Crown Ethers as Electrolyte Additives to Modulate Electrochemical

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Crown Ethers as Electrolyte Additives to Modulate Electrochemical Potential of Lithium Organic Batteries Yi-Fei Yang, Chun-Yu Chiou, Chuan-Wen Liu, Cheng-Lung Chen, and Jyh-Tsung Lee J. Phys. Chem. C, Just Accepted Manuscript • Publication Date (Web): 19 Aug 2019 Downloaded from pubs.acs.org on August 19, 2019

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The Journal of Physical Chemistry

Crown Ethers as Electrolyte Additives to Modulate Electrochemical Potential of Lithium Organic Batteries Yi-Fei Yanga, Chun-Yu Chioua, Chuan-Wen Liub, Cheng-Lung Chena, Jyh-Tsung Leea,c* a

Department of Chemistry, National Sun Yat-sen University, Kaohsiung 80424, Taiwan b

School of Defense Science, Chung Cheng Institute of Technology, National Defense

University, Dasi, Taoyuan 33551, Taiwan c

Department of Medicinal and Applied Chemistry, Kaohsiung Medical University, Kaohsiung 80708, Taiwan *

Corresponding author E-mail: [email protected]

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ABSTRACT Organic batteries have attracted much attention because of their flexibility, highpower densities, and highly designable structures of electrode-active materials. The electrochemical potential of the batteries can be modulated using different organic redox species or by structural modification of the redox unit centers. In this study, the electrochemical potential of lithium organic radical batteries is modulated, without the structural modification of the redox unit centers. Two different crown ethers, 12-crown4 (12C4) and 15-crown-5 (15C5), served as electrolyte additives to increase the electrochemical potential of the batteries. An average discharge voltage can be increase to 3.90 V through the electrolyte system using various concentrations of the electrolyte salt and crown ethers. The addition of one equivalent of 12C4 to the Li︱0.25 M LiClO4-ethylene

carbonate/diethyl

carbonate

(=

1/1,

v/v ︱ poly(2,2,6,6-

tetramethylpiperidin-1-oxy-4-yl methacrylate) cells significantly improved the capacity retention up to 21% after 300 cycles at a current rate of 3C. Furthermore, the structures and system energies of the lithium-crown ether complexes are investigated using density functional theory calculation.


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INTRODUCTION Organic batteries have attracted much attention because of their flexibility, high-power densities, resource abundance, low-cost, thin-film electrode, and long cycle-life.1-4 Poly(2,2,6,6-tetramethylpiperidin-1-oxy-4-yl methacrylate) (PTMA, see Figure S1) is a positive electrode-active material for organic batteries. PTMA contains a robust radical 2,2,6,6-tetramethylpiperidin-1-oxyl (TEMPO) group that is quite stable even under ambient conditions, due to the steric hindrance of its four methyl groups near the nitroxide center.5-7 PTMA exhibits a promising specific energy capacity of 111 mAh g– 1

and a high redox potential of 3.6 V vs. Li/Li+. Moreover, the PTMA electrode has

good C-rate performance during the charge-discharge process due to the small conformational changes between the redox species and the fast bimolecular electron self-exchange reaction rate in the polymer chain.6, 8-10 Today, electronic devices require relatively high energy densities and high power densities of rechargeable batteries. Increasing the working voltage of batteries is one of important approaches to increase the energy densities of batteries. The working voltages of batteries can be increased using different structures of positive electrodeactive materials with relatively high redox voltages. For example, in lithium-ion batteries, high-voltage positive electrode materials, such as LiNi0.5Mn1.5O4, LiCoPO4, Li3V2(PO4)3, Li2CoPO4F, and LiNixCoyMn1−x−yO2, have been studied.11 For organic radical batteries, several positive electrode-active materials have higher working voltage than conventional PTMA.8, 12 Moreover, by adjusting the functional groups near the radical center, different redox potentials of the nitroxide polymer can be obtained.13 For example, in poly[4-(N-tert-butyl-N-oxylamino)-3-trifluoromethylstyrene], the incorporation of the electron-withdrawing trifluoromethyl group would slightly increase its redox potential. Another approach is to copolymerize the TEMPOcontaining monomer with an anion-bearing monomer to form a copolymer such as 3 ACS Paragon Plus Environment


poly(TEMPO-methacrylate-stat-vinylsulfonic acid), in which the redox potential of the copolymer has been increased to 3.94 V with a specific capacity of 42 mAh g–1.14 However, these non-radical anion-bearing groups cannot participate in the redox reaction; thus, the incorporation of these groups will reduce the specific energy due to the additional molecular weight. Recently, concentrated electrolytes have gained much attention because of their wide electrochemical windows.15 However, few studies on the effects of the lithium salt concentration on the electrochemical potential of batteries have been reported. The lithium salt concentration affects the redox potential of nitroxide polymer positive electrodes.14,

