Environmental Stability of Li₆PS₅Cl_0.5Br_0.5 Electrolyte During Lithium Battery Manufacturing and a Simplified Test Protocol

Abstract

In this study, we investigate the environmental stability of the sulfide-based argyrodite solid electrolyte Li₆PS₅Cl₀.₅Br₀.₅, a promising candidate for all-solid-state lithium batteries due to its high ionic conductivity and favorable mechanical properties. Despite its potential, the material’s sensitivity to ambient air humidity presents challenges for large-scale battery manufacturing. Moisture exposure leads to performance degradation and the release of toxic hydrogen sulfide (H₂S) gas, raising concerns for workplace safety. The objectives of this study are to validate the electrolyte synthesis process, evaluate the effects of air humidity exposure on its reactivity and ionic conductivity, and establish a standardized protocol for assessing environmental stability. We report a synthesis method based on ball milling and heat treatment that achieves an ionic conductivity of 2.11 mS/cm, along with a fundamental study incorporating modeling and formulation approaches to evaluate the electrolyte’s environmental stability. Furthermore, we introduce a simplified testing method for assessing environmental stability, which may serve as a benchmark protocol for the broader class of argyrodite solid electrolytes.

1. Introduction

As lithium-ion batteries take center stage as a key component in global decarbonization, novel modification of lithium-ion technology in the form of solid-state batteries has recently gained much attention due to their potential safety, energy density, and cycle life benefits. Solid electrolytes are at the core of a solid-state cell and therein lies the appeal of this innovative technology. Solid electrolytes offer enhanced safety over flammable liquid electrolytes, necessitating specific thermal management considerations, particularly when used in electric vehicles [1,2,3]. They also have potentially slower degradation which leads to a higher cycle life. Another prospective advantage of solid electrolytes is the opportunity that they can provide to increase energy density at the battery level, which would be brought about by their adoption in all-solid-state lithium batteries (ASSLBs) [1]. For electric vehicles utilizing ASSLBS, higher energy densities would mean longer mile ranges per charge [4,5,6].

In the three main categories of solid electrolytes—oxides, sulfides, and polymers—sulfides are among the most promising due to their desirable mechanical properties, particularly ease of processability, while maintaining high ionic conductivity [6,7,8,9,10,11,12]. With a Young’s modulus of about 20 GPa, sulfide solid electrolytes can be considered semi-brittle ceramics, sitting between oxides and organic polymers [8,13]. The importance of a solid electrolyte with an intermediate Young’s modulus is twofold: it is sufficiently pliant to allow for processing methods like pressurized stacking, and at the same time, it maintains enough strength to resist lithium dendrite formation [5,6,12].

Oxide electrolytes such as perovskite-type Li₀.₅La₀.₅TiO₃ (LLTO) and newly developed compositionally complex perovskite oxides (CCPOs) offer excellent electrochemical and thermal stability [14,15]. However, their typically high grain boundary resistance and brittleness can limit processability and integration [16,17]. While CCPOs such as (Li₀.₃₇₅Sr₀.₄₃₇₅) (Ta₀.₃₇₅Nb₀.₃₇₅Zr₀.₁₂₅Hf₀.₁₂₅)O₃ − δ have demonstrated improved total ionic conductivities via grain boundary engineering and compositional complexity, their conductivities still remain modest compared to sulfides, particularly in bulk [14].

Polymer electrolytes, on the other hand, are valued for their flexibility and safety [17,18]. Recent advances incorporating ionic liquids and polyamide nanofibers into polyethylene oxide matrices have enhanced flame retardancy and an extended cycle life, with room-temperature performance still constrained by limited ionic conductivity (~10⁻⁶ S/cm) [19].

These efforts underscore the tradeoff between mechanical compliance and conductivity performance. Sulfides offer a rare balance which combines moderate mechanical properties with high ionic conductivity. Furthermore, among inorganic solid electrolytes, sulfide-based systems are most desirable for delivering high ionic conductivity at room temperature. In particular, the lithium argyrodite family, in the form of Li₆PS₅X (X = Cl, Br, I), as first established by Deiseroth et al. [20], has very high Li⁺ mobility and the capability of producing ionic conductivities above 10⁻³ S/cm, which is comparable to those of liquid electrolytes [20,21,22].

