li-iodine cells or by cyclic voltammetry for organic electrolytes that contain iodine. We constructed a beaker cell composed of a Li anode, a carbon cathode, and an electrolyte with various solvents that dissolved 10 mg of iodine. The discharge/charge curves of the Li-I2 battery using PC, DMSO, or TMP as a solvent are displayed in Figure 1. The redox potentials of iodine species are discussed according to the plateaus of the charge voltage. In the PC electrolyte, two small shoulders of the


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charge voltage were observed at 3.21 V and 3.74 V vs. Li+/Li, which were due to redox couples of 3I- → I3- + 2e- and 2I3- → 3I2 + 2e-.24 Since the decomposition potential of Li2CO3 calculated from the Nernst equation based on the reaction Li2CO3 → 2Li+ + 2e- + CO2 + 0.5O2 is 3.82 V vs. Li+/Li, it is suggested that iodine in the PC solution (using the redox potential of 2I3- → 3I2 + 2e-) cannot oxidize Li2CO3. Two redox potentials also appeared in the electrolytes using TMP and DMSO. The voltage plateaus for the redox potential of 2I3- → 3I2 + 2e- were observed clearly at 4.11 V (TMP) and 3.93 V (DMSO). These results suggest the possibility that iodine in the TMP and DMSO solutions could oxidize Li2CO3. A surprisingly high redox potential for 2I3- → 3I2 + 2ewas also detected in the CV test. It was at 4.25 V in the TMP system, whereas the PC electrolyte showed a value of 3.31 V. The CV curves for the TMP and PC electrolytes are summarized in the Supporting Information (Figure S4a). Raman studies were made to examine electrolyte decomposition during discharge in Li-I2 system using TMP. The Raman results suggested no decomposition of TMP molecule (Figure S4b). The discharge/charge curves of electrochemical cells using IBr-TMP and ICl-TMP solutions are shown in Figure S5.

Figure 1. Discharge/charge curves of carbon cathode/Li metal anode cells using iodine solutions with various solvents: propylene carbonate (blue), dimethyl sulfoxide (red), and trimethyl phosphate (black).



3.2 Computational details The redox potentials of iodine species in TMP or DMSO solution were higher than those in PC solution. We performed computational studies to understand why the potentials increased in the TMP and DMSO solutions. It is well known that iodine molecules are complexed with solvent in a molar ratio of 1:1. The 1:1 complexes are due to the interaction between the  bond of iodine and a lone pair of electrons in the solvent, i.e., from the P=O, S=O, or C=O unit. We calculated the free energy of complex formation and the Mulliken electric charge distribution in the complex. The sum of electronic and thermal free energies, G, for halogen, solvent and complex are summarized in Table S1. Mulliken electric charge distributions for all elements in some 1:1 complexes are displayed in Figure S6. The electric charge distribution in I2-TMP 1:1 complex is plotted in Figure 2a. The two iodine and one oxygen atoms are negatively charged. Although the electric charges of the two iodine elements in the I2-TMP 1:1 complex were -0.028 and -0.198, the two iodine elements in the I2-PC 1:1 complex showed the electric charges of -0.131 and -0.135. Therefore, a large deviation of electric charge (0.170 = -0.028 - (-0.198)) existed in the former complex. The I2-DMSO 1:1 complex also exhibited a moderate deviation of electric charge (0.045 = -0.123 - (-0.168)). A similar deviation of electric charge would be expected in interhalogens such as IBr and ICl. For reference, we calculated Mulliken electric charge distributions for IBr-TMP and ICl-TMP 1:1 complexes and estimated their redox potentials from their discharge/charge curves for Li-interhalogen cells The electric charges on iodine and bromine elements in the IBrTMP 1:1 complex were +0.051 and -0.364, respectively. The ICl-TMP 1:1 complex exhibited +0.190 on iodine and -0.231 on chlorine. Thus, the DFT calculations suggested a large deviation in the interhalogens as predicted. We next investigated the correlation between the deviation of electric charge on iodine molecules and the redox potential of iodine species (Fig.2b). The


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horizontal axis in Fig. 2b represents the differences in electric charges between the halogen elements in the complexes. The redox potential increased depending on the increase in the difference between electric charges. These results suggest that the surprisingly high redox potentials of iodine species in TMP and DMSO solutions are caused by the deviation of electric charge due to complex formation.

a)

b)

Figure 2. (a) Mulliken electric charge distributions in I2-TMP 1:1 complex, and (b) relationship between redox potentials of carbon/Li cells and electric charge differences on halogen elements. The electric charges on halogens for iodine-solvent 1:1 complexes were obtained by DFT calculations. The numbers in brackets represent electrical charges of the elements. The red dotted line represents the decomposition potential of lithium carbonate, which was calculated by the Nernst equation based on the following reaction: Li2CO3 → 2Li+ + 2e- + CO2 + 0.5O2.