16-17

The chelation of the lithium ion also affects the electrochemical

potential of batteries. Reportedly, glyme (CH3-O-(CH2-CH2-O)n-CH3) switches the redox potential upward in lithium batteries,18 Further, in organic radical batteries, the different concentrations of electrolytes are investigated.16-17 A high salt concentration can suppress the self-discharge of PTMA.17 Furthermore, the salt concentration affects the electrochemical potential of the batteries. However, the exact mechanism and qualification have not been fully clarified. In this study, the electrochemical behaviors of lithium organic radical batteries with different concentrations of lithium salt are investigated. Two different ionic solvation enhancers, 12-crown-4 (12C4) and 15crown-5 (15C5), of lithium salts act as additives to the electrolyte to modulate the electrochemical potential of lithium organic batteries. Furthermore, computer simulations are combined with experimental data to understand clearly the role of the electrolyte system in modulating the redox potential. This sheds light on that the redox potential can be regulated by the electrolyte composition.


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EXPERIMENTAL SECTION Materials. Ethyl 2-bromoisobutyrate (98%, Alfa Aesar), 2,2,6,6-tetramethylpiperidin-4-yl methacrylate (TMPM, 98%, TCI), 1,1,4,7,10,10-hexamethyltriethylenetetramine (HMTETA, 97%, Sigma-Aldrich), copper (I) bromide (CuBr, 99.999%, Sigma-Aldrich), lithium perchlorate (LiClO4, 99%, Alfa Aesar), tetrabutylammonium perchlorate (Bu4NClO4, > 98%, TCI), 1-methyl-2-pyrrolidinone (NMP, 99.9%, Mallinckrodt), ethylene carbonate (EC, 99%, Alfa Aesar), diethyl carbonate (DEC, > 99%, Alfa Aesar), poly(vinylidene fluoride) (PVDF, KF-1100, Kureha), Super P carbon black (Timcal), 12-crown-4 (12C4, 98%, Acros), and 15crown-5 (15C5, 98%, Acros) were used as received. m-Chloroperoxybenzoic acid (mCPBA, 70−75%, Acros) was purified by diethyl ether extraction. Instrumentation. The electrochemical measurements were performed on CHI 750A and CHI 627D electrochemical analyzers. Synthesis of poly(2,2,6,6-tetramethylpiperidin-4-yl methacrylate) (PTMPM). PTMPM was polymerized via atom transfer radical polymerization (ATRP).19 Ethyl 2-bromoisobutyrate (3 μL, 0.0204 mmol) as an initiator and HMTETA (5.5 μL, 0.020 mmol) as a ligand were dissolved in acetone (8 mL), and subsequently degassed by three freeze-pump-thaw cycles. The acetone solution was transferred by a cannula into a degassed flask containing CuBr (3.1 mg, 0.022 mmol). After sonication for 10 min, the resulting solution was once more degassed by two freeze-pump-thaw cycles, after which it was transferred to a degassed flask containing the TMPM monomer (2.764 g, 12.266 mmol) to undergo ATRP, with stirring at 60 °C for 24 h under nitrogen atmosphere. After polymerization, the solution was filtered through a column of aluminum oxide to remove the catalyst. The acetone solution in the filtrate was mostly removed using a rotavapor, and tetrahydrofuran (10 mL) was added as a diluent. Subsequently, the solution was dropped in hexane (200 mL) to precipitate the polymer. The PTMPM polymer was separated and dried at 60 °C for 12 h under vacuum.

