Cation-Directed Selective Polysulfide Stabilization in Alkali Metal

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Cation–Directed Selective Polysulfide Stabilization in Alkali Metal–Sulfur Batteries Qingli Zou, Zhuojian Liang, Guan-Ying Du, Chi-You Liu, Elise Y. Li, and Yi-Chun Lu J. Am. Chem. Soc., Just Accepted Manuscript • DOI: 10.1021/jacs.8b04536 • Publication Date (Web): 31 Jul 2018 Downloaded from http://pubs.acs.org on August 1, 2018

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Journal of the American Chemical Society

Cation–Directed Selective Polysulfide Stabilization in Alkali Metal–Sulfur Batteries Qingli Zou†, Zhuojian Liang†, Guan-Ying Du‡, Chi-You Liu ‡, Elise Y. Li ‡, and Yi-Chun Lu†* † Electrochemical Energy and Interfaces Laboratory, Department of Mechanical and Automation Engineering, The Chinese University of Hong Kong, Shatin, N.T. 999077, Hong Kong ‡ Department of Chemistry, National Taiwan Normal University, No. 88, Section 4, Tingchow Road, Taipei 116, Taiwan

ABSTRACT: Alkali-metal sulfur redox chemistry offers promising potential for high-energy-density energy storage. Fundamental understanding of alkali metal sulfur redox reactions is the prerequisite for rational designs of electrode and electrolyte. Here, we revealed a strong impact of alkali metal cation (Li+, Na+, K+ and Rb+) on polysulfide (PS) stability, redox reversibility, and solid product passivation. We employed operando UV-Vis spectroscopy to show that strongly negatively charged short-chain PS (e.g. S42-/S32-) is more stabilized in the larger cation (e.g. Rb+) than the smaller cation (e.g. Li+), which is attributed a stronger cationanion electrostatic interaction between Rb+ and S42-/S32- owing to its weaker solvation energy. In contrast, Li+ is much more strongly solvated by solvent and thus exhibits a weaker electrostatic interaction with S42-/S32-. The stabilization of short-chain PS in K+-, Rb+- sulfur cells promotes the reduction of long-chain PS to short-chain PS leading to high discharge potential. However, it discourages the oxidation of short-chain PS to long-chain PS leading to poor charge reversibility. Our work directly probes alkali metal-sulfur redox chemistry in operando and provides critical insights into alkali metal sulfur reaction mechanism.

INTRODUCTION Developing redox-active chemistry beyond lithium-based materials is critical to ensure long-term sustainability in energy storage. Elemental sulfur offers one of the highest gravimetric capacity among all electrode materials reported to date with large earth-abundance and relatively low cost.1-8 Unlike the most of the lithium-ion (Li-ion) intercalation materials whose reactions are based on one electron per transition metal ion (typically less than 300 mAh/g)9,10, reactions between elemental sulfur and metal sulfide stores two electron per sulfur atom yielding 1675 mAh/gsulfur.3,7,11,12 Much attention has been focused on lithium−sulfur (Li−S) batteries, which operate based on the formation of lithium sulfide (Li2S) (2Li + S → Li2S, E0 = 2.15 V), yielding a theoretical specific energy of 2600 Wh/kg4,11. Recently, other metal–sulfur batteries including roomtemperature sodium sulfur battery 13-25 ,potassium sulfur battery26,27, magnesium sulfur battery28-30, and calcium sulfur battery31 are emerging as alternatives to the Li-S batteries considering their high earth abundance. The development of metal sulfur batteries is limited by several problems including irreversible loss of active material, low coulombic efficiency, rapid capacity fading, limited rate capability, and high selfdischarge.1,3,4,7,12,32-34 These challenges are deeply related to the reactivity/stability of the soluble polysulfide intermediates and their subsequent side reactions with the metal anode (crossover). To tackle these challenges, mechanistic studies35-40 and material designs on sulfur cathodes41-46 have been intensively conducted for Li-S battery system.1,7,32,47 The stability and reactivity of polysulfide, the reaction pathways, and the influence of solvent molecules on these key factors have been studied in the Li+ containing electrolytes.33,34,37,39,48,49 Going be-

yond Li-S batteries, understanding the differences and correlations between the Li-S and other alkali metal-sulfur redox reactions is the key to leverage the knowledge in Li-S batteries and benefit other metal-sulfur systems. Alkali metal-sulfur batteries behave drastically different from each other and the underlying mechanism governing the differences are not well-understood.10,11,13,15,20,21,23,47 For instance, we compare representative voltage profiles of Li-33, Na-25, and K-26 sulfur batteries reported in the literature (all in tetraethylene glycol dimethyl ether, TEGDME), as shown in Figure S1. Clearly, the voltage pattern, achievable discharge capacity, coulombic efficiency, and polarization are extremely different. 25,26,33 However, these differences are originated from combinatorial effects from cathode and the anode (the two-electrode cell). Therefore, it is critical to systematically study the influence of cation on the sulfur redox reaction kinetics, reaction intermediates and reaction products excluding the contribution from the metal anodes. In this work, we investigate the influence of alkali cation on the sulfur redox reactions via operando UV–vis spectroscopy, rotating-ring disk voltammetry (RRDE) and a specially designed three-electrode sulfur cell (free from alkali metal’s interference). We select dimethyl sulfoxide (DMSO) as the model solvent as it offers reasonable solubility for alkali metal salt from LiClO4 to RbClO4. These techniques offer welldefined model systems for systematic evaluation of the differences and correlations between the Li-S and other alkali-sulfur redox reactions. We show that strongly negatively charged short-chain PS (e.g. S42-/S32-) is more stabilized in the larger cation (e.g. Rb+) than the smaller cation (e.g. Li+), which is attributed a stronger cation-anion electrostatic interaction between Rb+ and S42-/S32- owing to its weaker solvation energy. In contrast, Li+ is much more strongly solvated by solvent and

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thus exhibits a weaker electrostatic interaction with S42-/S32-. The difference in anion-cation electrostatic interaction of the solvated cations affects the stability of various polysulfide, leading to asymmetric of discharge and charge battery processes. The influences of cation on the polysulfide stability, redox reversibility and final product formation/passivation will be discussed and correlated to battery capacity and reversibility.

