Air-Breathing Aqueous Sulfur Flow Battery for Ultralow-Cost Long-Duration Electrical Storage

Air-Breathing Aqueous Sulfur Flow Battery for Ultralow-Cost Long-Duration Electrical Storage


Highlights

• •
  Chemical cost analyzed for 40 rechargeable couples developed over the past 60 years
• •
  Aqueous sulfur/sodium/air system identified with ultralow chemical cost of ∼US$1/kWh
• •
  Air-breathing flow battery architecture demonstrated at laboratory scale
• •
  Techno-economic analysis shows installed cost is comparable with PHS and CAES

Context & Scale

Wind and solar generation can displace carbon-intensive electricity if their intermittent output is cost-effectively re-shaped using electrical storage to meet user demand. Reductions in the cost of storage have lagged those for generation, with pumped hydroelectric storage (PHS) remaining today the lowest-cost and only form of electrical storage deployed at multi-gigawatt hour scale. Here, we propose and demonstrate an inherently scalable storage approach that uses sulfur, a virtually unlimited byproduct of fossil fuel production, and air, as the reactive components. Combined with sodium as an intermediary working species, the chemical cost of storage is the lowest of known batteries. While the electrical stacks extracting power can and should be improved, even at current performance, techno-economic analysis shows projected costs that are competitive with PHS, and of special interest for the long-duration storage that will be increasingly important as renewables penetration grows.
Summary

The intermittency of renewable electricity generation has created a pressing global need for low-cost, highly scalable energy storage. Although pumped hydroelectric storage (PHS) and underground compressed air energy storage (CAES) have the lowest costs today (∼US$100/kWh installed cost), each faces geographical and environmental constraints that may limit further deployment. Here, we demonstrate an ambient-temperature aqueous rechargeable flow battery that uses low-cost polysulfide anolytes in conjunction with lithium or sodium counter-ions, and an air- or oxygen-breathing cathode. The solution energy density, at 30–145 Wh/L depending on concentration and sulfur speciation range, exceeds current solution-based flow batteries, and the cost of active materials per stored energy is exceptionally low, ∼US$1/kWh when using sodium polysulfide. The projected storage economics parallel those for PHS and CAES but can be realized at higher energy density and with minimal locational constraints.
Introduction

The rapidly dropping cost of wind and solar electricity generation, as illustrated by levelized costs of electricity (LCOE) that are now competitive, or nearly so, with fossil fuel generation,^mce-anchor1 highlights the need for low-cost electrical storage that can transform intermittent renewable power into predictable and dispatchable electricity generation, and potentially even baseload power. Such a revolutionary outcome will require energy storage with costs well below the trajectory of current technology, while also being safe, scalable, long-lived, and sufficiently energy dense for widespread deployment, including in space-constrained environments. Emerging use-case studies suggest that installed costs of <$50/kWh, operating over multi-day or longer durations, will be required for renewable-based generation to compete economically with existing fossil plants on a drop-in basis. It is unclear whether electrochemical storage can meet these challenges.
Here, we first review the underlying chemical cost of energy storage for about 40 electrochemical couples representing all major classes of rechargeable batteries developed over the past 60 years. From this analysis, it is clear that the best opportunities for overcoming the above challenges reside with electrochemical couples that use ultralow-cost, highly abundant raw materials. Among these, sulfur has the 14^th highest crustal abundance and is widely available as a byproduct of natural gas and petroleum refining.^mce-anchor2 Sulfur also has the lowest cost per stored charge of known redox active materials, next to water and air (see Table S1; in US$/kAh, sulfur 0.15, zinc 3.66, graphite 32.27, and LiCoO₂ 292.14). In this work, we demonstrate an ambient-temperature, air-breathing, aqueous polysulfide flow battery that exploits sulfur's intrinsic advantages, and show using techno-economic analyses that such an approach has the potential to meet future storage needs for renewable energy.
Results and Discussion

Chemical Cost Comparison for Battery Electrochemistries

A reasonable starting point for bottom-up analysis of the economics of any battery technology is the cost of the cathode, anode, and electrolyte, normalized to the stored electrical energy. We define this quantity as the chemical cost of energy storage (abbreviated as chemical cost and given in US$/kWh), building on an earlier study^mce-anchor3 that analyzed the elemental costs of electrochemical couples. The chemical cost for about 40 battery chemistries is plotted in Figure 1 against the year that each electrochemistry was introduced (all costs are in 2017 US$). Although we have not attempted to exhaustively list all extant electrochemical couples, exemplars of each of the major classes of bulk rechargeable batteries are included. The numerical results plotted in Figure 1 are tabulated in Table S2, and details of the calculations, including input parameters, key assumptions, and literature sources, are given in the Supplemental Information.

