Synthesis of Li₇P₃S₁₁ solid electrolyte in ethyl propionate medium for all-solid-state Li-ion battery

Abstract

In this study, we report a facile wet chemical synthesis method for preparing Li₇P₃S₁₁ solid-state electrolyte using ethyl propionate as the solvent. The structures of the samples were investigated using X-ray diffraction, Raman spectroscopy, and ³¹P nuclear magnetic resonance. The electrochemical properties of the samples were evaluated using cyclic voltammetry, direct current, and alternating-current electrochemical impedance measurements. The Li₇P₃S₁₁ solid electrolyte (SE) exhibited a Li⁺ conductivity of 1.5 × 10⁻³ S cm⁻¹ at 25 °C. The electrochemical measurements confirmed that the synthesized electrolyte is a single Li⁺ conductor with stability up to 5.0 V vs. Li⁺/Li. Furthermore, a solid-state battery (SSB) cell incorporating LiNi_1/3Mn_1/3Co_1/3O₂ and the synthesized Li₇P₃S₁₁ maintained stability for up to 50 cycles, demonstrating the durability of Li₇P₃S₁₁ in high-voltage SSB applications. These results indicate that Li₇P₃S₁₁ SE is a promising candidate for rechargeable solid-state Li-ion batteries.

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1. Introduction

Solid-state batteries (SSBs) are a class of batteries that use solid electrolytes (SEs) instead of the liquid or gel electrolytes employed in conventional lithium-ion batteries.¹ SSBs offer significant advantages over conventional batteries, including higher energy density and improved safety, making them a promising technology for future energy storage, particularly in electric vehicles. SSBs can store more energy within the same volume or weight than conventional batteries, enabling longer ranges for electric vehicles.² The use of SEs, which are more stable and less prone to leakage or fire than liquid electrolytes, significantly improves safety.

Sulfide solid electrolytes are considered as promising materials for SSBs because of their high ionic conductivity, mechanical flexibility, and high energy density.^3,4 Among them, Li₇P₃S₁₁ SE has been extensively studied for its ionic conductivity ranging from 10⁻⁴ to 10⁻² S cm⁻¹ at room temperature.⁵ Li₇P₃S₁₁ synthesized via solid-state reactions has achieved a maximum ionic conductivity of 1.7 × 10⁻² S cm⁻¹ at room temperature.⁶ However, solid-state reactions are energy-intensive, which limits their scalability for mass production. In contrast, liquid-phase synthesis offers a more versatile approach for the scalable production and preparation of electrode composite materials for all SSBs.⁷

Liquid-phase synthesis of Li₇P₃S₁₁ has recently been reported, leveraging its compatibility with composite electrode preparation.^8,9 Among the solvents used in Li₇P₃S₁₁, acetonitrile (ACN) is the most widely used, as it enables SEs with ionic conductivity up to 9.7 × 10⁻⁴ S cm⁻¹ at 25 °C.¹⁰ In CAN, Li₂S reacts with P₂S₅ to form Li₃PS₄ and Li₄P₂S₇ precipitates in a 1 : 1 molar ratio.¹¹ After solvent removal, the residues Li₄P₂S₇·ACN and Li₃PS₄·CAN, yielded Li₇P₃S₁₁ upon heat treatment at elevated temperatures as ACN was removed. Recently, ethyl acetate (EA) was used to synthesize Li₇P₃S₁₁, yielding an ionic conductivity of 1.05 × 10⁻³ S cm⁻¹ at 25 °C.¹² Unlike in CAN, Li₂S and P₂S₅ completely dissolved in EA to form Li₇P₃S₁₁. Other solvents, such as 1,2-dimethoxyethane and tetrahydropyran, have also been used for the synthesis of Li₇P₃S₁₁ SE.^13,14 Ethyl propionate (EP) has previously been used to synthesize β-Li₃PS₄, Li₇P₂S₈I, and Li₃PO₄-doped Li₃PS₄, but not Li₇P₃S₁₁.^15–18

In this study, Li₇P₃S₁₁ SE was synthesized using EP to facilitate the reaction between Li₂S and P₂S₅. Li₂S and P₂S₅ (70 : 30 molar ratio) did not fully dissolve in EP, resulting in the formation of a white precipitate and a yellowish solution. After solvent removal, the residue was heated at 250 °C for 1 h to yield Li₇P₃S₁₁ SE. The prepared Li₇P₃S₁₁ SE exhibited an ionic conductivity of 1.5 × 10⁻³ S cm⁻¹ at 25 °C, comparable to previously reported values. A solid-state half-cell using Li₃InCl₆-coated LiNi_1/3Mn_1/3Co_1/3O₂ as the active material and Li₇P₃S₁₁ demonstrated stable cycling performance, indicating the compatibility of the synthesized Li₇P₃S₁₁ SE with high-voltage materials.