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Oxidation of PTMPM to PTMA. A 20-mL tetrahydrofuran solution of mCPBA (690.2 mg, 4 mmol) was added dropwise to a dichloromethane (20 mL) solution of PTMPM (450.7 mg, 2 mmol) in a water-ice bath. The resulting solution was stirred at room temperature for 3 h, after which it was transferred to a separatory funnel containing a 40-mL aqueous solution of NaHCO3 (672.1 mg, 8 mmol), to quench the oxidation reaction. After shaking and phase separation, the organic phase was washed 3 times with deionized water, dried over anhydrous MgSO4, and concentrated to 10 mL using a rotavapor. The dichloromethane solution was added to hexane (200 mL) to precipitate the polymer. After the precipitation, the PTMA polymer was dried at room temperature under vacuum. Electrochemical measurements. The PTMA composite was prepared by mixing PTMA/PVDF/Super P in a 1:1:8 weight ratio. NMP (1.5 mL) was added to the sample vial containing PTMA (20 mg) and PVDF (20 mg), and the resulting solution was stirred for 1 h at room temperature until a homogenous solution was obtained. Subsequently, Super P (160 mg) was added to the previous solution and stirred at room temperature for 8 h to form a slurry. A doctor-blade coating method was employed to cast PTMA composite electrodes. For the electrodes of electrochemical cells, the slurry was coated on an Au-coated silicon wafer (Ti as the adhesion layer); for the electrodes of coin cells, the slurry was coated on an aluminum foil. The mass loading of positive electrode is about 0.5 mg cm-2. For the cyclic voltammogram measurements for carbonate-based electrolytes, CR2032 coin cells were used. A PTMA-coated aluminum foil was used as a working electrode, while a lithium foil was used as both a counter electrode and a reference electrode. Without crown ether systems, the electrolytes were 0.1, 0.25, 0.5, and 1.0 M LiClO4 in EC/DEC (1/1, v/v). With crown ether systems, the electrolytes were 0.25 M LiClO4 with 1, 2, and 4 equivalents of 12C4 or 15C5 in EC/DEC (1/1, v/v).

The

cyclic voltammograms were measured between 3.0 V and 4.2 V vs. Li/Li+ at a scan rate of 0.1 mV s−1. For the cyclic voltammogram measurements of CH3CN-based electrolytes, a three-


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electrode system was used in an under an Ar-filled glovebox. A PTMA electrode was used as a working electrode; Pt was used as a counter electrode and a reference electrode. Ag/AgNO3 (0.01 M AgNO3 and 0.1 M Bu4NClO4 in CH3CN) was used as a reference electrode. Without crown ether systems, the electrolytes were 0.25, 0.5, and 1.0 M LiClO4 (or Bu4NClO4)-CH3CN; with crown ether systems, the electrolytes were 0.25 M LiClO4-CH3CN with 0, 1, 2, and 4 equivalents of 12C4.

The cyclic voltammograms were measured between 0.2 V and 0.7 V vs.

Ag/Ag+ at a scan rate of 5 mV s–1. For the charge-discharge and cycle-life experiments, CR2032 coin cells were used. The PTMA electrode served as the working electrode and a lithium foil served as both the counter electrode and the reference electrode. A Celgard 2500 separator was placed between the electrodes. The electrolyte was 0.25 M LiClO4 with 0, 1, and 2 equivalents of 12C4 or 15C5 in EC/DEC (1/1, v/v). The coin cells were fabricated under an Ar-filled environment. The cells were charged/discharged and cycled at a current rate of 3C. Theoretical methodology. To simulate the influence of crown ethers and LiClO4 on the redox potential of lithium-organic batteries, density functional theory (DFT) calculations were performed using the Gaussian0920 package. The fully optimized geometry and total electronic energy of each system were obtained using the B3LYP21-22 hybrid functional with the basis set of 6-31G(d,p), and the bulk solvent effect was considered using polarized continuum model (PCM).23-24 The whole redox reaction of a lithium-organic battery can be expressed as eq 1. (R+ + ClO4 − )p + Li(M) ⇌ (R∙ )p + (ClO4 − )s + (Li+ )s

(1)

where p denotes the polymer layer and s denotes the electrolyte solution. During the discharge process, the lithium metal anode is oxidized along with a phase change from a metal solid (Li(M)) to a cation solvated in the electrolyte solution (Li+(s)): Li(M) → (Li+ )s + e−

(2)

Meanwhile, at the positive electrode, the oxidized PTMA+ undergoes reduction:



(R+ + ClO4 − )p + e− → (R∙ )p + (ClO4 − )s

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(3)

Unlike the lithium metal anode, PTMA remains in the polymer layer with no phase change during the redox reaction. To calculate the redox potential of the complete reaction (Eredox), the absolute reduction potential (Eabs) of each half-reaction is adopted as shown in eq 4. 𝐸redox = 𝐸abs (R+ − R∙ ) − 𝐸abs (Li+ − Li)

(4)

in which the absolute reduction potential is equal to the change in the Gibbs free energy of each state during the reduction reaction (△Greduc.): 𝐸abs =

−Δ𝐺reduc. 𝑛𝐹

(5)

where n is the number of transferred electrons and F is the Faraday constant. To quantify the Gibbs free energy change of lithium, a thermodynamic cycle for the lithium electrode reaction is applied as follows.25-26

Scheme 1. Thermodynamic (Born-Haber) cycle for the lithium electrode during the reduction reaction. (ΔGreduc denotes the difference in Gibbs free energy of the lithium reduction reaction; ΔGsolv is the solvation free energy of the lithium ion; ΦM is the work function of the inert metal electrodes; ΔGe is the ionization energy; ΔGvap is the vaporization free energy; G(Li+(solv.)) is the Gibbs free energy of the lithium ion in solution.)

The reduction of the lithium ion can be represented in the Born-Haber cycle (Scheme 1). Because the free energy is a state function, it is independent of the path. Therefore, the Gibbs free energy of the reduction reaction can be written as eq 6:


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(6)

∆𝐺reduc. = −∆𝐺solv. + ΦM − ∆𝐺e − ∆𝐺vap. + + = 𝐺(Li+ (g) ) − 𝐺(Li(solv.) )+ΦM + 𝐺(Li(g) ) − 𝐺(Li(g) ) − ∆𝐺vap.

= 𝐺(Li(g) ) − 𝐺(Li+ (solv.) )+ΦM − ∆𝐺vap. = −𝐺(Li+ (solv.) ) + 𝑐𝑛𝑠𝑡. Obviously, according to eq 6, the difference between the Gibbs free energy change of the lithium reduction reaction in water and in the EC/DEC electrolyte solution relies only on the variation in the free energy of the lithium ion in the solution. Since (1) the relative standard reduction potential of lithium, E° (Li+ – Li), is equal to –3.04 V vs. standard hydrogen electrode (SHE) from the experimental data, and (2) the up-to-date value of the computed standard hydrogen potential, Eabs° (SHE), is equal to 4.28 V,27-28 the absolute reduction potential of lithium in water can be obtained by summing up the two terms, i.e., Eabsaqueous (Li+ – Li) = 1.24 V. Therefore, the absolute reduction potential of lithium in the EC/DEC electrolyte solution can also be efficiently obtained using this method as expressed in eq 7.29-30 aqueous

𝐸abs (Li+ − Li) = 𝐸abs

(Li+ − Li)

(7)

+ +(𝐸tot (Li+ solvent ) − 𝐸tot (Liaqueous ))

B3LYP/6-31G(d,p), PCM

At the PTMA electrode, however, due to the unavailability of a gas phase for the polymer, the thermodynamic cycle is not suitable for the quantification of the absolute reduction potential of PTMA. Considering that (1) there is no phase change of the polymer throughout the redox reaction, and (2) the difference between the vibrational and configurational entropy contributions of both the reduced and oxidized states of the redox radical unit at room temperature is assumed to be very small,29, 31 the Gibbs freeenergy change of the PTMA electrode during the reduction is approximated by the total electronic energy change between



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the reduced and the oxidized states, (Etot (R) and Etot (R+), i.e., the adiabatic ionization, as listed in eq 8. 𝐸abs (R+ − R∙ ) =

−Δ𝐺 0 ≃ −(𝐸tot (R∙ )−𝐸tot (R+ )) 𝑛𝐹

(8)

In addition, since the backbone structure of the polymer chain is assumed to have little impact on the redox potential,32-33 a methyl-terminated version of the monomer unit was adopted in the DFT calculation, and the dielectric constant of the electrolyte solution EC/DEC (1/1, v/v) was set to 33.6 at 298 K.34 Therefore, combined with eqs 7 and 8, the redox potential of the lithium organic battery can be calculated.