A striking result apparent in Figure 1 is that the chemical cost of new battery chemistries has in general systematically increased rather than decreased over 60 years of battery development. We believe that to a large extent this trend can be attributed to the pursuit of higher energy density, as exemplified by the deployment of Li-ion batteries in portable device and transportation applications once dominated by NiMH, NiCd, or Pb-acid batteries. Figure 1 shows that Li-ion technology itself spans nearly a 3-fold range of chemical cost, from ∼US$35/kWh (C₆/0.3Li₂MnO₃-0.7LiMn_0.5Ni_0.5O₂) to ∼ US$100/kWh (C₆/LiCoO₂). Given that energy density has generally increased over time, the rising chemical cost implies that the high cost of new materials has more than compensated. Figure 1 also shows that while several aqueous electrochemical couples have chemical cost below US$10/kWh, the lowest cost of which is Zn-air, which as a primary chemistry dates to the year 1878, several aqueous electrochemical couples have chemical costs greater than low-cost Li-ion. This again reflects the high cost of synthesized active materials relative to stored energy. The case of Na/S is instructive; while it has the lowest chemical cost in our plot, excluding present results, it is known that high-temperature Na/S batteries are among the most expensive at system level (∼US$800/kWh). This is due to the high cost of supporting components and balance of plant. Conversely, an ambient-temperature sodium-sulfur chemistry has the potential for exceptionally low system cost, given a starting chemical cost of ∼US$1/kWh. These considerations led us to explore new ambient-temperature alkaline-sulfur chemistries, culminating in the air-breathing aqueous polysulfide couples, denoted in Figure 1 as Li₂S_x/air and Na₂S_x/air, the lowest-cost members of which have Air-Breathing Aqueous Polysulfide Concept

Whereas nonaqueous lithium-sulfur^{mce-anchor4, mce-anchor5, mce-anchor6} and high-temperature sodium-sulfur batteries^mce-anchor7 use sulfur as the cathode, an all-aqueous system must use sulfur as the anode material to preserve aqueous stability while reaching a meaningful cell voltage. Solubilized aqueous sulfur electrodes have previously been paired with halogenated catholytes in flow batteries,^{mce-anchor8, mce-anchor9, mce-anchor10} used as the catholyte versus “protected” lithium metal anodes,^mce-anchor11 and used as the anolyte with lithium intercalation cathodes.^{mce-anchor12, mce-anchor13} (Here, “anode” and “cathode” refer to solid-phase active materials, whereas “anolyte” and “catholyte” refer to fluids with solubilized active materials.) In each of these cases, the chemical cost of storage is dominated by the nonsulfur electrode. Thus, we reasoned that if aqueous anolytes of highly soluble metal polysulfides (up to 12 M)^mce-anchor11 could be matched with a similarly low-cost catholyte, unprecedented low storage economics could be achieved. Our scheme, illustrated in Figure 2, is an aqueous flow battery that pairs a polysulfide anolyte with an oxygenated/aerated salt solution as the catholyte. At the anolyte side, charge transfer occurs to a negative current collector connected to the external circuit. At the catholyte side, a single or dual oxygen evolution/reduction cathode configuration is used. This configuration pairs two half-reactions, namely polysulfide oxidation/reduction and oxygen evolution reaction (OER)/oxygen reduction reaction (ORR), which unlike conventional rechargeable couples, do not share a common working ion.