2. Experimental

2.1 Chemicals

Li₂S (99.9%, Macklin), P₂S₅ (99%, Macklin), LiCl (99.9%, Macklin), InCl₃ (99.99%, Macklin), and super-dehydrated EP (Aldrich) were used as received without further treatment.

2.2 Liquid-phase synthesis of Li₇P₃S₁₁

A mixture of 2.0 g of Li₂S and P₂S₅ (7 : 3 molar ratio) was weighed and placed in a three-necked flask with 40 ml of EP. The mixture was stirred at 300 rpm and 50 °C for 24 h, yielding a suspension containing a white precipitate and a yellowish solution (Fig. S1). The solvent was evaporated from the suspension under reduced pressure at room temperature to obtain a residue (S-RT). The residue was then ground using an agate mortar and heat-treated in an Ar atmosphere for 1 h at 190 °C (S-190), followed by 1 h at 230 °C, 240 °C and 250 °C to obtain Li₇P₃S₁₁ SE (denoted as S-230, S-240, and S-250, respectively).

2.3 Li₃InCl₆-modified LiNi_1/3Mn_1/3Co_1/3O₂ preparation

Li₃InCl₆-modified LiNi_1/3Mn_1/3Co_1/3O₂ (NMC111, MTI) containing 2 wt% Li₃InCl₆ was prepared following a reported procedure.¹⁹ Li₃InCl₆ was synthesized according to a previously reported method.²⁰ Li₃InCl₆ was dissolved in anhydrous ethanol to form a uniformly transparent solution, into which NCM111 powder was added. After solvent evaporation at 80 °C, the precursor powder was further sintered at 400 °C under a nitrogen atmosphere to yield the composite powder (LIC@NCM111).

2.4 Structural characterization

The structure and morphology of the prepared samples were characterized using thermogravimetry-differential thermal analysis (TG-DTA, EVO II, Rigaku), X-ray diffraction (XRD, X8, Brucker), Raman spectroscopy (Horiba LabRam HR spectrometer, 532 nm), solid-state ³¹P nuclear magnetic resonance (³¹P NMR, Avance III 400, Bruker), scanning electron microscopy (SEM, S4800, Hitachi), and energy-dispersive X-ray spectroscopy (EDS, ULTIM MAX, Oxford Instrument).

The samples were prepared in an Ar-filled glove box and loaded into an airtight sample holder for characterization.

2.5 Electrochemical measurements

The electrical conductivity of a pellet prepared by uniaxially cold-pressing ∼200 mg of powder at 510 MPa was measured. Alternating-current impedance spectroscopy was performed using a potentiostat (PGSTAT302N, Autolab, Herisau, Switzerland) over a frequency range of 10 MHz–10 Hz.

DC conductivities were measured using blocking and nonblocking electrodes, in which stainless-steel rods and Li metal sheets were employed as the electrodes, respectively. The pelletized sample was prepared by cold-pressing S-250 powder at 510 MPa. For nonblocking electrodes, the Li metal sheets (∼8 mm diameter, 0.1 mm thickness) were attached to both faces of the pellet at room temperature. A 0.5 V DC was applied to the cells, and the current was measured using a potentiostat (SI 1287, Solatron) to determine the dominant mobile ions.

The electrochemical compatibility was tested using a Li|SE|Au cell (Au sputtered on a stainless-steel rod) at a scan rate of 5 mV s⁻¹ between 0.5 and 5 V with a potentiostat (SI 1287, Solatron).

The SSB half-cell was fabricated using the following process. The cathode composite was prepared by manually mixing LIC@NCM111 and S-250 in a 70 : 30 weight ratio using an agate mortar. A bilayer pellet (10 mm diameter) consisting of the electrode composite (12 mg) and S-250 (100 mg) was obtained by cold-pressing at 310 MPa. Indium (99.99%, Macklin) and Li foil were attached to opposite sides of the SE. The cell was assembled by sandwiching the pellet between stainless-steel rods and cycled in the constant-current mode at 0.1C between 3.70 and 2.40 V vs. Li–In.