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RESULTS AND DISCUSSION Ionic Conductivity. The ionic conductivity of the electrolyte is one of the key factors that determine the power density of a lithium-ion battery. To understand the electrolyte system further, the ionic conductivity of different concentrations of LiClO4 and various equivalents of 12C4 and 15C5 (with respect to the mole of lithium ion) in EC/DEC (1/1, v/v) was measured by AC impedance. The ionic conductivity (σ) of an electrolyte can be expressed by the following equation35: σ = ∑ 𝑛𝑖 𝑢𝑖 𝑍𝑖 𝑒

(9)

𝑖

where ni is the free ion number, ui is the ionic mobility, Zi is the ionic charge, and e is the electron charge. 𝑢𝑖 =

1 6𝜋𝜂𝑟𝑖

(10)

where η is the viscosity of the media and ri is the solvation radius of i. When the concentration of lithium salt is relatively low (< 1.0 M), the change in the viscosity can be neglected.36 Therefore, when the concentration of the salt (i.e., free ion number) increases from 0.1 to 1.0 M, the ionic conductivity increases by 235% as shown in Figure 1 (a). In the study of lithium-crown ethers, the concentration of the lithium salt is 0.25 M because the addition amount of crown ether is considered. In contrast, as 12C4 and 15C5 are regarded as the most efficient chelating ligands toward the lithium ion due to the fitness of their cavities (12C4: 1.5~1.7 Å ; 15C5: 1.7~2.2 Å )37 to the ionic diameter of the lithium ion (1.72 Å )38, adding these crown ethers to a 0.25 M LiClO4 solution can facilitate the dissociation of the lithium perchlorate ion pairs, thereby enhancing the ionic conductivity of the electrolyte. The lithium ion is solvated by the crown ethers to form lithium complexes confirmed by Raman spectra39 (Figure S5). For the 12C4 system, the ionic conductivity reaches its maximum with 2 equivalents of 12C4 (5.58 mS cm–1), after which it decreases with 4 equivalents of 12C4 (4.90 mS cm–1) (see Figure 1 (b)); for the 15C5 system, the ionic conductivity is maximized with 1



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equivalent of 15C5 (5.70 mS cm–1), slightly decreases to 5.17 mS cm–1 with 2 equivalents of 15C5, and further decreases to 4.35 mS cm–1 with 4 equivalents of 15C5 (see Figure 1 (c)). The stability of the ion-crown ether complexes increases in the order of [Li(12C4)+] ＞ [Li(15C5)+], which is reasonably due to the fitness of the 15C5 cavity and the higher coordinating oxygen number of 15C5. However, when the complexations of the crown ethers and the lithium ions reach saturation points, excessive amounts of 12C4 (4 equivalents) and 15C5 (2 and 4 equivalents)40 result in decreased ionic mobility due to the increased viscosities of 12C4 and 15C5 (9.54 mPa·s and 21.65 mPa·s, respectively, Table 1).

Figure 1. Ionic conductivity of (a) 0.1, 0.25, 0.5, and 1.0 M LiClO4 in EC/DEC (1/1, v/v); 0.25 M LiClO4 with 0, 1, 2, and 4 equivalents of (b) 12C4 and (c) 15C5 in EC/DEC (1/1, v/v).

Table 1. Viscosities and dielectric constants of EC, DEC, 12C4, and 15C5 at 298.15 K. substance

viscosity (mPa·s)

dielectric constant

references

EC

1.85 (40 °C)

95.3

34, 41

DEC

0.75

2.82

15,42-43

12C4

9.54

10.82

44-45

15C5

21.65

15.48

44-45

Electrochemical properties. Figure 2 shows the cyclic voltammograms of the PTMA electrode with different concentrations of LiClO4 at a scan rate of 0.1 mV s-1. The difference between the anodic and the cathodic peak potentials (Epa and Epc) of the PTMA electrode in