EXPERIMENTAL METHODS Operando UV-Vis spectroscopy cell assembling. The 1.0 mm micro cuvette was selected as cell body. The gold mesh was employed as a working electrode. The counter electrode was lithium metal foil. The Ag/Ag+ reference electrode (ALS; Japan) consisted of a glass tube filled with 0.01 M AgNO3, 0.1 M tetrabutylammonium perchlorate (TBAP) in acetonitrile. 2.0 mM S8 (Sigma Aldrich, >99.5%) were dissolved in pure DMSO (Sigma Aldrich, 99.7%) and stirred overnight. And 0.2 M lithium perchlorate (LiClO4) (Sigma Aldrich, 99.95%), sodium perchlorate (NaClO4)(Sigma Aldrich, 98.0%), potassium perchlorate(KClO4)(Sigma Aldrich, 99% ), and Rubidium perchlorate (RbClO4)(Alfa Aesar, 99.5%) were separately dissolved in the as prepared DMSO solution. The cell was assembled in an Argonfilled glove box (H2O < 1.0 ppm, O2 < 1.0 ppm, Etelux, China). The CV scans were obtained at a scanning rate of 1.0 mV/s. Operando UV-Vis spectroscopy measurement. The UV-Vis spectroscopic (SEC2000; ALS; Japan) was switched on at least 30 min before the test to have a stabilized light source. The reference was measured every time before and after the test. The UV-Vis measurement was started at the same time with the electrochemical test and the data was collected every 0.5 second. The UV-Vis cell was saturated with pure Argon gas (N5.0, HKO, Hong Kong) during the entire measurement. Preparation of polysulfide. Polysulfide samples with nominal composition “Li2S6”, “Li2S4” are prepared following the method developed by Rauh et al.50 Briefly, lithium sulfide (Sigma Aldrich, 99.98%) and sulfur were mixed with magnetic stirring at room temperature for 24 hours in various solvents in an Ar-filled glovebox to yield “Li2S6”, “Li2S4” with a concentration of 2mM following: Li2S + (n-1)/S8 → Li2Sn where n=4, 6. In order to evaluate the influence of alkali metal cations on the stability of polysulfide species, 0.2M alkali metal salt, including LiClO4, NaClO4, KClO4, and RbClO4, were dissolved in the as prepared polysulfide solution separately. RRDE Measurement. The nonaqueous RRDE configuration used in this study was adopted from that reported by Herranz et al.51. Briefly, the working electrode consisted of a PTFE embedded glassy carbon disk of 4.0 mm in diameter surrounded by a glassy carbon ring with an internal diameter of 5.0 mm and an external diameter of 7.0 mm (ALS; Japan). The reference electrode was assembled in the glovebox at least 30 min before putting together the rest of the electrochemical cell. 4 mM S8 was dissolved in pure and stirred overnight. And 0.2 M LiClO4, NaClO4, KClO4, and RbClO4 were separately dissolved in the as prepared DMSO solution. The RRDE working electrode, the Pt wire counter electrode, and the reference electrode were assembled inside the glovebox. Prior to the RRDE measurements.

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Measurements of Diffusion Coefficient of Elemental Sulfur. Diffusion coefficient of polysulfide in DMSO with alkali metal salt from LiClO4 to RbClO4 are determined using a method established in our previous studies40. Briefly, the diffusion coefficients of polysulfide is determined by fitting a relation between the inverse of the rotation speed and the transient time in potentialstepping experiments. Note that in all potential-stepping experiments, the disk potential was stepped from a defined potential that is close to OCV (no reaction happening) but not too far away from where it is stepping to. The diffusion coefficient of polysulfide is measured with a constant ring voltage that consume polysulfide on the ring (i.e., 0.5 VAg in present of LiClO4) and a stepping disk voltage that generate polysulfide on the disk (i.e., - 2.2 VAg in present of LiClO4). As soon as the stepping disk voltage is applied, the ring will oxidize the polysulfide sent from the disk. Metal-S three electrode catholyte cell Assembling. 4.0 mM S8 was first dissolved in DMSO at 50 ° C and stirred for 12 hours. Subsequently, 0.5 M LiClO4, NaClO4, KClO4, and 0.4M RbClO4 was then added in the sulfur containing electrolyte. The catholyte was then cool down to room temperature for more than 24 hours prior to use. Pt wire (ALS; Japan) was placed on the stainless steel negative cell case. Then 0.3 ml DMSO electrolyte with 0.5 M alkali metal perchlorate was added onto the Pt wire. One piece of glass fiber (QMA, Φ16, Whatman) was placed onto the Pt wire followed by pieces of carbon paper (HCP010N, Shanghai Hesen Electric Co. Ltd., China) as both current collector and inter layer to eliminate the shuttle of polysulfide species. And 10 µl catholyte (4 mM S8 –0.5 M alkali metal perchlorate in DMSO) was added. The Ag/Ag+ reference was placed between the cathode and anode. The three electrode metal-S catholyte cell was assembled in the Argon-filled glove box. Computation Density functional theory (DFT) calculations are performed by the Gaussian 16 program, details are listed in supporting information.