3. Results and discussion

Fig. 1a shows the TG-DTA curves of the S-RT. The TG curve shows two distinct stages of the mass loss process: Stage 1 occurs from approximately 50 °C to 160 °C, and Stage 2 from 270 °C to 290 °C. The DTA curve displays several endothermic peaks between 30 °C and 150 °C, along with two exothermic peaks in the ranges of 150–200 °C and 270–290 °C. The mass loss in Stage 1 is attributed to the removal of crystallized solvent molecules. The P₂S₇⁴⁻ group in Li₇P₃S₁₁ is known to decompose at elevated temperatures, forming sulfur and P₂S₆⁴⁻.²¹ Therefore, the exothermic peak and accompanying mass loss in Stage 2 (270–290 °C) are ascribed to the decomposition of the P₂S₇⁴⁻ group in the sample. Fig. 1b shows the first-order derivatives of the TG and DTA curves. The first-order differential curves of TG-DTA show that the EP solvent removal process finishes at ∼170 °C. The d(DTA) curve illustrates an endothermic phenomenon between 230 °C and 250 °C without any mass loss, indicating the formation of Li₇P₃S₁₁. Notably, another endothermic phenomenon started at ∼250 °C, accompanied by mass loss, corresponding to the decomposition of P₂S₇⁴⁻ ions to P₂S₆⁴⁻. Based on these TG-DTA results, the experimental samples obtained after removal of the EP solvent at room temperature were heated at 190 °C for 1 h, followed by heat treatments at 230 °C, 240 °C, and 250 °C for 1 h to examine the formation of Li₇P₃S₁₁.

Fig. 1 (a) TG and DTA curves of S-RT; (b) TG and DTA first-order derivatives of S-RT.

Fig. S2 shows the XRD patterns of the precipitate obtained after decanting the yellow supernatant (Fig. S1), along with Li₂S for comparison. The precipitate pattern shows no residual Li₂S, indicating complete consumption of Li₂S in the reaction with P₂S₅ in the EP medium under specific conditions. Notably, the precipitate pattern resembled that of the Li₃PS₄ precursor synthesized from Li₂S and P₂S₅ in the EP medium.¹⁵ Fig. 2 shows the XRD patterns of Li₂S, P₂S₅, and the prepared samples. The Li₂S pattern exhibits peaks at 2θ ≈ 26° and 31°, while P₂S₅ appears nearly amorphous to XRD. The residue pattern after solvent removal at room temperature (S-RT) shows many peaks that do not correspond to known crystal structures. The absence of Li₂S in the S-RT pattern indicates that Li₂S has completely reacted with P₂S₅ in the EP medium. The pattern of the sample obtained after heating S-RT at 190 °C for 1 h (S-190) was nearly amorphous to XRD. Notably, it has been reported that the XRD pattern of 70Li₂S–30P₂S₅ recovered from EA after heating at 100 °C showed the crystal structure of β-Li₃PS₄.¹² The patterns of the samples obtained after heating S-190 at 230 °C and 240 °C for 1 h (S-230, S-240) exhibit peaks consistent with β-Li₃PS₄ and Li₇P₃S₁₁.^15,22 The patterns obtained after heating S-190 at 250 °C for 1 h (S-250) displays features of Li₇P₃S₁₁ along with some peaks of Li₄P₂S₆.^23–25 The formation of Li₇P₃S₁₁ in S-250 was consistent with the TG-DTA results, confirming that EP is an effective medium for synthesizing Li₇P₃S₁₁.

Fig. 2 XRD patterns of Li₂S, P₂S₅, and the prepared samples S-RT, S-190, S-230, S-240, and S-250.

