Systems and methods of processing brine to capture lithium ions

Abstract

The present disclosure provides for methods and systems for processing brine to obtain lithium. The present disclosure provides for methods that leverages intercalation materials and drives lithium capture using a chemical approach to treating brine as opposed an electrochemical approach, which can be complicated and expensive. The present disclosure provides for a chemical redox-driven approach to selectively extract Li+ from brine.

Description

[0001] SYSTEMS AND METHODS OF PROCESSING BRINE TO CAPTURE LITHIUM IONS

[0002] CROSS-REFERENCE TO RELATED APPLICATION

[0003] This application claims priority to U.S. provisional application entitled “SELECTIVE LITHIUM EXTRACTION FROM BRINE VIA CHEMICAL REDUCTION OF IRON PHOSPHATE WITH AQUEOUS IRON COMPOUNDS,” having serial number 63 / 604,965 filed on December 1, 2023, which is entirely incorporated herein by reference.

[0004] BACKGROUND

[0005] Demand for lithium (Li) continues to increase due to the increased use of Li-ion batteries. Sources of Li include ores, brines and seawater. Extracting Li from these sources can be expensive and harm the environment. Alternative methods of sourcing Li are needed.

[0006] SUMMARY

[0007] The present disclosure provides methods and system for processing brine to obtain lithium and the like.

[0008] In an aspect, the present disclosure provides methods of processing a brine solution to capture lithium ions (Li+). The method includes mixing the brine solution with an Feecompound and a complexing agent to form a first mixture that includes a dissolved Feecomplex including the Li+; exposing the first mixture to a plurality of particles, wherein the particles comprise an intercalation material, wherein the intercalation material includes a transition metal cation compound, wherein a portion of the transition metal cations in the intercalation material are reduced and Li+is intercalated into the intercalation material; and exposing the plurality of particles to an oxidizing solution, wherein the Li+in the intercalation material is released when the portion of the transition metal cations in the intercalation material are oxidized to reform the intercalation material and a Li+-solution. In an aspect, the Fe2+compound is selected from the group consisting of: FeCI2, FeSO4, Fe(NO3)2, or a mixture thereof, the complexing agent is selected from the group consisting of: citrate, ethylenediaminetetraacetic acid, and a mixture thereof, and / or the intercalation material contains a Fe3+transition metal and is an intercalation Fe3+-solid material. In an aspect, the brine solution is selected from the group consisting of: a natural brine, a synthetic brine, or a combination thereof and in particular, the brine solution is selected from the group consisting of: a geothermal brine, an oil field brine, a brine from hard rock lithium mining, or a combination thereof.

[0009] In an aspect, the present disclosure provides for a system for chemically treating a brine solution to capture lithium ions (Li+), comprising: a first structure for mixing a brine solution with an Fe2+-compound and a complexing agent to form a first mixture including a Fe2+-complex; a flow system to flow the first mixture; and a second structure that receives the first mixture, wherein the second structure functions to contact a first mixture with a plurality of particles, wherein the first mixture includes a Fe2+-complex, wherein the particles comprise an intercalation material, wherein the intercalation material includes a transition metal cation compound, wherein a portion of the transition metal cations in the intercalation material are reduced and Li+is intercalated into the intercalation material to form Li+intercalated intercalation material; wherein the second structure also functions to mix the plurality of particles with an oxidizing solution after the first mixture is removed from the second structure, wherein the Li+in the Li+intercalated material is released when the portion of the transition metal cations in the intercalation material are oxidized to reform the intercalation material and a Li+-solution.

[0010] BRIEF DESCRIPTION OF DRAWINGS

[0011] Further aspects of the present disclosure will be more readily appreciated upon review of the detailed description of its various embodiments, described below, when taken in conjunction with the accompanying drawings.

[0012] FIG. 1 shows cyclic voltammetry (CV) curves measured using two platinum disc electrodes as the working and counter electrodes and Ag / AgCI with saturated KOI as the reference electrode with scan rate of 200 mV s1in chloride brine solutions containing 39 mM of Li+and 3.04 M of Na+. The redox mediators were 0.1 M ethylenediaminetetraacetic acid (EDTA)-Fe2+with 1.3 M of KOH to adjust pH to 7 (blue), 0.1 M of Citrate-Fe2+(black), and 0.1 M of Fe2+ / Fe3+(red). The Fe concentration was prepared to contain initially equimolar Fe2+and Fe3+in the brine. The dashed line is located at the half-wave potential of LFP / FP.

[0013] FIGs. 2A-B show cation uptake (adsorption capacity) from chemical reduction of FP using brines containing redox mediators. The redox mediators were (FIG. 2A) EDTA-Fe2+and (FIG. 2B) Citrate-Fe2+. All the salts were alkali metal chlorides, with each monovalent cation concentration of ~1 M in the mixed salt solution. The two temperatures of the chemical redox reaction were 23 °C and 75 °C, with the relevant condition of the reaction indicated on the figure.