RESULTS AND DISCUSSION Cation’s Influence on Sulfur Redox Potentials. We first examine the sulfur reduction and oxidation reactions by cyclic voltammetry on a stagnant GC electrode in 4 mM S8 in DMSO with alkali ions containing Li+, Na+, K+ and Rb+ (Figure 1). The first reduction and oxidation peak centers at ≈ -0.8 VAg and ≈ -0.62 VAg , respectively, which is independent of cation. This redox reaction could be attributed to52-59 S8+ 2e- ↔ S82-

(eq 1)

The overlapping reduction/oxidation potentials for all cations in the first redox pair (eq 1) suggests that the stability/energetics of S8 and S82- is insensitive to cations in DMSO. Interestingly, strong cation influence is observed on the subsequent reduction wave37,40,48,52,53,55,59: S82-+ 2e- → 2S42-

(eq 2)

The second reduction and oxidation potential increases from Li+, Na+, K+ to Rb+. This suggests that the free energy of S42- is strongly affected by cations. We hypothesize that the Rb+ solvated S42- polysulfide is more energetically stable than the Li+ solvated S42- polysulfide, leading to a much reversible potential of the S82/S42- couple in the presence of Rb+. In addition, the oxidation cur-


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rent decreases from Li+ to Rb+, which indicates strong cationinfluence on the reversibility of low-chain-polysulfide to longchain-polysulfide.

Current (mA/cm2disk)

Figure 1. Comparison of sulfur CV curves in DMSO in present of 4mM S8 and 0.2M salt (LiClO4, NaClO4, KClO4, RbClO4, and TBAClO4) with scanning rate at 50mV/s. An Operando UV-Vis Approach to Probe Sulfur Reaction

1 0

To gain more insights into the reaction mechanism and the type of polysulfide involved in each redox step, we performed operando UV-Vis measurements along with the CV on gold mesh electrode (Figure 2). First, the redox potential and the peak features obtained on gold electrode in the UV-Vis cell are consistent with that observed on the GC electrode. Each stage was separated and highlighted with colors and labels (Figure 2).

Li Na K Rb TBA

-1

-2.0 -1.5 -1.0 -0.5 0.0 0.5 Potential [V vs. Ag/Ag+] +

Potential [V vs. Ag/Ag ] -2.0

-1.5

-1.0

-0.5

0.0

0.5 -2.0

Current (µA)

40

-1.5

-1.0

-0.5

0.0

0.5

a.

b.

0

Reduction

R1 -40

R2

With LiClO4

R3

2-

S4 2S3

2-

S6

2-

S3

S3

2-

S4

0.6

S6

2-

2-

S6 2S8

With RbClO4

R3

c.

S4 2S3

1.8 1.2

Reduction

R1

R2

-80 2-

d.

22-

S4

.-

S3

2-

S6 2S8

.-

S3

R1Li

R1Rb

0.0

Absorbance (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


e.

f.

1.8 1.2 0.6

R2Li

0.0

R2Rb

h.

g.

1.8 1.2 0.6

R3Li

R3Rb

0.0 300

450

600

750

300

450

600

750

Wavelength (nm)

Figure 2. Operando CV curve of sulfur in DMSO with (a) 0.2M LiClO4 and (b) 0.2M RbClO4. Operando UV-vis spectra of each reduction steps (colors from dark to light in each figure represent the bands’ evolution over time) in present of (c, e, g) LiClO4 and (d, f, h) RbClO4. During the first reduction wave, the UV absorbance with both Li and Rb salt shows continuous increase at S3•– (617 nm), S82–(492 nm), and S62–(475, 350 nm). This indicates that similar reaction intermediates are produced in the presence of Li+ and Rb+. The formation of S82–, S62– and S3•– can be attributed to the following steps37,40,52-55,59:

bands (R2 and R3). The first part (R2, Li salt at ≈ -1.3 VAg, Rb salt at ≈ -1.1 VAg ), shows an increase in S3•– (617 nm) accompanied by a decrease in S82– (492 nm) and an slightly increase at 325 nm (with Li and Rb salt) and 420nm (only with Rb salt), which can be assigned to S42–. According to pervious reports37,40,48,52,53,55 these cloud be attributed to the following reactions:

S8+ 2e- ↔ S82-

(eq 1)

S82-+ 2e- → 2S42-

(eq 2)

S82- → S62-+ 1/4 S8

(eq 3)

S42-+ 1/4 S8 ↔ 2S3•-

(eq 5)

S62- → 2S3•-

(eq 4)

The second reduction wave of sulfur with Li and Rb salts can be separated in to two steps according to the changes in the UV

During the second part (R3) with both Li and Rb salt, the S3•– (617 nm) decreased sharply with an increase in S32– (270 nm, 340nm). These suggest a reduction reaction of S3•– to form S32– (eq 6)



S3•-+ e- → S32-

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(eq 6)

As same polysulfide evolution processes were separately observed in present of Li+ and Rb+, similar reaction steps are suggested during sulfur reduction. To further evaluate the physical origins responsible for the distinct CV responses observed with various cations, the evolution of the UV absorbance of S3•– (617 nm) and S82– (492 nm), which are the main intermediate product, during the CV scans are compared in Figure 3. The changes of absorbance values are related to the changes in polysulfide concentration during negative-going scan. First, upon negative-going scan during R1 (-0.78 to -1.34 VAg), the absorbance of S82– and S3•– in both Li+ to Rb+ start increasing at around -0.85 VAg with almost same increasing rate. This is in consistent with the overlapped first reduction wave shown in CV scan (Figure 1). S8+ 2e- ↔ S82-

(eq 1)

S82- → S62-+ 1/4 S8

(eq 3)

S62- → 2S3•-

(eq 4)