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0.5, 1.0, and 2.0 M LiClO4-EC/DEC (1/1, v/v) electrolytes is 6, 10, and 16 mV, respectively, which is a Nernstian adsorbate-like behavior.46-47 When the scan rate (10 mV s-1) is too fast, the difference between the Epa and Epc potentials of the PTMA electrode dramatically increases (Figure S6), which may due to the viscosity of the electrolyte and the diffusivity of the charge carriers. Notably, the voltages of Epa and Epc decline at high concentrations of LiClO4 in the electrolyte solution. The mid-peak potential (Emid = (Epa + Epc)/2), which is often regarded as the redox potential in the electrochemically reversible reaction, is reduced to 3.608 V vs. Li/Li+ in the 2.0 M LiClO4 electrolyte solution. In addition, the cathodic and anodic peak currents (ipc and ipa) decrease significantly in the 2.0-M LiClO4 solution, which may be ascribed to the increasing ionic aggregation of the LiClO4 ion pairs and the relatively high viscosity of the 2.0 M electrolyte solution. The shift of the redox peaks may be due to the lithium anode, which is used as both the reference electrode and counter electrode in the two-electrode electrochemical cell, or to the concentration of the lithium salt. To clarify the shift, a three-electrode electrochemical cell employing the PTMA composite electrode as the working electrode, and a Pt wire and Ag/AgNO3 (0.01 M AgNO3 and 0.1 M Bu4NClO4 in CH3CN) as counter and reference electrodes, respectively, were adopted in both Bu4NClO4 and LiClO4 electrolyte systems. Figures 3a and b show the cyclic voltammograms of the PTMA electrodes with different concentrations of Bu4NClO4 and LiClO4 in CH3CN. A similar shift of Emid is observed, indicating that this decrease in the redox potential may be mainly owing to the increased concentration of ClO4– in the electrolyte solution with respect to the PTMA electrode, and the half-reaction at the PTMA electrode interface (eq 11) can be described by the Nernst equation (eq 12): (R∙ )p + (ClO4 − )s ⇌ (R+ + ClO4 − )p + e−

(11)

where R is the nitroxide radical; R+ is the oxoammonium cation, p is the PTMA polymer layer, and s is the electrolyte solution.



𝐸PTMA =

°′ 𝐸PTMA

[R+ ClO4 − ]p 𝑅𝑇 + 2.3 log ( ∙ ) 𝐹 [R ]p [ClO4 − ]s

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(12)

where E is the redox potential, E°’ is the formal potential, R is the gas constant, T is the absolute temperature, and F is the Faraday constant.

At the half-wave potential where [R] is equal to [R+], E1/2 can be expressed as below: °′ 𝐸1/2 = 𝐸PTMA − 2.3

𝑅𝑇 log[ClO4 − ]s 𝐹

(13)

As can be deduced from eq 13, an increase in [ClO4–]s results in a decrease in E1/2. In contrast, an upward shift of Emid is observed in Figure 4 when the crown ethers are added to the 0.25 M LiClO4-EC/DEC (1/1, v/v) electrolyte. Relatively large equivalents of crown ethers in the electrolyte solution containing a fixed concentration of ClO4– in the two-electrode electrochemical cell are associated with high redox potentials (3.93 V and 4.04 V for 4 equivalents of 12C4 and 15C5, respectively). This ability of crown ethers to modulate the redox potential cathodically in the lithium organic battery is quite interesting and has not been reported previously. To investigate the cause of this phenomenon, the aforementioned threeelectrode electrochemical cell was adapted as well.

Figure 2. Cyclic voltammograms of the PTMA electrode in (a) 0.5, (b) 1.0, and (c) 2.0 M LiClO4-EC/DEC (1/1, v/v) electrolytes at a scan rate of 0.1 mV s–1. Li was used as the counter and reference electrodes.


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Figure 3. Cyclic voltammograms of (a) PTMA electrodes in 0.1, 0.5, and 1.0 M Bu4NClO4CH3CN electrolytes and (b) PTMA electrode in 0.25, 0.5, and 1.0 M LiClO4-CH3CN electrolytes at a scan rate of 5 mV s–1. Ag/AgNO3 (0.01 M AgNO3 and 0.1 M Bu4NClO4 in CH3CN) was used as a reference electrode; Pt was used as a counter electrode.

As shown in Figure 5, if the reference electrode was not lithium metal, the cathodic shift of Emid of the cells with the 12C4 additive would not occur; instead, a slight anodic shift of Emid has been observed (from 0.44 V to 0.43 V for 0 and 2 equivalents of 12C4, respectively), implying that the upward redox potential is mainly due to the effect of 12C4 on the lithium metal anode interface; otherwise, a downward mediation would occur due to the ionic solvation



enhancing ability of 12C4 that results in a slight increase in [ClO4–] in the electrolyte solution, and, in turn, a minor decline in Emid.