Fig. S3 shows the Raman spectra of the precipitate obtained after decanting the yellow supernatant (Fig. S1), with the spectra of EP, Li₂S, and P₂S₅ included for comparison. The inset highlights an enlarged view of the white precipitate spectrum, which displays an intense peak at approximately 421 cm⁻¹, indicating the presence of PS₄³⁻ ion.²⁶ No Li₂S signal was detected in the spectrum of the white precipitate, confirming the complete consumption of Li₂S in the reaction with P₂S₅. Fig. 3 shows the Raman spectra of Li₂S, P₂S₅, and the prepared samples S-RT, S190, S-230, S-240, and S250. The inset on the left shows an enlarged view of the 3100–2800 cm⁻¹ region for the S-RT and S-190 spectra. The Li₂S spectrum exhibited a single intense peak at 370 cm⁻¹, while the P₂S₅ spectrum exhibited multiple peaks between 150 and 320 cm⁻¹. The S-RT spectrum shows peaks at 406, 420, 2940, and 2965 cm⁻¹, with no main peaks from Li₂S or P₂S₅ detected, indicating a complete reaction between Li₂S and P₂S₅. This observation is consistent with the XRD results. The peaks at 2940 and 2965 cm⁻¹ correspond to CH₃ group vibrations in EP.¹⁵ The peaks at 406 and 420 cm⁻¹ are attributed to the local structure units of P₂S₇⁴⁻ and PS₄³⁻ in Li₄P₂S₇ and Li₃PS₄, respectively.²⁶ The S-190 spectrum exhibits peaks at 402 and 416 cm⁻¹, indicating the presence of P₂S₇⁴⁻ and PS₄³⁻ in Li₄P₂S₇ and Li₃PS₄, respectively. The absence of CH₃ vibration peaks indicates the effective removal of EP after heating at 190 °C for 1 h. The positions of the PS₄³⁻ and P₂S₇⁴⁻ peaks in S-190 were shifted relative to those in the S-RT sample, indicating strong interactions between Li₃PS₄ and Li₄P₂S₇ with the EP solvent molecules. This observation suggests that EP removal from S-RT resulted in the formation of an amorphous structure in S-190, which is consistent with the XRD results. The spectra of S-230 and S-240 show the presence of P₂S₇⁴⁻ and PS₄³⁻ with peaks at 402 and 417 cm⁻¹, respectively. The spectrum of S-250 exhibits peaks at 417, 401, and 379 cm⁻¹, corresponding to the local structure units of PS₄³⁻ and P₂S₇⁴⁻ in Li₇P₃S₁₁, and P₂S₆⁴⁻ in Li₄P₂S₆, respectively.^8,26,27 These Raman results confirm the presence of Li₇P₃S₁₁ and Li₄P₂S₆ phases in S-250, which is consistent with the XRD measurements.

Fig. 3 Raman spectra of Li₂S, P₂S₅, and the prepared samples S-RT, S-190, S-230, S-240, and S-250. The figure on the right-hand side shows the enlargement in the range of 3050 to 2800 cm⁻¹ of the spectra of S-RT and S-190.

Powder XRD is effective for identifying crystalline phases but cannot characterize amorphous components. Therefore, ³¹P NMR was employed as a complementary technique to investigate the local structure of samples S-240 and S-250 (Fig. 4a and b, respectively). The S-240 spectrum exhibited peaks at the characteristic positions of the PS₄³⁻ and P₂S₇⁴⁻ groups in Li₇P₃S₁₁.^28–30 The deconvolution result revealed peaks at 86.9, 87.8, 88.9, and 91.7 ppm. The peak at 86.9 ppm corresponds to PS₄³⁻ unit in β-Li₃PS₄,³¹ while the peaks at 87.8 and 88.9 ppm are attributed to PS₄³⁻ and P₂S₇⁴⁻ units in Li₇P₃S₁₁, respectively.^28–30 The broad peak at 91.7 ppm is attributed to the P₂S₇⁴⁻ group in amorphous Li₄P₂S₇,³² indicating the presence of β-Li₃PS₄, Li₇P₃S₁₁, and amorphous Li₄P₂S₇ in S-240. The S-250 spectrum exhibits peaks at the characteristic positions of the PS₄³⁻ and P₂S₇⁴⁻ groups in Li₇P₃S₁₁ and P₂S₆⁴⁻ in Li₄P₂S₆.^28–30 The deconvolution result revealed peaks at 86.3, 87.8, 89.9, 91.8, 105.2, and 108.7 ppm. The peak at 86.9 ppm corresponds to the PS₄³⁻ of the structural unit in β-Li₃PS₄.³¹ The peaks at 87.8 and 89.9 ppm are attributed to PS₄³⁻ and P₂S₇⁴⁻ groups in Li₇P₃S₁₁, respectively, while the broad peak at 91.7 ppm arises from P₂S₇⁴⁻ group in amorphous Li₄P₂S₇. Notably, the area ratio between the P₂S₇⁴⁻ peak in amorphous Li₄P₂S₇ and Li₇P₃S₁₁ of S-250 is smaller than that of S-240, indicating the transformation of amorphous Li₄P₂S₇ into crystalline Li₇P₃S₁₁. The peaks at 105.2 and 108.7 ppm correspond to the P₂S₆⁴⁻ structural unit in Li₄P₂S₆ (ref. 32–34). The formation of Li₄P₂S₆ suggests that amorphous Li₄P₂S₇ was partially decomposed during the heat treatment at 250 °C. These findings are consistent with the XRD and Raman results.