3.3. Chemical decomposition of Li2CO3 Figure 2b suggests that Li2CO3 will be decomposed by iodine in TMP or DMSO solution. Heating tests of iodine solutions with suspended Li2CO3 powders were conducted to examine the occurrence of Li2CO3 chemical decomposition. The test samples were stored at 25, 60 or 80 °C. Li2CO3 powders were present in overwhelming excess in the test samples compared to iodine. After the heating tests, the iodine solutions became lighter in color (Figure S7). The amount of Li in the filtrate was measured using inductively coupled plasma atomic emission spectroscopy (ICPAES) and the results are summarized in Table S2. The amount of Li was below the detection limit



in pure solvents free of iodine; therefore, Li2CO3 could not be dissolved in the solvents. On the other hand, Li on the order of micromoles was quantified in the filtrates of the TMP and DMSO solutions that contained iodine. Thus, the iodine in the TMP or DMSO solutions decomposed the Li2CO3 powder. A small amount of Li (0.39 mol) was detected in the PC filtrate. The mechanism for the dissolution of Li is not clear at the present stage. The amount of Li+ in the filtrate increased with temperature. The activation energy (Ea) estimated from the Nernst equation was Ea = 22.1 kJ/mol (Figure 3).

Figure 3. Arrhenius plot for decomposition reaction of Li2CO3 by iodine solutions: TMP (black) and DMSO (red).

3.4. Decomposition scheme The ICP-AES results indicate that Li2CO3 powders were decomposed by iodine. When iodine can oxidize Li2CO3, its decomposition pathway is clarified. The likely reaction is explained by Equation 1. In Equation 1, the generation of oxygen will be observed. When the pathway follows Equation 2 or 3, oxygen will not be detected. Therefore, we analyzed gases in the test tubes by GC-MS spectrometry after the heating test. 3I2 + Li2CO3 → 2I3- + 2Li+ + CO2 + 0.5O2

(1)


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3I2 + 2Li2CO3 → 2I3- + 2Li+ + 2CO2 + Li2O2

(2)

9I2 + 4Li2CO3 → 6I3- + 6Li+ + 4CO2 + 2LiO2

(3)

The gases in the test tubes after the heating test were analyzed using GC/MS spectrometry. The GC-MS spectrum for the I2-TMP system after the test at 60 °C is shown in Figure S8a. The signal at an elution time of 4.25 min was due to carbon dioxide, methane gas appeared at 10.3 min, and oxygen and nitrogen were detected at 9.30 and 9.67 min, respectively. The detection of N2 indicates that air penetrated the GC/MS spectrometer during sample gas injection. The amount of oxygen was corrected by subtracting the effect of the air. The following gases were detected: O2, 0.0015 mmol; CO2, 0.163 mmol; CH4, 0.0016 mmol; H2, 0.0017 mmol (Table S3). The amount of oxygen generated by decomposition in the TMP solution was 0.0015 mmol, whereas the amount of CO2 (0.163 mmol) was larger than that of oxygen by a factor of one hundred. The result indicates that the most likely oxidative reaction, 3I2 + Li2CO3 → 2I3- + 2Li+ + CO2 + 0.5O2, does not appear to be active in this experiment. In Eq. 2, lithium peroxide Li2O2 is formed in the Li2CO3 decomposition process. However, Li2O2 must be decomposed by iodine and subsequently oxygen will be detected. In Eq.3, an oxygen radical LiO2 is generated in the Li2CO3 decomposition process. Since LiO2 has high reactivity, it attacks solvent molecules. To identify the decomposition pathway, the Raman profiles of the residues in the test tube were measured. Raman spectra of the residues span the region from 100 to 3800 cm-1(Figure 4a). A strong signal in Fig. 3a is observed at 1091 cm-1 in the Li2CO3 sample. The signals from the residue at 114 and 170 cm-1 are due to the stretching vibrations of I- and I3-, respectively.25, 26 The signal at 346 cm-1 may be assigned to twisting mode of SO2. The residue exhibited no signal due to Li2O at 522 cm-1, which indicates that the decomposition of Li2CO3 to CO2 and Li2O did not occur. Raman profiles of commercial Li2O2, Li2O, and Li2CO3 are also