Although synthesis methods may vary, such as milling or sintering in the solid or liquid phase, the bulk ionic conductivity results of sulfide argyrodite solid electrolytes are very promising. Investigating the effect of halides on the aforementioned argyrodite solid electrolytes, Rao et al. found that mechanical milling followed by annealing led to ionic conductivity in the order of 10⁻¹ mS/cm for Li₆PS₅Cl [23]. Boulineau et al. reported mechanochemical synthesis of Li₆PS₅Cl, obtaining a high ionic conductivity of 1.33 mS/cm [24]. For the same electrolyte composition but a different synthesis method, Wang et al. reported an ionic conductivity of 3.15 mS/cm by optimizing the sintering duration [20,25].

While there are many theoretical advantages to the adoption of sulfide solid electrolytes, there are still several limitations to their application in large-scale lithium-ion batteries. Most prominent of these are high interfacial resistance with electrodes [26,27] and instability in ambient air moisture and oxygen [28]. It should be noted that, in addition to humidity and oxygen, ambient air also contains CO₂, which may react with argyrodite electrolytes. While the softness of sulfide solid electrolytes allows for good contact with electrode materials, it has a low electrochemical stability window, leading to undesired oxidation and reduction reactions at the interface [9,29,30].

Most pertinent to this study is the chemical instability and hydrolysis of sulfides when exposed to ambient conditions. Examples of possible hydrolysis reactions reported in the literature which may occur with exposure to some molarity of H₂O can be found in Equations (1) and (2) [6,31]. Both reactions lead to the release of hydrogen sulfide (H₂S) gas, which is toxic.

$${\mathrm{Li}}_6{\mathrm{PS}}_5{\mathrm{Cl}}_{0.5}{\mathrm{Br}}_{0.5}+6{\mathrm{H}}_2\mathrm{O}\to 5{\mathrm{H}}_2\mathrm{S}+{\mathrm{Li}}_3{\mathrm{PO}}_4+{\displaystyle \frac{1}{2}}\mathrm{LiCl}+{\displaystyle \frac{1}{2}}\mathrm{LiBr}+2\mathrm{LiOH}$$

(1)

Li₆PS₅Cl + 4 H₂O → 4 H₂S + Li₃PO₄ + LiCl + Li₂S

(2)

In the first reaction, when LPSCB is exposed to ambient air containing humidity, oxygen, and CO₂, the potential formation of Li₂O, Li₂CO₃, Li₂S, and H₃PO₄ could be additionally anticipated alongside the products mentioned in Equation (1). While there are several variations in hydrolysis reactions proposed in the literature, our modeling is only dependent upon H₂S generation regardless of reaction pathway and subsequent byproducts [32]. As such, Equations (1) and (2) are exclusively illustrative, and our analysis is conducted assuming a generalized reaction illustrated in Equation (3).

Li₆PS₅Cl_0.5Br_0.5 + 6 H₂O → 5 H₂S + byproducts

(3)

Advancing the endeavor to stabilize sulfide electrolytes in ambient air, many studies report some improvements through the doping of elements such as oxygen and copper [21,33,34,35,36]. Jiang et al. reported improved air stability and high ionic conductivity of a copper bromide-doped sulfide argyrodite electrolyte with the composition Li_5.6Cu_0.2PS_4.8Br_1.2 [33]. Several studies have reported observations of decreased ionic conductivity with increasing air exposure, which is expected since hydrolysis reactions deteriorate the functional composition of the electrolyte [37,38]. For example, Jang et al. reported zinc oxide doping of Li₆PS₅Cl, which showed 41.7% ionic conductivity retention after 30 min [39]. Nonetheless, the reaction mechanisms behind these improvements require more detailed exploration, especially regarding the clarification of the foundational chemistry for improved air stability.