Second, upon further reduction to R2, the UV absorbance of S82starts to decrease owing to electrochemical reduction of S82- for all cations: S82-+ 2e- → 2S42-

(eq 2)

Interestingly, the decrease of S82- absorbance occurs at the highest potential in Rb+ (-1.27 VAg) followed by K+ (-1.32 VAg), Na+ (1.35 VAg), and Li+ (-1.42 VAg). This directly supports that the high reduction potential of the second reduction wave observed for Rb+ is attributed to S82-+ 2e- → 2S42-

(eq 2)

Meanwhile, the potential at which S3•– starts to increases is the highest for Rb+ followed by K+ ≈ Na+ and Li+ (Figure 3b), which is consistent with S42-+ 1/4 S8 ↔2S3•–

(eq 5)

Third, when the potential further reduces (R3), the intensity of S3•– reduces (indicated with dash line in Figure 3b) as a result of electrochemical reduction of S3•–: S3•–+ e- → S32-

(eq 6)

The potential at which S3•– starts to reduce is the highest for Rb+ followed by K+, Na+ and Li+, suggesting that S32– (strongly negatively charged) is more stable in Rb+ than in Li+. In short, detailed comparison of the evolution of the UV absorbance upon reduction suggests that the presence of cation Rb+ promote the electrochemical reactions: S82-+ 2e- → 2S42-

(eq 2)

S3•–+ e- → S32-

(eq 6)

Figure 3. Evolution of UV-vis absorbance at (a) 492 nm (S82-) and (b) 617 nm (S3•–) during negative-going scan in all electrolytes. These experimental observations can be explained by the different strength of cation-anion electrostatic interaction of different cations. The originally larger cation Rb+ is much more weakly solvated by DMSO than the smaller Li+ due to lower solvation energy60, which is supported by our DFT calculation (Table 1). Due to weak solvation energy of Rb+, its electrostatic interaction with strongly negatively charged polysulfide (S42-/S32-) is stronger than Li+. By contrast, Li+ is strongly solvated by DMSO and its electrostatic interaction with S42-/S32- is weaker (similar to TBA+). This is supported by the observation that the CV response in Li+ electrolyte resembles to that in TBA+ (Figure 1). This explanation is supported by a spectroscopy study by J. Gill reporting that the extent of ion-pairing is greatest with the largest cations and decreases markedly with decreasing ionic radius: Cs+ > Rb+ > K+ > Na+ > Li+.60 The low ion-pairing ability of Li+ to anions is attributed to the high solvation energy of Li+.60 To further investigate how cation affect the stability of various polysulfides, we perform UV–vis measurements on a series of chemically synthesized polysulfides (Figure 4).


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.-

2-

S3

S4

2.4

a.

2-

S6

1.8

Absorbance(a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


2-

'S6 '

2-

S6

1.2

2-

S4

0.6 0.0 Li Na K Rb

2.4 1.8

b. 2-

'S4 '

1.2 0.6 0.0 300

450

600

Wavelength(nm)

750

Figure 4. UV-visible spectra of 2mM chemical synthesized (a) S62- and (b) S42- solutions in present of 0.2M salt from LiClO4 to RbClO4.

Table 1. Reference data of radius and solvation numbers of cations in DMSO solvent61-63, and the computed average binding energies between different cations (Eb) and n explicit DMSO molecules as well as the theoretically estimated binding numbers (nb). Eb (eV/molecule) Cation

Ionic radius (pm)

Apparent Stoke’s radius (pm)

Solvation numbers n=1

n=2

n=3

n=4

Estimated binding numbers (nb)

Li+

60

318

3.3

0.67

0.60

0.46

0.34

4

Na+

96

303

3.1

0.57

0.46

0.32

0.23

3

K+

133

291

2.8

0.37

0.26

--

--

1

Rb+

147

259

2.3

-

-

TBA+

494

374

-

-

-

Nominal polysulfides S62– and S42– (2 mM) were synthesized using in DMSO by mixing corresponding amount of Li2S and S850. Subsequently, 0.2 M of LiClO4, NaClO4, KClO4 and RbClO4 were added in the chemically prepared polysulfide. All the UV– vis spectra show a large absorption peak at 617 nm, which is attribute to S3•–37,48,53,64-67. The formation of S3•– from S62- and S42could be achieved via disproportionation reactions specified earlier such as

S42-+ 1/4 S8 ↔2S3•–

(eq 5)

Interestingly, the absorbance of the S3•– significantly decreased from Li+ to Rb+, indicating an increasing stability of S42– from Li+ to Rb+. The evolution of polysulfide ions during the positivegoing scan of sulfur CV with Li salt or Rb salt were monitored by UV–vis spectroscopy, as shown in figure 5.



Page 6 of 11

+

Potential [V vs. Ag/Ag ] -2.0

Current (µA)

40

-1.5

-1.0

Oxidation

-0.5

O4

0.5 -2.0

a.

O5

-1.0

-0.5

O4

Oxidation

0.0

O5

O2 O1

With LiClO4 2-

1.8 2-

0.5

b.

O3

O1

-40

1.2

-1.5

O3

O2

0

0.0

-80

With RbClO4 2-

c.

S4 2S3 2S6

S3

2-

S4

d.

S4 2S3 2S6

2-

S3

2-

0.6

S6 2S8

2-

2-

S4

.-

S3

S6 2S8

O1Li

.-

S3

O1Rb

0.0

e.

f.

1.8 1.2 0.6

O2Rb

O2Li 0.0

Absorbance (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

g.

h.

1.8 1.2 0.6

O3Li

O3Rb

0.0

i.

j.

1.8 1.2 0.6

O4Rb

O4Li 0.0

k.

l.