shown in Figure S9a. A signal due to the asymmetric stretching of SO2 was observed at 1464 cm-1. Figure 4b shows the enlarged Raman spectrum between 700 and 860 cm-1. In the residue sample, the peak top is given at 782 cm-1 and a shoulder appears in the vicinity of the peak top. The Raman profile located near 800 cm-1 was separated and curve-fitted by three components labeled A, B, and C with peaks at 782, 800, and 834 cm-1, respectively (Figure S10). The spectral assignments for the two signals at 782 and 834 cm-1 are not clear; however, the oxide radical, LiO2, would be generated by the Li2CO3 decomposition process (9I2 + 4Li2CO3 → 6I3- + 6Li+ + 4CO2 + 2LiO2). The LiO2 radical with high reactivity attacks the solvent and produces unidentified products. The signals for P-O symmetric stretching are given between 730 and 780 cm-1; therefore, the 782 cm-1 signal may be a degraded compound of TMP.27-29 The degraded product was analyzed by Raman, FTIR (Figure S11), and 1H, 13C, and 31P-NMR techniques (Figure S12). Based on the results of these analyses, we inferred degradation scheme and final products (Figure S13). The first step of the degradation scheme is the proton extraction from OCH3 unit in TMP by LiO2. Second, the oxygen molecule generated from HOO radical is added to OCH2 unit in TMP. The obtained material, (CH3O)2P(=O)O-CH2O2-Li+, is finally dimerized or polymerized through the two routes. The two final products are CH3O-[CH3OP(=O)O]2-Li and [(CH3O)2P(=O)-OC(=O)]2. The latter is the minor product, which is supported by small detection of C=O due to FTIR and 13C-NMR,

and a trace amount of hydrogen by GC/MS.

The Raman signal at 800 cm-1 may probably be due to vibration of Li2O2. The commercial Li2O2 in this study is characterized by signals at 792 and 260 cm-1. Zhang et al. reported that amorphous Li2O2 synthesized by a reaction of tetramethylammonium superoxide and LiClO4 had a Raman band of 788 cm-1 assigned to O-O stretching of Li2O2, and no remarkable signal due to crystalline Li2O2 near 260 cm-1.30 The peak at 800 cm-1 in Fig. S13 (red line) is most likely due to amorphous


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Li2O2, which suggests the possibility that the decomposition reaction (3I2 + 2Li2CO3 → 2I3- + 2Li+ + 2CO2 + Li2O2) occurred. However, its Raman signal due to vibration of Li2O2 was very weak. The existence of Li2O2 was also confirmed with a chromogenic method using titanium (IV) oxy sulfate [Ti(IV)OSO4].31 Li2O2 reacts with H2O in Ti(IV)OSO4 aqueous solution (Li2O2 + 2H2O → 2LiOH + H2O2). The subsequent reaction between H2O2 and Ti(IV)OSO4 forms a yellow [Ti(O22-)]2+ complex. When the residue after the Test was immersed in Ti(IV)OSO4 solution, the solution turned yellow lightly. The measurement of UV absorption spectrum for the yellow solution was conducted with a Shimadzu UV absorption spectrometer. It exhibited an absorption peak at 405 nm due to [Ti(O22-)]2+ complex (Figure 5). It followed from this evidence that the residue contained a small amount of Li2O2. The amounts of Li2O2 estimated by the absorbance of UV absorption spectrum were 0.24 μmol (TMP), and 0.13 μmol (DMSO) (Figure S13). In a coexistent system of Li2O2 and iodine, Li2O2 must be decomposed by iodine. The reason for the existence of Li2O2 in the residue may be a blocking effect of the degraded compound against attack by iodine species on Li2O2.

Figure 4. Raman spectra for the residues in the test tubes using I2-TMP (black) and I2-DMSO (red): (a) 100 - 3800 cm-1, and (b) 700 - 860 cm-1. The experimental data were curve-fitted by three components.



Figure 5. UV-VIS absorption spectra of Ti(IV)OSO4 solutions after the chromogenic reaction with the residue from TMP (black) and DMSO (red) solutions.