The thorough investigation of environmental reactivity is essential for the successful integration of sulfide solid electrolytes into the manufacturing processes of solid-state batteries [4,38,40]. The methodology of analyzing the reactivity of sulfide electrolytes in a controlled environment is most easily evaluated by using an H₂S sensor and a desiccator as a sealed container of fixed volume to hold the released H₂S gas, as utilized by previous researchers [31,41]. Some researchers also place a micro fan and a small container of water such as a crucible within the desiccator [31,34,41,42]. However, there is a discrepancy between studies as to the size of the desiccator and therefore the volume of air the electrolyte is exposed to. Furthermore, previous studies reported exposure of electrolyte amounts ranging from 50 to 100 mg to air at varying relative humidities with a broad range of 15 to 80% RH [42,43,44]. Although one study by Singer et al. in 2022 investigated the hydrolysis of Li₆PS₅Cl in the form of powder and separator sheets, whether the electrolyte should be tested in pellet or powder form mostly appears arbitrary throughout the literature, even though each sample formation has different surface areas in contact with the testing environment and will both be used in distinct stages of manufacturing [31]. The resulting exposure and H₂S generation data from these varying environments and sample formats would require normalization for any accurate comparative analysis of the electrolyte material between studies [45]. The many differences in experimental procedures expose the necessity for a formal and standardized method for evaluating the effect of hydrolysis on the air stability of sulfide solid-state electrolytes.

In this study, after optimizing the synthesis process of LPSCB electrolytes, synthesized by high-energy ball milling and heat treatment, and achieving a relatively high ionic conductivity, we investigate the effects of ambient air exposure on the electrolyte through comparisons of X-ray diffraction (XRD) and scanning electron microscopy (SEM) with energy-dispersive spectroscopy (EDS) characterization data, the evolution of H₂S gas, and ionic conductivity before and after air exposure. Data are fitted to appropriate models to understand the reaction rate and time constants that represent the environmental stability of LPSCB. Moreover, this study introduces an environmental testing protocol that could serve as a benchmark for future research in this field.

2. Materials and Methods

The LPSCB was synthesized using mechanical milling and subsequent heat treatment. Stoichiometric ratios of Li₂S (MSE Supplies, 99.9%), P₂S₅ (Sigma-Aldrich, 99%), LiCl (Sigma-Aldrich, 99%), and LiBr (Sigma-Aldrich, ≥99%) precursors were ground in an agate mortar for 5 min. Using a vertical planetary ball mill (TMAX Battery Equipment) with 45 mL vacuum jars, the powder was mixed to homogeneity at 110 rpm for 1 h. High-energy milling at 600 rpm for 20 h with 10 min of rest at 1 h intervals was used to induce the mechanochemical reactions producing the LPSCB compound. The resulting powders were annealed in ceramic crucibles in a muffle furnace (TMAX Battery Equipment) at 550 °C for 5 h, with a ramp rate of 25 °C/min. After the annealing sequence, the furnace was left to naturally cool.

To test the air stability of LPSCB, ~100 mg of electrolyte powder was exposed to air for 2, 5, 10, 20, 40, 70, 100, 180, 240, and 300 min in separate trials. The effect of exposure to ambient air was investigated using a controlled-volume desiccator (21.6 × 33.8 × 25.4 cm, ~18.5 L). Experiments were performed under standard laboratory conditions without strict control of relative humidity. This was deemed acceptable because, even at low relative humidities (~20% RH), the amount of H₂O vapor present in the desiccator volume exceeds the stoichiometric requirement to fully hydrolyze the 100 mg of electrolyte used per test. Therefore, precise humidity control was not necessary for evaluating the extent of hydrolysis or H₂S evolution. The setup within the desiccator is shown in Figure 1a, including the electrolyte powder evenly dispersed in a vial, a small battery-controlled fan, a H₂S sensor (Forensics Detectors), and a digital atmospheric sensor (Sensor Push), which recorded temperature, humidity, and pressure. For each exposure time, the H₂S sensor ppm output was video recorded in real time.