1.8 1.2 0.6 O5Rb

O5Li 0.0 300

450

600

750

300

450

600

750

Wavelength (nm)

Figure 5. Operando CV curve of sulfur in DMSO with (a) 0.2M LiClO4 and (b) 0.2M RbClO4. Operando UV-vis spectra of each oxidation reaction steps (Color from dark to light in each figure represent the bands’ evolution over time) in present of (c, e, g, i) LiClO4 and (d, f, h, j) RbClO4. During oxidation, the type of polysulfide observed were similar in both electrolytes, but appeared at different potentials. During initial oxidation (O1), UV spectra in Li salt showed obvious increase of S3•– at 617nm, which is in reverse with the prior reduction process. Therefore, it could be attribute to the oxidation reaction of S32– form S3•– S32- → S3•–+ e-

(eq 7)

In addition, S42- increases during O1, which could be resulted from a reverse chemical equilibrium of eq 5 owing to increase in S3•– concentration S42-+ 1/4 S8 ↔2S3•–

(eq 5)

In contrast, UV spectra in Rb salt showed a minor reduction in S3•–, which could be considered as the unfinished process of R3. As the oxidation continues, the S3•– increased as the oxidation proceed into O2 and O3.

During the second oxidation peak (O2) in the Li+-electrolyte, S82– and S62– continued to increase, which is similar in the Rb+electrolyte. This could be attributed to the oxidation of S42- to S82or to S62- following the reverse reaction of eq 2 and/or followed by eq 3 forming S622S42- → S82-+ 2e-

(eq 8)

S82- → S62-+ 1/4 S8

(eq 3)

During O3, S3•– in Li salt start to decrease with an increase in S82–, which can be assigned to the electrochemical oxidation of S3•–. The S3•– in Rb salt was still decreasing, which can be attributed to the unfinished process of O2 (time delay again). The S3•– in Rb salt started to decrease when entering the O4 stage. Finally, after the last oxidation peak (O4 and O5), all UV bands continuously decreased and returned to a similar level as the pristine state, indicating oxidation of polysulfides to elemental sulfur.


Page 7 of 11

To visualize such delay, we compare the evolution of the UV absorbance of S3•– (617 nm) during the positive-going CV scans (Figure 6). In general, the S3•– increases as we scan positively for all cations. Interestingly, the potential at which the S3•– starts to increase is the highest in Li+, followed by Na+, K+, and Rb+. This – suggests that the strongly negatively charged S32 is less stable in Li+ than in Rb+, which is consistent with the observation during negative-going scan (Figure 3).

Table 2. Diffusion coefficient of polysulfide ( Dpolysulfide, reduction product at the potential of -2.2V vs. Ag/Ag+) calculated by RRDE potential stepping experiment. Li

Na

Diffusion coefficient

4.889

(10-6 cm2/s)

± 0.054

K

Rb

5.157

5.774

6.151

± 0.051

± 0.053

± 0.101

Three Electrode Catholyte Cell

during positive-going scan in all electrolytes.

RRDE Approach: Diffusion Coefficient We further apply rotating-ring-disk electrode to reveal how the cation-complex affect the sulfur redox, especially the diffusion processes of the polysulfide. The diffusion coefficient was determined by stepping the disk potential from open-circuit voltage into the polysulfide generated potential, while holding the ring current at a positive potential where polysulfide is being oxidized continuously. From the time delay (Ts) between stepping the disk potential and observing a change in the ring current, certain polysulfide diffusion coefficient can be obtained via eq 9, where K is a proportionality constant K = 10.1 rpm-s . 1/3

Ts = K(υ/D) ω

−1

(eq 9)

The experimental Dpolysulfide values in present of salt from LiClO4 to RbClO4 were shown in Table 2 with error bar. We show that with the decreasing size of solvent-ion complex from Li+−(DMSO)n to Rb+−(DMSO)n (Table 1)60 the diffusion coefficient of polysulfide increase, which we believe explains the more sever polysulfide shuttling in Na-S batteries than in Li-S batteries24.

Figure 7 shows the galvanostatic discharge/charge profiles of alkali metal–sulfur catholyte cells in DMSO at 0.15C rate (2.5µA). All conditions in these cells are equivalent except the cations in the electrolyte. We note that the lack of flat region in Figure 7 (DMSO) is similar to reported Li-S cell in high-donicity solvents such as DMSO37,40 and DMF37. This can be attributed to slower reaction kinetics (due to high-donicity)40.

+

Figure 6. Evolution of UV-vis absorbance at 617 nm (S3•–)

To investigate the influence of cations on sulfur electrodes in the metal-sulfur batteries, it is critical to eliminate effects from the different metal anodes (e.g. overpotential in Li, Na and K metal). Most of Li, Na and K sulfur batteries shown in the literature are performed with pure Li, Na or K metal foil as anode1324,26,27,33,59,68 , which makes it difficult to separate the influence of anode from the sulfur electrode on the overall cell behaviors. We design a three-electrode closed-cell using 4 mM S8 together with the Pt counter electrode and Ag+/Ag reference electrode. The closed-cell configuration confines soluble polysulfide around the working electrode, which allows the full utilization of the soluble polysulfide when possible.