Figure 4. Cyclic voltammograms of PTMA electrodes in 0.25 M LiClO4-EC/DEC (1/1, v/v) electrolytes with 0, 1, 2, and 4 equivalents of (a) 12C4 and (b) 15C5 at a scan rate of 0.1 mV s–1. Li was used as the counter and reference electrodes.


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Figure 5. Cyclic voltammograms of the PTMA electrode in 0.25 M LiClO4-CH3CN electrolytes with 0, 1, 2, and 4 equivalents of 12C4 at a scan rate of 5 mV s–1. Ag/AgNO3 (0.01 M AgNO3 and 0.1 M Bu4NClO4 in CH3CN) was used as a reference electrode; Pt was used as a counter electrode.

According to the half-reaction at the interface of the lithium metal anode (eq 14), the redox potential of the lithium anode (ELi) can be expressed using the Nernst equation as follows: Li(M) ⇌ (Li+ )s + e−

(14)

(Li(M): lithium metal) °′ 𝐸Li = 𝐸Li + 2.3

𝑅𝑇 log[Li+ ]s 𝐹

(15)

Due to the strong chelating abilities of 12C4 and 15C5 toward the lithium ions, the incorporation of crown ethers will reduce the free lithium-ion concentration in the solution ([Li+]s); thus, the redox potential of the lithium anode increases. This difference between the effects of the presence and absence of the crown ethers can be observed clearly in Figures 6 and 7. Consequently, since the cell voltage is equal to the subtraction of the positive electrode potential from the anode potential, the cell voltage of the lithium organic battery will be increased when the anode potential is decreased.



Figure 6. Schematic of the lithium ions and perchlorate anions transfer during the charging process in the Li︱ LiClO4-EC/DEC︱PTMA cell.

Figure 7. Schematic of lithium ions and perchlorate anions transfer during the charging process in the Li︱ LiClO4-EC/DEC with 1 equivalent of crown ether︱PTMA cell.


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DFT calculation. The Eabs values of PTMA with counter anions and that of lithium with 12C4 and 15C5, calculated by DFT calculations using B3LYP, are summarized in Table 2. The Eabs values of the PTMA electrode exhibit a descending trend (4.82, 4.63, and 4.54 V) with 0, 1, and 2 equivalents of ClO4–, which is attributed to the stabilizing effect of ClO4– on the radical cation (see the charge distribution maps in Figure 8). At the lithium anode, where the lithium ion is complexed to the crown ether, the Eabs of the lithium anode shifts from 1.31 to 0.97 V for [Li(12C4)]+ and to 0.88 V for [Li(15C5)]+. The optimized structures of these complexes are shown in Figure 9, and the relative calculated results are listed in Table 3. Obviously, the total energy of the [Li(15C5)]+ complex is lower than that of the [Li(12C4)]+ complex, indicating the higher stability due to the better fit of the 15C5 cavity to the ionic diameter of the lithium ion and to the higher number of donating oxygen atoms coordinated with the lithium ion. Overall, the redox potential of the lithium organic battery is up-shifted with the complexations of the lithium ion to 12C4 (3.85 V) and to 15C5 (3.94 V), and down-shifted with the stabilization of ClO4– to the radical cation. The results of the DFT calculations are in good agreement with the experimental data.



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Table 2. Absolute reduction potential (Eabs) of PTMA with counter anions, lithium with 12C4/15C5, and the redox potential of the lithium organic battery calculated at the B3LYP level of theory with the basis set of 6-31G(d,p) using the Gaussian09 package. PTMA electrode System

Redox Potential

Eabs (R+– R.)

Eabs (Li+ – Li)

Eabs, 12C4 (Li+ – Li)

Eabs, 15C5 (Li+ – Li)

Eredox

Eredox, 12C4

Eredox, 15C5

(V)

(V)

(V)

(V)

(V)

(V)

(V)

3.51

3.85

3.94

3.33

3.67

3.76

3.23

3.57

3.66

PTMA repeating unit

4.82

PTMA repeating unit + ClO4– × 1

4.63

PTMA repeating unit + ClO4– × 2

4.54

(a)

Li electrode

1.31

0.97

0.88

(b)

(c)

Figure 8. Electrostatic potential maps (isovalue = 0.02): (a) PTMA repeating unit; (b) PTMA repeating unit with a counter anion of ClO4–; (c) PTMA repeating unit with two counter anions of ClO4– (red, blue, grey, light grey, and green spheres represent oxygen, nitrogen, carbon, hydrogen, and chlorine atoms, respectively).