To test the electrochemical properties of pristine and exposed samples, electrolyte powder in quantities of ~100 mg was pelletized at 540 MPa into a ~1 mm thick disc of 10 mm diameter using two stainless-steel plungers encased in a PEEK sleeve placed in a TMAX cell press as seen in Figure 1b. Using electrodes attached to the stainless-steel plungers, electrochemical impedance spectroscopy (EIS) measurements were taken at an operating pressure of 250 MPa and room temperature through a potentiostat (Biologic SP-150). XRD of samples with no exposure, 10, 100, and 300 min of exposure was performed using a (Rigaku Ultima IV). Along with SEM imaging, elemental quantification was performed using EDS. The raw data were processed with EDAX software, which applied eZAF Smart Quant corrections to account for atomic number, absorption, and fluorescence effects, ensuring accurate quantification. XRD samples were prepared and sealed in the glovebox using Kapton tape to mitigate moisture exposure and then transferred to the X-ray machine. The same procedure was performed for the SEM samples to minimize contamination. We applied the same experimental and characterization procedures, including EIS measurements and H₂S-evolution analysis, to commercial LPSCB, using it as a benchmark to supplement the validation of our environmental stability protocol. All synthesis and EIS measurements were conducted under Argon atmosphere in a glovebox with controlled oxygen and humidity levels.

3. Results and Discussion

3.1. Initial Characterization

After synthesis of LPSCB, XRD was used to analyze the structure of the resulting powder and confirm the formation of the argyrodite phase. The XRD spectra comparing the produced LPSCB and commercial LPSCB electrolyte are shown in Figure 2. The structure shows the produced electrolyte has main diffraction peaks at 2θ = 25.4°, 29.8°, 31.1°, 39.5°, 44.7°, 47.5°, and 52.2°, with smaller peaks at 58.9°, 61.1°, 61.4°, 71.2°, 73.2°, 77.1°, 82.5°, 85.3°, and 85.4° which indicate a face-centered cubic structure of the space group of F$\overline{43}$m [25]. Commercial electrolyte exhibited almost identical main diffraction peaks, demonstrating that the desired argyrodite structure was achieved through our synthesis procedures. The use of Kapton tape, while preserving sample integrity, produces a broad scattering background at low angles which can partially obscure weaker reflections. Background subtraction algorithms used during processing may slightly distort peak intensities in this region. Despite this, the major characteristic peaks (for example 25.3°, 29.7°, and 31.1°) match well between both samples and confirm that the lab-based LPSCB follows the argyrodite structure of commercial LPSCB.

3.2. Ambient Humidity Exposure and Reactivity

After validating the structure of the synthesized electrolyte, air stability tests were conducted for each exposure duration. For each exposure test, ppm levels of H₂S gas, shown in Figure 3, were collected using a sensor for 100 mg samples of the produced LPSCB and commercial LPSCB. As exposure time increased, ppm levels showed a rapid increase followed by a plateau as a maximum amount of H₂S gas was produced. Since the 100 ppm upper limit of the H₂S sensor range was reached, the ppm values for the commercial electrolyte after 240 and 300 min might truly be above 100 ppm.

Utilizing the chemical reaction of LPSCB and water (Equation (3)), we developed a formula to obtain the mass of electrolyte participating in the reaction (Equation (4)):

$$M_E=P_a{V}_d{MW}_E{\displaystyle \frac{H}{5RT}}$$

(4)

where M_E is mass of electrolyte participating in the reaction, P_a is the average ambient pressure, V_d is the volume of the desiccator, MW_E is the molar weight of the electrolyte, H is the maximum ppm level of H₂S released after exposure, R is the universal gas constant, and T is the average temperature during the test. In the reaction, the mass is proportional to ${\displaystyle \frac{1}{5}}$ due to the 1:5 molar ratio of LPSCB to H₂S in the hydrolysis reaction of Equation (3). This equation could be modified to calculate the reaction mass in the hydrolysis reaction of other electrolytes by changing the molecular weight of the material reacting and, if necessary, the proportionality.