Voltage [V vs. Ag/Ag ]

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.0 Li Na K Rb

-0.4 -0.8 -1.2 -1.6 -2.0 0

400

800

1200

1600

S-Capacity (mAh/g) Figure 7. Galvanostatic discharge and charge profiles of three electrode S catholyte cell consist of 10µL 4mM S8 in a closed three-electrode cell. The mismatch between the measured capacity (1700 mAh/g) and the theoretical capacity of S (1675 mAh/g) in the Rb-S cell can be attributed to small variation in catholyte volume used in the three-electrode cell (10 µL). During discharge in the first plateau (≈-0.75 VAg), all the cations show the same discharge voltage with same capacity, which is consistent with our observation in the CV measurements that cation has very limited influence on the first reduction wave (Fig-

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ure 1). In contrast, the cation shows significant influence on the subsequent discharge profiles in terms of voltage and capacity. The second discharge plateau in Rb+ starts at the highest potential (around -1.0VAg), followed by K+, Na+, and Li+, which is in line with the onset potential observed in the CV (Figure 1). One critical information provided by the three-electrode closed-cell is the total achievable capacity. The achieved discharge capacity in Na and K are much smaller than that of Li, which is consistent with many literature reporting much smaller capacity of Na and K compared to Li-S batteries.16,18,23,25,26 This is attributed to the formation of different discharge products in Li, Na and K-S batteries.14,26,69 For instance, Ahn et al.14 and Chen et al.26, apply Xray diffraction and reported that the final product of Na-S and K-S battery is Na2S2 with Na2S3 and K2S3 instead of metal disulfide in the Li-S batteries. Based on the observed discharge capacity of the catholyte cells (Figure 7), Li2S is the dominant product in the Li-S cell to account for a discharge capacity of ~1600 mAh/g; a mixture of Na2S2/Na2S and K2S2/K2S could be the final products in the Na-S and K-S catholyte cells to account for a discharge capacity of ~1000 mAh/g. We note that the higher discharge capacity observed in our catholyte cell for Na-S and K-S cells compared to literature using solid sulfur electrodes14,26 could be attributed to higher degree of sulfur utilization and conversion since only dissolved sulfur (4 mM S8) was used as oppose to highsurface-area solid sulfur electrodes. Notably, the discharge capacity of the Rb-S cell is similar to that of the Li-S, approaching to the theoretical capacity of formation of metal disulfide (1675 mAh/g). We hypothesize that the high discharge capacity of the Rb-S cell could be related to the highly soluble nature of the Rb2S.To support this, we evaluated the dissolution enthalpy of M2S (∆Hdis), as well as the geometry parameters of M2S(DMSO) (Table S1). A general trend that the ∆Hdis decreases from Li to Rb can be observed, indicating that the Rb2S can be more soluble than other M2S in DMSO solution. We note that the absolute number of the dissolution enthalpy may depend on the chosen basis set, but all calculations confirm the same trend (Table S2). Therefore, we believe that the Rb2S, as well as other product formed in the Rb-S battery could be more soluble in DMSO than all other M2S/M2S2, which offer a plausible explanation to the high capacity of Rb2S in Figure 7. The influences of cation on polysulfide stabilities, sulfur reduction pathways and discharge products are summarized in Figure 8.

Page 8 of 11

Figure 8. Schematic illustration of the cation-directed sulfur reduction reactions.

During recharge, the charge capacity reduces from Li+ to Rb+ , resulting coulombic efficiency of 85%, 74%, 68%, 40% from Li+ to Rb+. Interestingly, the first charge plateau potential is quite similar in all cases and the major charge irreversibility in Rb+, K+ and Na+ cells occurred at the later stage at higher potentials. While many factors will lead to poor rechargeability, we believe that the poor charge behavior, especially for Rb+, could be attributed to the strongly stabilized discharge product (M2S/M2S2), and thus reduces the rechargeability. Further studies are on-going to investigate the origins responsible for the poor rechargeability of Rb-S system. Considering the drastic difference in polysulfide equilibrium induced by different cations, introducing mixed cations in sulfur cells could be a potential direction to manipulate the desired behaviors of the polysulfide in sulfur batteries.70 To investigate if our hypothesis is applicable in a more commonly used solvent in Li-S batteries such as 1,2-Dimethoxyethane (DME), we conducted the CV measurements of in DME (only Li+, Na+, TBA+ are conducted since K+ and Rb+ salts are not soluble in DME, Figure S2). As shown in Figure S3a, the CV patterns are largely similar with major differences in the reduction peak separation where TBA+ has the largest peak separation (580 mV), followed by Li+ (370 mV) and Na+ (290 mV).The large peak separation in TBA+ could indicate a strong stabilization of the longchain polysulfide, which delays its subsequent reduction to midor short- chain polysulfide40. Therefore, the larger peak separation observed in Li+ than in Na+ suggests that the DME-solvated Li+ has a weaker electrostatic interaction to polysulfides than DME-solvated Na+, which is consistent with the observations in DMSO. The galvanostatic voltage profiles of Li+ and Na+ in DME (Figure S3b) agree with that obtained in DMSO (Figure 7) that Li+ and Na+ cells exhibit similar first discharge voltage and the Li+ cell has a higher discharge capacity than Na+. Figure S3c shows the ex-situ UV-vis spectra of chemically prepared polysulfide with addition of 0.4 M of the corresponding cations. The spectra agree with the CV results that the nature of the polysulfide in TBA+ resemble that in Li+ rather than in Na+. These observations in DME and DMSO suggest that solvent environment and solvation energy of the cation should be considered when evaluating the electrostatic interaction between cation and anions. We believe that our findings could be applied to understand other types of solvents and cations (Mg2+, Ca2+) and their interactions with polysulfides.

CONCLUSION Our study directly probes alkali metal-sulfur redox chemistry in operando without interferences of metal anodes and reveals the critical role of cation on battery capacity and reversibility. We show that strongly negatively charged short-chain PS (S42-/S32-) is more stabilized in the larger cation (e.g. Rb+) than the smaller cation (e.g. Li+),which is attributed a stronger cation-anion electrostatic interaction between Rb+ and S42-/S32- owing to its weaker solvation energy. Such stabilization explains the asymmetric nature of discharge and charge profiles in K+- and Rb+- sulfur cells. Our study provides critical insights into alkali metal sulfur reaction mechanism and underlying mechanism governing their distinct electrochemical behaviors.