The following is a summary of the above. The chemical decomposition takes place in accordance with no oxygen evolution. The most common oxidation reaction of Li2CO3 (3I2 + Li2CO3 → 2I3+ 2Li+ + CO2 + 0.5O2) is unlikely in the presence of iodine. The decomposition pathway accompanied to the formation of Li2O2 described by Equation 2 is the minor one. Since the potential of Li2O2 is 2.96 V vs. Li+/Li, Li2O2 must be decomposed by iodine (Li2O2 + I2 →2Li+ + 2I- + O2); however this pathway is remote in the extreme because oxygen occurred very slightly. We propose a decomposition pathway via LiO2 radical (Figure 6). Yang et al. reported that CO2 was released during the electrochemical decomposition of Li2CO3 at 4.5 V, while oxygen was not detected in a Li-air/CO2 battery.14 They also proposed a mechanism for electrochemical decomposition through a LiO2 radical. The LiO2 radical caused degradation of the tetraglyme in their electrolyte. The chemical decomposition scheme in this study is something a little different from them. We treated LiO2 radical as electrically neutral material, The first step of the degradation is expressed by Equation 3. The LiO2 radical has high reactivity; therefore, it attacks the electrolyte directly, which leads to the formation of unknown products (Step 2). Mahne et al. and Schafzahl et al. detected singlet oxygen 1O2 during the discharge/charge processes in lithium-oxygen


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batteries.32,33 The singlet oxygen 1O2 is formed via LiO2 dimers or protons. There is a possibility that singlet oxygen 1O2 is the main cause of electrolyte degradation in the present system. Lithium oxygen radical reacts readily with water in the electrolyte, whereby LiOH and LiOOH are continuously produced.34-36 However, Raman spectroscopy studies in this work revealed no remarkable peaks between 3500 and 3700 cm-1. Therefore, it is quite unlikely that the reaction Li2O2 + H2O → LiOH + LiOOH proceeded. The formation of Li2O2 from LiO2 (2LiO2→ Li2O2+O2) is neglected.

Figure 6. Proposed mechanism via LiO2 radical for the decomposition of Li2CO3 by iodine.

3.5. Catalytic Cycle Test Finally, we combined the Li2CO3 chemical decomposition behavior with the redox couple of the iodine species, which works as a mediator. The Catalytic Cycle Test started from charging. Initially, triiodide ion I3- is generated by the reaction of I2 with Li2CO3. Second, the triiodine ion recovers iodine molecules at cathode upon charging of the cell. The iodine generated by electrical charging diffuses and reaches Li2CO3 powders. Since the iodine has a high redox potential by complex formation, it can decompose the Li2CO3 powders. This catalytic cycle can be repeated ideally until there is no Li2CO3 powder left. Figure 7 shows the charging curves for the cells using I2-PC, I2-DMSO, I2-TMP electrolytes, and TMP electrolyte free of I2. The cells with the DMSO



(red) and TMP (black) electrolytes exhibited a long plateau in cell voltage near 4.0 V, indicating that the catalytic cycle for chemical decomposition of Li2CO3 powders by iodine species was functioned. The plateau is lower than the decomposition voltages of Li2CO3 in the previous studies.37,38 The charge capacity up to 4.4 V in the TMP system was 15.88 mAh. If the charging capacity is only due to Li2CO3 decomposition, it corresponds to 0.296 mmol of Li2CO3 decomposition. Since the prepared amount of iodine in the TMP electrolyte was 0.0315 mmol, the catalytic cycle was repeated eight times (0.296/0.0315 = 9.39).

Figure 7. Charging curves of the cells using various iodine electrolytes with Li2CO3: TMP (black), DMSO (red), and PC (blue). The black dotted line represents the cell using TMP electrolyte free of iodine.

To confirm the decomposition pathway, the gases in the upper part of the cathode after the charging test were detected using GC/MS (Fig.S8b). The strong signal at an elution time of 4.25 min is due to carbon dioxide. Carbon monoxide was not detected. A small amount of hydrogen at 9.03 min was included. Oxygen and nitrogen appear at 9.30 and 9.68 min, respectively. The amount of oxygen was corrected by subtracting the effect of the air. The amount of oxygen generated by Li2CO3 decomposition in the TMP solution was 0.0023 mmol, whereas the amount of CO2 (0.284 mmol) generated was roughly 100 times more than that of oxygen. Almost all of the gas was CO2, which suggests that 0.284 mmol of Li2CO3 powder was decomposed by iodine. The amount of


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decomposed Li2CO3 is quantitatively matched with that estimated by the charge capacity (0.296 mmol). Thus, the Catalytic Cycle Test was started, which indicates that the charge capacity was produced slightly by other effects.