Fig. 4 ³¹P NMR spectra of (a) S-240 and (b) S-250.

Fig. 5a shows the impedance spectra of S-250 over a frequency range of 10 Hz–10 MHz at temperatures ranging from 20 °C to 125 °C. At 20 °C, the spectrum displays a semicircle and a low-frequency tail, which is consistent with the Li⁺ blocking effect, indicating the ionic conductivity. The diameter of the semicircle decreases with increasing temperature and becomes almost negligible at 50 °C. The total impedance of the electrolyte pellet was determined from the intersection of the semicircle and the x-axis in the intermediate frequency region.³⁵ The temperature-dependent ionic conductivities of S-240 and S-250 were calculated from the total impedance values and are plotted in Fig. 5b. Notably, log₁₀ σ exhibits an almost linear dependence on the inverse temperature, following the Arrhenius equation σ = σ₀ exp(−E_a/(k_BT)). The ionic conductivities of S-240 and S-250 at 25 °C are 7.2 × 10⁻⁴ and 1.5 × 10⁻³ S cm⁻¹, respectively. These values are comparable to the reported values of conductivities for Li₇P₃S₁₁ synthesized via liquid-phase methods using ACN and EA, which are higher than those obtained with 1,2-dimethoxyethane (Table 1).^10,12,13 The calculated activation energies E_a for S-240 and S-250 are 0.18 and 0.27 eV, respectively. Previous studies have reported that a higher crystalline fraction of Li₇P₃S₁₁ enhances ionic conductivity and reduces E_a,³⁰ whereas the presence of Li₄P₂S₆ in Li₇P₃S₁₁ increases E_a.³⁶ The area fractions of amorphous Li₄P₂S₇, crystalline Li₇P₃S₁₁, Li₄P₂S₆, and β-Li₃PS₄ were calculated to evaluate the effect of Li₄P₂S₆ on the conductivity and activation energy of S-250 (Table 2). Therefore, the higher ionic conductivity of S-250 compared with S-240 is attributed to increased Li₇P₃S₁₁ crystallinity, while its higher E_a is attributed to the presence of Li₄P₂S₆.

Table 1 Summary of the liquid-phase processing data of Li₇P₃S₁₁ solid electrolytes

Solvent	Ionic conductivity (25 °C)/S cm^−1	Processing temperature/°C	Activation energy Ea/eV	Electrochemical window/V vs. Li^+/Li	Reference
ACN	9.7 × 10^−4	250	0.323	5.0	10
EA	1.05 × 10^−3	260	—	—	12
EP	1.5 × 10^−3	250	0.27	5.0	This work

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Fig. 5 (a) The impedance spectra of S-250 obtained at various temperatures from around room temperature (20 °C) to 125 °C; (b) temperature dependence of the ionic conductivity of S-240 and S-250.

Table 2 Area fraction of the NMR peaks for S-240 and S-250

Sample	Phase	Area (%)
S-240	Li4P2S7 amorphous	49.6
	Li7P3S11	21.8
	β-Li3PS4	28.6
S-250	Li4P2S7 amorphous	13.5
	Li7P3S11	69.4
	β-Li3PS4	5.5
	Li4P2S6	11.6

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Fig. 6a shows the variation of DC with time when a voltage of 0.5 V (DC) was applied to the S-250 sample. With nonblocking electrodes, a constant current was observed, whereas blocking electrodes initially exhibited polarization followed by a nearly constant current. The current with nonblocking electrodes was approximately four orders of magnitude higher than that with blocking electrodes, indicating that S-250 is a single-ion conductor with a lithium-ion transport number estimated to be higher than 0.999. Fig. 6b shows the cyclic voltammogram of S-250. The small inset shows an enlarged cyclic voltammetry curve in the range of 2.0–5.0 V vs. Li⁺/Li. The cathodic current, corresponding to the Li reduction, started at approximately 0 V, confirming the stability of the prepared SE against Li metal. Except for the cathodic anodic peaks, no additional peaks were observed within the scanned range, indicating the compatibility of the S-250 with Li metal up to 5 V vs. Li⁺/Li.

Fig. 6 (a) Change in DC with time when a voltage of 0.5 V (DC) was applied to the S-250 sample; (b) cyclic voltammogram of S-250 obtained at a scan rate of 5 mV s⁻¹ between 0.5 and 5 V.