ASSOCIATED CONTENT


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Supporting Information. Include the computation detail. This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Author * Email: [email protected]

Notes The authors declare no competing financial interests.

ACKNOWLEDGMENT The work described in this paper was fully supported by a grant from the Research Grant Council of the Hong Kong Special Administrative Region, China (Project No. T23-601/17-R).

REFERENCES (1) Manthiram, A.; Fu, Y.; Chung, S.-H.; Zu, C.; Su, Y.-S. Chem. Rev. 2014, 114, 11751. (2) Xiao, J. Adv. Energy Mater. 2015, 5, 1501102. (3) Ji, X.; Nazar, L. F. J. Mater. Chem. 2010, 20, 9821. (4) Xu, R.; Lu, J.; Amine, K. Adv. Energy Mater. 2015, 5, 1500408. (5) Peng, H. J.; Huang, J. Q.; Cheng, X. B.; Zhang, Q. Adv. Energy Mater. 2017, 7, 1770141. (6) Borchardt, L.; Althues, H.; Kaskel, S. Current Opinion in Green and Sustainable Chem. 2017, 4, 64. (7) Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J. M. Nat. Mater. 2012, 11, 19. (8) Hassoun, J.; Scrosati, B. Adv. Mater. 2010, 22, 5198. (9) Whittingham, M. S. Science 1976, 192, 1126. (10) Whittingham, M. S. Chem. Rev. 2004, 104, 4271. (11) Ellis, B. L.; Lee, K. T.; Nazar, L. F. Chem. Mater. 2010, 22, 691. (12) Yin, Y.-X.; Xin, S.; Guo, Y.-G.; Wan, L.-J. Angew. Chem. Int. Ed. 2013, 52, 13186. (13) Lee, D.-J.; Park, J.-W.; Hasa, I.; Sun, Y.-K.; Scrosati, B.; Hassoun, J. J. Mater. Chem. A 2013, 1, 5256. (14) Ryu, H.; Kim, T.; Kim, K.; Ahn, J.-H.; Nam, T.; Wang, G.; Ahn, H.-J. J. Power Sources 2011, 196, 5186. (15) Carter, R.; Oakes, L.; Douglas, A.; Muralidharan, N.; Cohn, A. P.; Pint, C. L. Nano Lett. 2017, 17, 1863. (16) Bauer, I.; Kohl, M.; Althues, H.; Kaskel, S. Chem. Commun. 2014, 50, 3208. (17) Yu, X.; Manthiram, A. J. Phys. Chem. C 2014, 118, 22952. (18) Yu, X.; Manthiram, A. ChemElectroChem 2014, 1, 1275. (19) Yu, X.; Manthiram, A. J. Phys. Chem. Lett. 2014, 5, 1943. (20) Manthiram, A.; Yu, X. Small 2015, 11, 2108. (21) Yu, X.; Manthiram, A. Adv. Energy Mater. 2015, 5, 1500350. (22) Kohl, M.; Borrmann, F.; Althues, H.; Kaskel, S. Adv. Energy Mater. 2016, 6, 1502185. (23) Lu, Q.; Wang, X.; Cao, J.; Chen, C.; Chen, K.; Zhao, Z.; Niu, Z.; Chen, J. Energy Storage Mater. 2017, 8, 77.