The reaction mass represents the amount of electrolyte that participated in the reaction over time, which is shown in Figure 4 for the LPSCB produced in a laboratory and the commercial LPSCB electrolyte. In this figure, the y-axis is the ratio of the electrolyte mass reacted with air humidity, $M_E\left(t\right)$, divided by the initial mass of the electrolyte, $M_{initial}$, which is 100 mg. The curves were fitted to a saturating exponential model, in the form of Equation (5), to investigate how rapidly the mass reaction occurs and consequently quantify the time constant, τ, or the duration it would take for 63.2% of the maximum mass, $M_{max},$ to react for each material. We selected this model because it aligns with our laboratory observations, indicating the formation of a protective layer over the electrolyte. This layer appears to slow down the reaction rate progressively until it becomes sufficiently thick to halt the reaction entirely.

$$NormalizedMassofReactedElectrolyte={\displaystyle \frac{M_E\left(t\right)}{M_{initial}}}=M_{max}\left(1-e^{-{\displaystyle \frac{t}{\tau }}}\right)$$

(5)

For the LPSCB produced in the laboratory and the commercial LPSCB electrolyte, the time constant is about 110 and 365 min, respectively. This indicates that the protective layer forms nearly three times faster on the produced LPSCB than on the commercial LPSCB, effectively halting the reaction. Additionally, Figure 4 shows that the produced LPSCB reacts more slowly with air humidity compared to the commercial LPSCB. Although the composition of the electrolytes is identical, the difference in reactivity would be attributed to variations in the synthesis procedure. Evidence of a structural discrepancy can be seen in Figure 2, where commercial LPSCB presents a broad halo around 25–27° 2θ. Our synthesis includes a 550 °C anneal that fully crystallizes the powder, eliminating any reactive amorphous phases. The commercial powder, if prepared without this high-temperature step, would retain some amorphous structures.

3.3. SEM and EDS Characterization

Through SEM imaging, significant agglomeration was observed as exposure time increases and has been reported in previous studies [31]. The SEM images in Figure 5 from the laboratory-produced LPSCB electrolyte clearly show this trend. As the pristine electrolyte powders (Figure 5a) are exposed to air for longer periods, there is a noticeable increase in the size of some particles, leading to a more pronounced difference between the larger agglomerates and the smaller particles surrounding them. This is most clearly seen in Figure 5c, where the electrolyte is exposed for 100 min. Agglomeration of particles and the consequential variation in particle size is crucial to understanding the effects of air exposure on ionic conductivity, especially as there is varying evidence on the impact of particle size on battery performance [46,47]. The SEM imaging of the laboratory-produced LPSCB exposed for 300 min (Figure 5d) deviates from the aforementioned trend and shows two distinct morphologies, providing insight into the extreme effects of prolonged air exposure. This data point underscores the potential for significant morphological changes over extended periods, which could impact the interpretation of long-term stability data.

The block-like formations observed in the 300 min exposed sample, shown in Figure 6, were further analyzed using EDS to compare with the more abundant spherical particles. The elemental analysis revealed no Cl and limited amounts of Br (1.8 wt.%), indicating the volatilization of halide dopants. This observation suggests a point of recrystallization may have occurred, negating the effects of earlier annealing procedures.

However, further analysis deemed 300 min as an overexposure of material to air for practical applications. This extended exposure highlights the potential for atypical changes in the material that are not representative of manufacturing environments or routine use. Since the focus of the current study is solely on the manufacturability of lithium batteries with LPSCB electrolytes, the 300 min SEM and EDS exposure data are excluded, and a maximum exposure time of 100 min is recommended for future studies.

To ensure precision in our findings, EDS analysis (Figure 7) was focused primarily on pristine samples and those exposed for 10 and 100 min. These findings are critical to understanding the elemental composition changes within the context of practical air exposure durations. Lithium is not included among these elements, as the EDS instrument is not sensitive enough to detect it. Therefore, discussion of EDS data will be limited to change in trends for the detectable elements.