Conclusions In growing calls for electric vehicles, the re-use of lithium resources is required. In this paper, a catalytic cycle to reproduce lithium from lithium carbonate, which is one of the de-activated compounds in Li-ion batteries, was proposed. It has been accepted that iodine with a redox potential near 3.5 V cannot oxidize Li2CO3 with a decomposition potential of 3.82 V. We observed unexpectedly high redox potentials of iodine species in some organic solvents such as TMP and DMSO. The high redox potential was produced by deviation of the electric charge in the iodine molecule, which was brought about by its complexion with solvent. The iodine species with a high potential was applied as a redox mediator and combined with electrical charging. We demonstrated the continuous decomposition of Li2CO3 at room temperature. Thus, the iodine-solvent complex formation, which has been studied for a long time, opens the door toward lithium reproduction.

■ASSOCIATED CONTENT Supporting Information. Experimental setup, CV curves of iodine solutions, DFT calculation data of 1:1 complexes, Raman, FTIR, and NMR spectra, Li2CO3 decomposition data, Metal corrosion. This material is available free of charge via the Internet at http://pubs.acs.org. ■ AUTHOR INFORMATION Corresponding Author



* Tohru Shiga E-mail: [email protected]. ORCID Tohru Shiga: 0000-0001-7331-3380 Yoko Hase: 0000-0001-8805-7922 ■ NOTES The authors declare no competing financial interest. ■ Author Contributions T.S. conceived and carried out the experiments, analyzed the data and wrote the paper, Y.K performed DFT calculations, Y.K. carried out Raman analysis, and Y.H. proposed the catalytic cycle mechanism. ■ ACKNOWLEDGMENTS The authors thank Keiko Fukumoto of Toyota CRDL for NMR measurements. ■ REFERENCES (1) Kim, J-S.; Lee, Y-H.; Choi, S.; Shin, J. Dinh, H-C.; Choi, J.W. An Electrochemical Cell for Selective Lithium Capture from Seawater. Enviro. Sci. technol. 2015, 49, 9415-9422. (2) Ryu, T.; Haldorai, Y.; Rengaraj, A.; Shin, J.; Hong, H-J.; Lee, G-W.; Han, Y-K.; Huh, Y.S.; Chung, K-S. Recovery of Lithium Ions from Seawater Using a Continuous Flow Adsorption Column Packed with Granulated Chitosan-Lithium Manganese Oxide. Ind. Eng. Chem. Res. 2016, 55, 7218-7225. (3) Xu, K. “Electrolytes and Interphases in Li-ion Batteries and Beyond” Chem. Rev. 2014, 114, 11503-11618. (4) Aravindan, V.; Gnanaraj, J.; Lee, Y-S.; Madhavi, S. Insertion-Type Electrodes for Nonaqueous Li-Ion Capacitors. Chem. Rev. 2014, 114, 11619-11635. (5) Reddy, M.V.; Subba Rao, G.V.; Chowdari, B.V.R. Metal Oxides and Oxysalts as Anode Materials for Li Ion Batteries. Chem. Rev. 2013, 113, 5364-5467. (6) Nakahara, K.; Iwasa, S.; Satoh, M.; Morioka, Y.; Iriyama, J.; Suguro, M.; Hasegawa, M. Rechargeable Batteries with Organic Radical Cathodes. Chem. Phys. Lett. 2002, 359, 351-354. (7) Choi, W.; Harada, D.; Oyaizu, K.; Nishide, H. Aqueous Electrochemistry of Poly(ninylantraquinone) for Anode-Active Materials in High-Density and Rechargeable Polymer/Air Batteries. J. Am. Chem. Soc. 2011, 133, 19838-19843. (8) Zhu, Y.G.; Wang, X.; Jia, C.; Yang, J.; Wang, Q. Redox-Mediated ORR and OER Reactions: Redox Flow Lithium Oxygen Batteries Enabled with a Pair of Soluble Redox Catalysts. ACS Catal. 2016, 6, 6191-6197. (9) Hase, Y.; Ito, E.; Shiga, T.; Mizuno, F.; Nishikoori, H.; Iba, H.; Takechi, K. Quantitation of Li2O2 Stored in Li-O2 Batteries Based on Its Reaction with an Oxoammonium Salt. Chem. Commun. 2013. 49,


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Iodine Mediator with Anomalously High Redox Potential and Its

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