(24) Wang, Y. X.; Zhang, B.; Lai, W.; Xu, Y.; Chou, S. L.; Liu, H. K.; Dou, S. X. Adv. Energy Mater. 2017, 7, 1602829. (25) Kim, I.; Park, J.-Y.; Kim, C. H.; Park, J.-W.; Ahn, J.-P.; Ahn, J.-H.; Kim, K.-W.; Ahn, H.-J. J. Power Sources 2016, 301, 332. (26) Zhao, Q.; Hu, Y.; Zhang, K.; Chen, J. Inorg. Chem. 2014, 53, 9000. (27) Lu, X.; Bowden, M. E.; Sprenkle, V. L.; Liu, J. Adv. Mater. 2015, 27, 5915. (28) Vinayan, B. P.; Zhao-Karger, Z.; Diemant, T.; Chakravadhanula, V. S.; Schwarzburger, N. I.; Cambaz, M. A.; Behm, R. J.; Kubel, C.; Fichtner, M. Nanoscale 2016, 8, 3296. (29) Yu, X.; Manthiram, A. ACS Energy Lett. 2016, 1, 431. (30) Zhang, Z.; Cui, Z.; Qiao, L.; Guan, J.; Xu, H.; Wang, X.; Hu, P.; Du, H.; Li, S.; Zhou, X.; Dong, S.; Liu, Z.; Cui, G.; Chen, L. Adv. Energy Mater. 2017, 7, 1602055. (31) See, K. A.; Gerbec, J. A.; Jun, Y. S.; Wudl, F.; Stucky, G. D.; Seshadri, R. Adv. Energy Mater. 2013, 3, 1056. (32) Gao, J.; Abruña, H. c. D. J. Phys. Chem. Lett. 2014, 5, 882. (33) Zhang, S.; Ueno, K.; Dokko, K.; Watanabe, M. Adv. Energy Mater. 2015, 5, 1500117. (34) Pang, Q.; Liang, X.; Kwok, C. Y.; Nazar, L. F. Nat. Energy 2016, 1, 16132. (35) Conder, J.; Bouchet, R.; Trabesinger, S.; Marino, C.; Gubler, L.; Villevieille, C. Nat. Energy 2017, 2, 17069. (36) Liang, X.; Kwok, C. Y.; Lodi‐Marzano, F.; Pang, Q.; Cuisinier, M.; Huang, H.; Hart, C. J.; Houtarde, D.; Kaup, K.; Sommer, H. Adv. Energy Mater. 2016, 6, 1501636. (37) Zou, Q.; Lu, Y.-C. J. Phys. Chem. Lett. 2016, 7, 1518. (38) Patel, M. U.; Demir‐Cakan, R.; Morcrette, M.; Tarascon, J. M.; Gaberscek, M.; Dominko, R. ChemSusChem 2013, 6, 1177. (39) Barchasz, C.; Molton, F.; Duboc, C.; Leprêtre, J.-C.; Patoux, S.; Alloin, F. Anal. Chem. 2012, 84, 3973. (40) Lu, Y.-C.; He, Q.; Gasteiger, H. A. J. Phys. Chem. C 2014, 118, 5733. (41) Ji, X.; Lee, K. T.; Nazar, L. F. Nat. Mater. 2009, 8, 500. (42) Jayaprakash, N.; Shen, J.; Moganty, S. S.; Corona, A.; Archer, L. A. Angew. Chem. 2011, 123, 6026. (43) Oschatz, M.; Borchardt, L.; Pinkert, K.; Thieme, S.; Lohe, M. R.; Hoffmann, C.; Benusch, M.; Wisser, F. M.; Ziegler, C.; Giebeler, L. Adv. Energy Mater. 2014, 4, 1300645. (44) Zhou, G.; Paek, E.; Hwang, G. S.; Manthiram, A. Nat. Commun. 2015, 6, 7760. (45) Liang, X.; Garsuch, A.; Nazar, L. F. Angew. Chem. Int. Ed. 2015, 54, 3907. (46) Yao, H.; Zheng, G.; Hsu, P.-C.; Kong, D.; Cha, J. J.; Li, W.; Seh, Z. W.; McDowell, M. T.; Yan, K.; Liang, Z. Nat. Commun. 2014, 5, 3943. (47) Akridge, J. R.; Mikhaylik, Y. V.; White, N. Solid State Ionics 2004, 175, 243. (48) Kim, B. S.; Park, S. M. J. Electrochem. Soc. 1993, 140, 115. (49) Zhu, W.; Paolella, A.; Kim, C. S.; Liu, D.; Feng, Z.; Gagnon, C.; Trottier, J.; Vijh, A.; Guerfi, A.; Mauger, A.; Julien, C. M.; Armand, M.; Zaghib, K. Sustainable Energy Fuels 2017, 1, 737. (50) Rauh, R. D.; Shuker, F. S.; Marston, J. M.; Brummer, S. B. J. Inorg. Nucl. Chem. 1977, 39, 1761.



Page 10 of 11

(51) Herranz, J.; Garsuch, A.; Gasteiger, H. A. J. Phys. Chem. C 2012, 116, 19084. (52) Bonnaterre, R.; Cauquis, G. J. Chem. Soc., Chem. Commun. 1972, 5, 293. (53) Martin, R. P.; Doub, W. H.; Roberts, J. L.; Sawyer, D. T. Inorg. Chem. 1973, 12, 1921. (54) Badoz-Lambling, J.; Bonnaterre, R.; Cauquis, G.; Delamar, M.; Demange, G. Electrochim. Acta 1976, 21, 119. (55) Fujinaga, T.; Kuwamoto, T.; Okazaki, S.; Hojo, M. Bull. Chem. Soc. Jpn. 1980, 53, 2851. (56) Baranski, A. S.; Fawcett, W. R.; Gilbert, C. M. Anal. Chem. 1985, 57, 166. (57) Leghié, P.; Lelieur, J. P.; Levillain, E. Electrochem. Commun. 2002, 4, 406. (58) Leghié, P.; Lelieur, J. P.; Levillain, E. Electrochem. Commun. 2002, 4, 628. (59) Cuisinier, M.; Hart, C.; Balasubramanian, M.; Garsuch, A.; Nazar, L. F. Adv. Energy Mater. 2015, 5, 1401801. (60) Gill, J. B. Pure Appl. Chem. 1981, 53, 1365. (61) Matsuura, N.; Umemoto, K.; Takeda, Y. Bull. Chem. Soc. Jpn. 1975, 48, 2253. (62) Paul, R.; Johar, S.; Banait, J.; Narula, S. J. Phys. Chem. 1976, 80, 351. (63) Yang, R. T. Adsorbents: fundamentals and applications; John Wiley & Sons, 2003. (64) Seel, F.; Guttler, H.-J.; Simon, G.; Wieckowski, A. Pure Appl. Chem. 1977, 49, 45. (65) Gaillard, F.; Levillain, E. J. Electroanal. Chem. 1995, 398, 77. (66) Kawase, A.; Shirai, S.; Yamoto, Y.; Arakawa, R.; Takata, T. Phys. Chem. Chem. Phys. 2014, 16, 9344. (67) Yu, X.; Manthiram, A. Phys. Chem. Chem. Phys. 2015, 17, 2127. (68) Yu, X.; Manthiram, A. Chem. Eur. J. 2015, 21, 4233. (69) Waluś, S.; Barchasz, C.; Colin, J.-F.; Martin, J.-F.; Elkaïm, E.; Leprêtre, J.-C.; Alloin, F. Chem. Commun. 2013, 49, 7899. (70) Gao, T.; Noked, M.; Pearse, A. J.; Gillette, E.; Fan, X.; Zhu, Y.; Luo, C.; Suo, L.; Schroeder, M. A.; Xu, K. J. Am. Chem. Soc. 2015, 137, 12388.


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Cation-Directed Selective Polysulfide Stabilization in Alkali Metal

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