Fig. 7a shows SEM-EDS images of the prepared LIC@NMC111 and S-250 mixture, illustrating that the SE particles are well-distributed among NMC111 particles. To assess the compatibility of S-250 with high-voltage SSBs, NMC111 was employed as the active material to evaluate cycling performance at a constant current of 0.1C in the voltage range of 2.4–3.7 V vs. Li–In at room temperature (Fig. 7b). The charge–discharge capacities of the 1st cycle were 122 and 101 mAh g_NMC⁻¹. In the 10th cycle, the capacities were 138 and 137 mAh g_NMC⁻¹, which were slightly higher than those of the 1st cycle. At the 30th cycle, the capacities were 120 and 119 mAh g_NMC⁻¹, respectively. The initial coulombic efficiency (CE) was approximately 83%. From 2nd cycle onwards, the CE was higher than 99%, indicating cell stability. Fig. 7c illustrates the charge–discharge capacities and CE of the prepared cell over 50 cycles, with the discharge capacity of the 50th-cycle nearly identical to that of the 1st cycle, further confirming the stability of the cell. These results demonstrate that the prepared S-250 SE is a promising candidate for SSBs.

Fig. 7 (a) SEM-EDS images of the mixture of LIC@NMC11 and S-250; (b) 1st, 10th, 30th, and 50th charge–discharge curves of the prepared solid-state cell cycling at 0.1C between 3.7 and 2.4 V vs. Li–In; (c) charge–discharge capacities and CE of the prepared cell for 50 cycles.

Fig. 8 shows the differential capacity curves (dQ/dV) of the NMC111 cathode at different cycles to analyze the evolution of the voltage profile. In the 1st cycle (Fig. 8a), distinct redox peaks at 3.81/3.71 V vs. Li⁺/Li were observed, attributed to Ni³⁺/Ni⁴⁺ oxidation/reduction associated with Li de-insertion/insertion.^37,38 In the 10th cycle (Fig. 8b), these peaks shifted slightly to 3.80/3.70 V vs. Li⁺/Li. Subsequently, the oxidation and reduction peaks shifted toward higher and lower voltage regions, respectively, indicating continuous capacity degradation (Fig. 8c). The redox peaks at 4.18/3.95 V vs. Li⁺/Li in the 1st cycle are attributed to Co⁴⁺/Co³⁺ oxidation/reduction.^37,38 These peaks shifted to the lower voltage region, which were again consistent with continuous capacity degradation. These results indicate that the electrode polarization increases with cycling. The redox peaks at about 3.84/3.55 V vs. Li⁺/Li observed in the 10th cycle, which was gradually shifted to lower voltage region, is related to the oxidation/reduction of Ni³⁺/Ni²⁺.^38,39 These peaks were not clearly visible in the 1st cycle, and their appearance correlates with the increase in capacity from the 1st to the 10th cycle.

Fig. 8 dQ/dV curves of the NMC111 cathode at different cycles: (a) 1st, (b) 10th, (c) 1st, 10th, 30th, and 50th.

4. Conclusion

This study demonstrates that Li₇P₃S₁₁ SEs can be synthesized using EP to promote the reaction between Li₂S and P₂S₅. The residue obtained after solvent removal at 190 °C is amorphous to XRD, but Raman spectroscopy confirms the presence of PS₄³⁻ and P₂S₇⁴⁻ groups. Heating this residue at 230 °C, 240 °C, and 250 °C leads to the formation of Li₇P₃S₁₁ SE. ³¹P NMR spectra results indicate that amorphous Li₄P₂S₇ reacts with β-Li₃PS₄ to yield Li₇P₃S₁₁. The prepared Li₇P₃S₁₁ SE exhibited an ionic conductivity of 1.5 × 10⁻³ S cm⁻¹ at 25 °C with an activation energy of 0.27 eV. The electrochemical measurements show that the prepared SE is stable up to 5.0 V vs. Li⁺/Li, demonstrating its suitability for SSB applications.

Author contributions

TAT: investigation, formal analysis, data curation, formal analysis, resources; NHHP: conceptualization, methodology, writing – original draft, visualization, supervision, project administration, writing – review & editing.

Conflicts of interest

The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper.

Data availability

The data itself is presented in the form of figures and tables in this manuscript.

Supplementary information is available. See DOI: https://doi.org/10.1039/d5ra05281e.

Acknowledgements

We acknowledge Ho Chi Minh City University of Technology (HCMUT), VNU-HCM for supporting this study. The authors thank Rio Jefferson (Enago; https://www.enago.com/vnuhcm/) for the English language review.

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