Specifically, during operation, Li⁺ and Na⁺ (or other metal ions) shuttle between the electrolytes. In the anolyte, these working ions participate in the polysulfide redox reactions. In the catholyte, in order to maintain electroneutrality, ions are generated or consumed through oxygen electrochemistry, using water as both solvent and reactant. The reversible capacity of the cell is determined by either the total concentration of alkali-metal ions, or by the sulfur concentration in the anolyte, whichever is limiting. The generation and consumption of protons (for acid catholyte) or hydroxyls (for alkaline catholyte) lead to pH swings in the catholyte. In the alkaline catholyte case, hydroxyl crossover is not anticipated to be detrimental to performance, while in the acid catholyte case, protons must be confined to the catholyte chamber to prevent mixing with the alkaline anolyte. Here, we use ceramic membranes as the separators. Note that this scheme is not a sulfur-air battery, as there is no direct reaction between any sulfur species and oxygen. We searched for related concepts in the literature, and found a recently proposed Zn-air battery that also utilizes acidic catholyte and alkaline anolyte separated by a solid electrolyte (LiSICON) membrane,^mce-anchor14 but in a stationary (non-flowing) design. We also found a published patent application^mce-anchor15 that proposes an all-alkaline flow cell in which an air cathode is paired with a solid sulfur-polysulfide anode; here the working ions are hydroxyl ions and the system requires an anion exchange membrane.
In the present approach, cells may be assembled in the discharged state, in which case the alkali-metal working ion is provided by one of several possible low-cost salts dissolved in the catholyte. In the examples below we use Li₂SO₄ and Na₂SO₄. Cells may also be assembled in the charged state, in which case the working ion is provided by lithium or sodium sulfide dissolved in the anolyte. To avoid H₂S formation at the anolyte, adequately high pH (≥12) is maintained by the addition of a suitable base (here, LiOH or NaOH). The catholyte may be either acidic or alkaline, providing cell voltages of ∼1.68 V and ∼0.85 V, respectively, at standard state equilibrium, as explained via the cell reactions.
Acidic Catholyte





Alkaline Catholyte





In between these limits, the equilibrium cell voltage varies continuously with pH as shown in the Pourbaix diagram in Figure S1; i.e., it is ∼1.26 V at neutral pH.
In order to determine the chemical cost, the stable speciation range of the polysulfide must be taken into account. Although polysulfide solubilities in aqueous solutions can reach as high as 12 M sulfur concentration,^mce-anchor11 stability issues (see below) may limit the practical capacity to less than that for complete reduction of sulfur according to the reaction 2A + S → A₂S (where A is Li or Na). However, even with a more limited range of sulfur reduction, exceptionally low chemical cost is attainable while reaching energy densities higher than many previous flow batteries. This is shown in Tables S4 and S5, where energy density and chemical cost for catholytes and anolytes having 5 M of Li or Na and 5 M of S, respectively, have been calculated. If the entire sulfur reduction capacity is utilized, the energy densities are 121 and 58 Wh/L for acidic and alkaline catholyte, respectively. Increasing the anolyte S concentration to 10 M raises the upper bound to about 145 Wh/L and 70 Wh/L for acidic and alkaline catholyte, respectively. If cycling is restricted to just 25% of the full sulfur reduction capacity, the corresponding energy densities are 60 and 29 Wh/L when using acidic and alkaline catholyte, respectively. This capacity range corresponds to the solution range Li₂S₂-Li₂S₄ or Na₂S₂-Na₂S₄, for which we later demonstrate stable cycling over ∼1,000 hr. For sodium-polysulfide chemistry, the chemical cost is remarkably low, only US$0.4–1.7/kWh (using acidic catholyte at 5 M S), depending on the utilization of sulfur theoretical capacity (100%–25%). At 50% utilization or higher, one reaches the lowest chemical cost to our knowledge of any rechargeable battery (Figure 1). When using lithium polysulfide rather than sodium polysulfide, the chemical cost is US$2.0–5.0/kWh, still lower than existing flow batteries.
In the remainder of the paper, we first demonstrate key performance characteristics of the proposed electrochemical couples. Since both catholyte and anolyte are fluids, our electrochemical couples lend themselves to a flow battery design, which is also demonstrated. An attribute of flow batteries is the ability to independently size the power stack and chemical storage capacity to meet desired energy to power ratios. In such architecture, the contribution to system cost of the power stack can be minimized by sizing it only as large as is necessary to meet the maximum power requirement. Using the methodology in a recent techno-economic analysis of flow batteries,^{mce-anchor16, mce-anchor17} we show that the proposed electrochemical storage system has attractive performance and cost attributes very similar to those of pumped hydroelectric storage (PHS) and underground compressed air energy storage (CAES), which are currently the lowest-cost energy storage technologies.^mce-anchor18 Finally, the exceptionally low chemical cost allows electrochemical storage to address a new domain of long-duration discharge that was not previously accessible in a cost-effective manner.
Verifying the Catholyte and Anolyte Reactions