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Figure 9. Optimized structures of (a) [Li(12C4)]+ and (b) [Li(15C5)]+ calculated by the B3LYP level of theory with the basis set of 6-31G(d,p) in conjunction with PCM.

Table 3. Parameters of the Li+ complexes with different crown ethers in EC/DEC (1/1, v/v) calculated by the B3LYP level of theory with the basis set of 6-31G(d,p) in conjunction with PCM. System

Etot (Hartree)a

q(Li+)b

rLi-O (Å )c

O-Li-O bond anglec

[Li(12C4)]+

−622.604152

0.408

1.89

92.08

[Li(15C5)]+

−776.396757

0.432

2.20

a

b

74.28 + c

Total energy of the complexes. Charge distribution on Li . Average value of bond lengths and bond angles.



Cell performance. To further check the electrochemical performances of Li ︱0.25 M LiClO4-EC/DEC (1/1, v/v) with 0, 1, and 2 equivalents of crown ethers︱PTMA coin cells, charge-discharge and galvanostatic cycling tests were conducted. As shown in Figure 10, the discharging plateau shifts to higher voltages (from 3.78 to 3.90 V) in the presence of 0 to 2 equivalents of 12C4 in the electrolyte solution. Moreover, when 1 equivalent of 12C4 is added to the electrolyte, the discharging capacity at a 3C rate reached 92.3 mAh g–1. However, when the amount of 12C4 increased to 2 equivalents, the discharging capacity decreases to 72.9 mAh g–1 (by ~20%). From the cycle-life performances of the coin cells charged/discharged at a current rate of 3C (see Figure 11), the cycle life of the cells is slightly decay, which may be due to that the molecular weight of PTMA is not high enough.48 However, it is noted that all those with 1 and 2 equivalents of crown ethers exhibit higher capacity retentions after 300 cycles in comparison with those without the crown ethers; the cell with 1 equivalent of 12C4 possesses the highest capacity retention (72%) and the highest Coulombic efficiency. The superior cycle-life performances may be ascribed to the enhanced lithium anode cyclability due to the complexation of the lithium ions to the crown ethers, which suppresses the formation of lithium dendrite.49


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Figure 10. Charge-discharge characteristics of Li︱0.25 M LiClO4-EC/DEC (1/1, v/v) with (a) 0, (b) 1, and (c) 2 equivalents of 12C4︱PTMA coin cells at a charge/discharge rate of 3C.



Figure 11. Cycle-life performances of Li︱0.25 M LiClO4-EC/DEC (1/1, v/v) with (i) 0, (ii) 1, and (iii) 2 equivalents of (a) 12C4 and (b) 15C5︱PTMA coin cells at a charge/discharge rate of 3C.


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CONCLUSIONS In this study, the redox potential of organic batteries was controlled by varying the LiClO4 and crown ethers concentrations in the electrolyte system. The addition of the crown ethers, particularly 15C5, induced a significant up-shift of the redox potential further than that induced by 12C4; conversely, increasing the lithium salt concentration led to a down-shift of Emid. The change in the electrochemical potential was well-explained by thermodynamics and DFT calculations. The results of the DFT calculations were consistent with the experimental results. By simply adding 1 equivalent of 12C4 to the 0.25 M LiClO4 electrolyte, the capacity retention was improved up to 21%. Appropriate incorporation of the crown ethers into the electrolyte system modified the redox potential and afforded relatively high energy densities and excellent cycle performances in the lithium organic batteries.

AUTHOR INFORMATION Corresponding Author *Email: [email protected]; tel: +886-7-525-3951. Notes The authors declare no competing financial interest. ACKNOWLEDGMENTS The authors gratefully acknowledge support of the Ministry of Science and Technology in Taiwan through Grant MOST 106-2113-M-110-004.



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Crown Ethers as Electrolyte Additives to Modulate Electrochemical

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