Some trends in the chemical stability of LPSCB can be elucidated from EDS data. From the consistent decrease in P and S as the material experiences longer exposures to air, we can infer the removal of these elements from the surface due to the hydrolysis reaction. The products of this reaction could be volatile such as H₂S gas or soluble in the base material such as Li₃PO₄. Particularly for sulfur, an additional cause of the decrease in content could be due to the initial volatility of the sulfide ions as seen in the ionic formula (Li⁺)₆(PS₄³⁻)S²⁻X⁻ [20]. This readily available sulfur easily combines with moisture to form H₂S.

Furthermore, the simultaneous increase in Cl and Br weight percentages is most likely due to relative concentrations which make up for the loss of P and S, especially as the halides can default to their original precursors in the structure as LiCl and LiBr without changing surface compositions. On the other hand, a jump in O content is observed after 10 min of air exposure, which is obviously due to the initial addition of O with the base material elements. There is a subsequent and counterintuitive decrease in O weight percentage as exposure times increase. This trend might have several causations. We can infer based on theoretical reaction mechanisms that Li will be redistributed into the reactant compounds and will initially react to form LiOH, which will introduce a large amount of O to the sample. Further hydrolysis will lead to the formation of Li₃PO₄ which solubilizes O within the structure and decreases its concentration on the surface.

3.4. X-Ray Diffraction

As seen in XRD patterns in Figure 8, the main diffraction peaks of the F$\overline{43}$m structure begin to diminish after 100 min of exposure for both laboratory-produced and commercial LPSCB electrolytes. The appearance of some amorphization is also prevalent at the range of 15° to 20° diffraction angles. In addition to the onset appearance of impurities, there are slight shifts in peaks ranging from 5 to 50° towards higher diffraction angles. In similarity with EDS analysis, after 100 min of exposure, the peak shifts to higher diffraction angles are indicative of the inclusion of O atoms, which are around half the size of base material elements or smaller [48]. After 300 min of exposure, the initial amorphization has substantially increased and several new peaks have formed, indicating increased impurities such as LiCl, LiBr, LiOH, and Li₂S.

3.5. Ionic Conductivity

After measuring average resistances of solid electrolyte pellets by fitting electrochemical impedance spectroscopy graphs to Nyquist plots, ionic conductivity was calculated using Equation (6):

$$\sigma ={\displaystyle \frac{\delta }{RA}}$$

(6)

where $\sigma$ is the ionic conductivity of the solid electrolyte, $\delta$ is the thickness of the pellet, $R$ is the bulk resistance of the pellet, obtained from fitting the Nyquist plot, and $A$ is the cross-sectional area of the pellet. The measurement setup is shown in Figure 1b.

The resulting ionic conductivities of the laboratory-produced LPSCB and commercial LPSCB electrolytes are shown in Figure 9, which also displays the equations of each data set as fitted to an exponential decay model, in the form of Equation (7). As before, the time constant (τ) for the exponential decay equation was evaluated from models fitted to each material. In this case, the time constant is the inverse of the rate of decay of the exponential term as presented in Equation (7):

$$\sigma \left(t\right)={\sigma }_f+\left({\sigma }_0-{\sigma }_f\right)e^{-{\displaystyle \frac{t}{\tau }}}$$

(7)

where ${\sigma }_0$ and ${\sigma }_f$ represent the initial (pristine) and final (after long-term exposure) ionic conductivities of the electrolyte, respectively.

As seen, the ionic conductivity of the produced electrolyte is consistently higher than that of the commercial electrolyte at all exposure times. The time constants for degradation of the laboratory-produced and the commercial LPSCB electrolyte are 100 and 50 min, respectively, meaning that 63.2% of the decay of ionic conductivity will occur for each electrolyte after these respective durations.

The lab-synthesized LPSCB electrolyte loses approximately 1%, 9%, and 58% of its initial ionic conductivity after exposure to ambient air humidity for 1 min, 10 min, and 100 min, respectively. In comparison, the commercial LPSCB exhibits greater degradation, losing about 2%, 18%, and 85% of its initial ionic conductivity over the same exposure durations.