Using H cells (Figure 3A), we confirmed that OER/ORR occurs at the catholyte side by observing the changes in cell voltage induced by changing the gas composition or by cycling the gas flow on and off. An immediate response of the cell voltage to a change in gas flow conditions is clearly seen in Figures 4A and 4B for acidic and alkaline catholyte, respectively. Then, to determine whether the charge-storing capacity of the catholyte corresponds to the concentration of added salt (Li₂SO₄ or Na₂SO₄), cells with an excess of anolyte were prepared using the modified H-cell design in Figure 3B. Results in Figures 4C and 4D show that the measured capacity of these cells corresponds almost exactly to the theoretical capacity calculated from the Li⁺ or Na⁺ concentration. In addition, the cell voltage is the same for both alkali ions, as expected from the cell reactions. In the anolyte, polysulfide redox activity was readily discerned from changes in color with state of charge.


Cell Efficiency, Impedance, and Power Density

The voltage efficiency of the cells is determined by the OER/ORR reaction as well as other contributions to cell impedance. To separate these, we first measured the voltage efficiency as a function of current density and temperature, using the standard H cells as well as a cell with separate reference electrodes for the cathode and anode. Cell tests were conducted in the temperature range 25°C–70°C, within which increasing the temperature was observed to dramatically reduce cell impedance. Increasing temperature also sharply increases Na₂SO₄ solubility in the catholyte beginning at about 35°C.^{mce-anchor19, mce-anchor20}
Galvanostatic step charges and discharges produced voltage efficiency results shown in Figures 5A and 5B for Li and Na chemistry, respectively. Values plotted are the round-trip voltage efficiency, i.e., the discharge voltage divided by the charge voltage. Using the acidic catholyte and dual catalysts (IrO₂ for OER and Pt black for ORR), the voltage efficiency reaches 71%–74% for both Li and Na solutions at low current density (0.065 mA/cm²). Here, efficiency is limited by the OER/ORR reaction. Using a less efficient, Pt mesh, single cathode rather than the dual cathodes with IrO₂/Pt black catalysts, and holding all other cell parameters constant, the voltage efficiency at 70°C is about 20% lower, as shown in the Supplemental Information (Figure S2).

With increasing current density, the voltage efficiency decreases since additional contributions to impedance become significant. The results in Figure 5A were obtained using a 150-μm-thick LiSICON membrane, with a measured ionic conductivity of 0.28 mS/cm at room temperature and 0.6 mS/cm at 55°C, whereas the results in Figure 5B were obtained using a NaSICON membrane of higher ionic conductivity, 2 mS/cm at room temperature,^mce-anchor21 but also much greater thickness of 1,000 μm. We conducted most of our experiments using Li chemistry simply due to the greater availability of LiSICON membranes. To separate the membrane resistance from other contributions, an experiment was conducted using two different reference electrodes, an alkaline Hg/HgO (in 1 M LiOH) reference electrode at the anolyte side and a Hg/Hg₂SO₄ (in saturated K₂SO₄) reference electrode at the catholyte side. A stainless-steel mesh anode current collector and dual cathodes with IrO₂ and Pt black catalysts were used. Step-galvanostatic scans consisting of sequential 5 min galvanostatic charge or discharge at various current densities were performed while measuring voltage. The overall polarization is obtained from the voltage difference between cathode and anode, the cathode contribution from the voltage between cathode and Hg/Hg₂SO₄ reference electrode, and the anode contribution from the voltage between anode and Hg/HgO reference electrode. Figure 5C plots the total polarization and the cathode, anode, and membrane contributions as a function of current density, using the 150-μm-thick LiSICON membrane. It is seen that the cathode and anode polarization remain relatively constant, with the anode polarization being lower by at least a factor of five over the measured current density range (up to 7 mA/cm²). The membrane polarization, on the other hand, increases with current density, and is about 50% of the total impedance at current density of 3 mA/cm².
Further evidence that polarization and power density in the current unoptimized design are primarily limited by membrane resistance appears in Figure 6, where the polarization and power density of H cells were measured for two LiSICON membranes with different thicknesses (both at 55°C using a single Pt mesh cathode coated with Pt black). With the thinner membrane of 50 μm thickness, decreased polarization is evident and a higher power density with peak value of 5.1 mW/cm² at 7.1 mA/cm² was reached. In comparison, a 150-μm-thick membrane of the same composition yielded a peak power density of 3.4 mW/cm² at 5.5 mA/cm². These values of power density are low compared with those for polymer-membrane-based redox flow batteries or proton exchange membrane (PEM) fuel cells, and may be significantly improved if membrane resistance is lowered. Nonetheless, we use these experimentally observed values as starting input parameters in the techno-economic (T-E) model for the flow battery, discussed later.

Another important cost factor is the cost of catalyst required to support a given power density. Several non-platinum group metal (PGM) catalysts for OER and ORR in acid media have been proposed,^{mce-anchor22, mce-anchor23, mce-anchor24, mce-anchor25} which we are in the process of evaluating. However, we have also investigated the impact of Pt black loading on the charge-transfer resistance of the experimental flow cell (Figure 3D), using galvanostatic electrochemical impedance spectroscopy (GEIS). The US Department of Energy's (DOE) future PEM fuel cell total PGM target for Pt loading is 0.15 mg/cm²,^mce-anchor26 but this loading accommodates about 100 times higher power density (800 mW/cm²) than is assumed in our model. Experiments were conducted using two Pt black loadings, one higher (0.21 mg/cm²) than the DOE target and one a factor of five lower (0.03 mg/cm²). We found that the impact of temperature was much greater than the impact of loading. Figure 7A shows selected Cole-Cole plots from these measurements, in which the higher frequency arc is due to membrane impedance and the lower frequency arc is correlated with ORR kinetics.^mce-anchor27 A small inductive loop at the lowest frequencies is observed in each case, consistent with the literature,^mce-anchor27 which we did not attempt to fit. Figure 7B summarizes the GEIS results as a plot of total cell area-specific resistance (ASR) versus temperature. It is seen that the curves converge with increasing temperature such that at 50°C and above, neither current density nor loading significantly changes the ASR. At lower temperature where results are more differentiated, increasing discharge current density (comparing the red and black curves in Figure 7A) primarily decreases the ORR charge-transfer resistance, consistent with Butler-Volmer behavior whereby a larger overpotential results in faster kinetics. Increasing temperature (comparing blue and black curves in Figure 7A) decreases both membrane and ORR impedance. Increasing the catalyst loading with other parameters held constant (comparing blue and green curves) decreases the ORR charge-transfer resistance, but only at the lower temperatures, and does not change the membrane resistance. We use these results later for T-E modeling of the flow battery. Unlike Pt black, IrO₂ is not readily available as high surface area powder, so the impact of its loading on OER kinetics during charging was not investigated. In the T-E modeling, we assume similar loadings as for Pt black.

Deep Cycling Tests of Catholyte and Anolyte Stability

The ability of the catholyte and anolyte to undergo sustained deep cycling, and the durability of the cell components in contact with catholyte and anolyte, were tested using modified H-cell designs (Figures 3B and 3C). To test the catholyte, cells were assembled in which the catholyte capacity was lower than the anolyte capacity (provided by a 4 M S solution). Using a single cathode (Pt mesh), galvanostatic cycling (0.325 mA/cm²) was conducted at room temperature, with the capacity during each cycle being limited to 96% of the catholyte theoretical capacity, based on the starting alkali ion concentration. Results are shown in Figure 8A, plotted as voltage-capacity curves for the 1^st, 10^th, 20^th, and 30^th galvanostatic charge/discharge cycles, spanning a total test time of over 1,600 hr. Dry air was continuously flowed into the catholyte chamber via a dispersion tube during discharge, and water was periodically added to the catholyte to compensate for evaporation. Initially, we chose to use capacity-limited cycling out of concern that proton intercalation into the LiSICON might occur upon the exhaustion of Li⁺ without a clear voltage indication. However, as seen in Figure 8A, and in expanded scale in Figure S3, the charge curves are almost invariant after the first charge, indicating that the results would be nearly identical if a voltage cutoff were used. The slope of the voltage profile clearly deviates toward the end of the charge (Figure S3) as expected from an increased mass transfer resistance resulting from the depletion of working ion (Li⁺). The minor variations in polarization for the discharge curves were found to be correlated with fluctuations in catholyte water level and air flow rate, which are expected to affect ORR kinetics. Over the >2 month duration of the test, the cell impedance did not grow detectably, pointing to good stability of the LiSICON membrane in contact with the acidic catholyte as well as stability of catholyte and anolyte with their respective electrodes, Pt mesh and stainless steel. (The voltage efficiency in these tests is low since they are conducted at room temperature and only use the single Pt mesh cathode for OER and ORR.)