Using the fitted models derived earlier (Equations (5) and (7)), we can project the evolution of both the reacted-mass fraction and ionic conductivity out to longer times (Figure 4 and Figure 9). For the lab-produced LPSCB (τ ≈ 110 min for mass reaction; τ ≈ 100 min for ionic conductivity), extrapolation to 300 min predicts that roughly 0.13 wt% of the powder will have reacted, only ~0.03 wt% more than at 100 min, because the reaction rate is already approaching its passivation-controlled plateau. In contrast, ionic conductivity continues to decay exponentially, falling to ≈5% of the pristine value (≈0.11 mS cm⁻¹) at 300 min. The commercial powder, with a shorter conductivity time constant (τ ≈ 50 min), is expected to retain only ≈2% of its initial conductivity after the same period. Beyond ~3τ (≈300 min for the lab material) additional degradation becomes increasingly marginal because the growing hydrolysis layer inhibits further reaction. Practically, this means that accidental long-term exposures—for example, a ruptured battery module left open to humid air—could render the electrolyte nearly non-conductive, whereas routine manufacturing steps that limit exposure to ≤10 min remain in the low-degradation regime (>90% conductivity retention).

The variation in degradation times may be attributed to a difference in synthesis procedures between commercial lab-produced LPSCB. Additionally, the formation of different reaction byproducts, as discussed in Equations (1)–(3), may contribute to conductivity decay. Moreover, delays in degradation and sustained conductivity for a window of time suggest that LPSCB can tolerate brief air exposure during manufacturing processes without a significant loss in performance. Further investigation into the complete reaction mechanisms, the resulting products, and the factors contributing to improved air stability is necessary. It will also be important to assess how short-term air exposure influences overall cell performance in continuations of this study.

3.6. Simplified Testing for Environmental Stability

This study demonstrated that the reaction rate and ionic conductivity degradation of the electrolyte under ambient air humidity follow saturating exponential and exponential decay behaviors, respectively. As a result, testing at only two time points is sufficient to extract the key model parameters, enabling prediction of environmental stability over a wide range of exposure durations with over 90% accuracy. To validate the model, one or two additional time points are recommended for comparing predicted values with experimental results. This approach significantly simplifies and accelerates environmental stability testing. We recommend using 10 min and 100 min exposures to establish the model equations and 1 min and 30 min exposures to validate the model’s accuracy and confirm the reliability of the equations’ extrapolation and interpolation capabilities. Due to challenges in achieving repeatable results, a 1 min exposure test is recommended only when a highly robust and sensitive experimental setup is available. If test repeatability is an issue, a 5 min exposure can be used.

4. Conclusions

In this study, we investigated the behavior of the sulfide-based argyrodite solid electrolyte Li₆PS₅Cl₀.₅Br₀.₅ under ambient air humidity to validate its synthesis process, evaluate the effects of humidity on its reactivity and ionic conductivity, and establish a standardized protocol for assessing the environmental stability of argyrodite electrolytes. The results showed that the saturating exponential equation effectively models the reaction rate of electrolyte upon exposure to ambient air humidity. Additionally, the exponential decay equation provides the best fit for predicting the decline in ionic conductivity as a function of exposure time. While exposure to air humidity does lead to performance degradation, our synthesized electrode exhibits notable resilience, retaining about 99%, 91%, and 42% of its pristine ionic conductivity when it is exposed to the ambient air humidity for 1 min, 10 min, and 100 min, respectively. The fraction of electrolyte mass that reacts with ambient air humidity is approximately 0.001%, 0.014%, and 0.096% of the initial mass after 1, 10, and 100 min of exposure, respectively. These values offer a useful approximation for scientists and engineers to estimate performance changes in Li₆PS₅Cl₀.₅Br₀.₅ when exposed to air humidity during battery manufacturing.

Based on extensive environmental stability tests, we proposed a simplified protocol: 10 min and 100 min exposures are sufficient to establish model equations, while 1 min and 30 min tests validate the model and confirm its interpolation and extrapolation accuracy. The 1 min test is only recommended with robust, sensitive setups; otherwise, a 5 min exposure is a more reliable alternative. We recommended these durations as a standard protocol for evaluating the environmental stability of argyrodite electrolytes, enabling consistent and effective comparison across different electrolytes.