In the Supplemental Information, Figure S4, we show that extreme variations in either water loss or water excess (dilution of the catholyte), or interruptions in gas flow, cause large deviations in coulombic efficiency, but also that the cells are tolerant of such variations and return to stable operation once the water content and gas flow at the cathode are restored to their initial values.
Cycling stability of the polysulfide anolyte was tested in cells with two gas-tight chambers of differing volume containing the same anolyte, thus deeply cycling the smaller chamber (Figure 3C). LiSICON membrane and stainless-steel electrodes were used, and an alkaline Hg/HgO in 1 M LiOH reference electrode was placed in the larger chamber. Selection of the appropriate speciation range over which to cycle the anolyte requires consideration of complex equilibria.^{mce-anchor28, mce-anchor29, mce-anchor30} Aqueous alkali-metal polysulfide solutions contain a wide range of species, including the alkali-metal cations (Li⁺, Na⁺, or K⁺), H₂O, OH⁻, H⁺, H₂S, HS⁻, S²⁻, S₂²⁻, S₃²⁻, S₄²⁻, and S₅²⁻. Polysulfide solubility and stability are strongly dependent on pH,^{mce-anchor12, mce-anchor29} alkali-metal cation,^mce-anchor30 nominal polysulfide speciation, and concentration^mce-anchor31 as well as temperature.^mce-anchor32 At low pH (<7), HS⁻ is the primary species, and the H₂S molecule is the predominant reduced product in the polysulfide solution.^{mce-anchor29, mce-anchor33} At intermediate alkalinity (pH 9–14), the primary polysulfide species are S₄²⁻ and S₅²⁻ instead of HS⁻,^mce-anchor29 although without good sealing H₂S may still be generated at pH ≈ 12, especially during the reduction reaction.^mce-anchor12 In a recently reported lithium polysulfide/LiMn₂O₄ battery, adding porous SBA-15 silica adsorbent to the polysulfide anolyte was found to reduce irreversible capacity loss and to improve capacity retention.^mce-anchor13 The improvements were attributed to the suppression of gaseous H₂S release, preventing continuous loss of sulfur. In highly alkaline polysulfide solution (e.g., >3 M OH⁻), the predominant species are S₃²⁻and S₂²⁻.^mce-anchor34 A Li₂S₄ solution under these conditions may disproportionate into S⁰ and S₂²⁻/S₃²⁻ during long-term storage.^{mce-anchor29, mce-anchor34} At temperatures above 80°C, thiosulfate is readily formed via the reaction S_n²⁻ + mOH⁻ → S_n-mO_m^2- + mHS⁻.^{mce-anchor32, mce-anchor35} This parasitic disproportionation reaction is detrimental to the stability of polysulfide anolyte, but slow at moderate temperature.^{mce-anchor34, mce-anchor36}
Taking into account these prior studies, we prepared starting polysulfide solutions containing nominal stoichiometry Li₂S₄ and 1 M or 3 M LiOH at room temperature, and operated cells from room temperature up to 70°C, such that the starting predominant sulfur species is S₄²⁻ and the major degradation reactions are minimized. To date, we have found that the most stable cycling is obtained when we constrain the composition range to Li₂S₂ to Li₂S₄ or Na₂S₂ to Na₂S₄; see Figure 8B for cycling results at 5 M S concentration obtained over 720 hr. In the interest of increasing energy density and further lowering cost, the ability to reversibly cycle beyond the Li₂S or Na₂S solubility limit is desirable. Our group has previously shown that percolating nanocarbon suspensions can be used to improve charge transfer in fluid electrodes to the extent of realizing reversible precipitation of Li₂S.^{mce-anchor37, mce-anchor38} A similar strategy could be applied here to extend the polysulfide capacity range.


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