[0014] FIGs. 3A-D show various graphs illustrating, in FIGs. 3A-B, (FIG. 3A) Uptake of Li+and Na+and (FIG. 3B) enduring chemical reduction of FP with redox mediators EDTA-Fe2+. FIGs. 3C-D show (FIG. 3C) Uptake of Li+and Na+and (FIG. 3D) a^aafter reduction by Citrate-Fe2+. The mediator concentration and reduction temperatures are indicated on the figures. The surrounding brine had a 1 :78 Li+:Na+molar concenration. FIG. 4 shows XRD patterns for the as-received LFP powder, FP after chemical oxidation of the LFP, and chemically reduced FP after contact with brine containing Citrate- Fe2+mediator and EDTA-Fe2+mediator, respectively. The mediator concentration was 0.1 M and the brine had 1 :78 Li+:Na+stoichiometry. The brine reaction was at 45 °C for a duration of 24 hours. Indexes were assigned based on PDF-01-078-7908 for LFP and PDF-04-017- 0610 for FP.28

[0015] FIGs. 5A-C show Li+(black circles) and Na+(red circles) adsorption capacity as a function of time for extraction of Li+from 1 :78 Li+:Na+brine using (FIG. 5A) EDTA-Fe2+and (FIG. 5B) Citrate-Fe2+. The extraction experiments were conducted at 45 °C. FIG. 5C shows selectivity factors (aa) for Li+extraction from EDTA-Fe2+(yellow-green circles) and Citrate- Fe2+(blue circles).

[0016] FIGs. 6A-B show Li+and Na+uptake into FP during multiple cycles of cation release via chemical oxidation of the powder and cation insertion via chemical reduction in brine with mediators (FIG. 6A) EDTA-Fe2+and (FIG. 6B) Citrate-Fe2+.

[0017] FIGs. 7A-F show SEM images of (FIGs. 7A and 7B) the initial FP prepared by chemical oxidation, (FIGs. 7C and 7D) FP after 5 cycles of Li+extraction / release using redox mediator EDTA-Fe2+, and (FIGs. 7E and 7F) FP after 5 cycles of Li+extraction / release using redox mediator Citrate-EDTA-Fe2+. Scale bars indicated in the images, where the top row is at relatively lower magnification relative to the bottom row of images.

[0018] FIGs. 8A-F show cyclic voltammograms at scan rates from 10 mV s1to 200 mV s1using brine solutions containing only added Fe2+ / Fe3+(chlorides) or the Fe with additional ligands. The mediator concentration was 0.1 M for all cases, and the mediators used were (FIG. 8A) EDTA-Fe2+, (FIG. 8C) Citrate-Fe2+, and (FIG. 8E) Fe2+ / Fe3+chloride. The total Fe concentration in each brine solution was equimolar of Fe2+and Fe3+. Cathodic peak currents as a function of square root of scan rate determined from the CV scans are also shown for (FIG. 8B) EDTA-Fe2+, (FIG. 8D) Citrate-Fe2+, and (FIG. 8F) Fe2+.

[0019] FIGs. 9A-B show the electrochemical reduction of FP to capture cations from a brine solution containing Li+:Na+molar ratio of 1 :78. FIG. 9A shows the cations adsorption capacity of the working electrode (FP) and FIG. 9B shows the selectivity factor a^afor FP was assessed after 50% extent of reduction of the FP at constant voltages of 0, -0.1 , -0.2, and -0.3 V (vs. Ag / AgCI).

[0020] FIGs. 10A-B show (FIG. 10A) Adsorption capacities for Li+and Na+and (FIG. 10B) selectivity factors a^ain FP from brine that more closely simulated Salton Sea brine (more specific composition in Table 3). Cation uptake was driven by the redox mediator indicated on the Figures. FIGs. 11 A-D show calibration curves for peak area as a function of prepared cation concentration using ion chromatography. (FIG. 1 1 A) Li+, (FIG. 1 1 B) Na+, (FIG. 11 C) K+, and (FIG. 11 D) Mg2+are shown, and the solutions containing these cations were prepared as chloride salts. The targeted solution concentration range was 0.01 to 1 mM.

[0021] FIG. 12 shows a cyclic voltammogram for LFP material as the working electrode measured in 1 M IJ2SO4 solution. The scan range was -0.25 V to 0.75 V (vs. Ag / AgCI), and the scan rate was 0.2 mV s-1.

[0022] FIG. 13 shows an XRD pattern (bottom) for FP reduced in 1 M of LiCI solution at 75 °C for 24 hours with 0.1 M of FeCI2and no added ligands. XRD for LFP (top) and FP (middle) powders also provided.

[0023] FIGs. 14A-B show XRD patterns for FP powder after chemical reduction at (FIG. 14A) 23 °C and (FIG. 14B) 75 °C in brine solutions using EDTA-Fe2+redox mediators. Each monovalent cation concentrations in the brine were ~1 .3 M, and the molar ratios of the cations in the brine are indicated on the Figures. At the top of each Figure are the patterns for the as received LFP powder and the LFP powder after chemical oxidation with hydrogen peroxide to form FP.

[0024] FIGs. 15A-B show XRD patterns for FP powder after chemical reduction at (FIG. 15A) 23 °C and (FIG. 15B) 75 °C using brine solutions containing Citrate-Fe2+redox mediators. Each monovalent cation concentrations in the brine were ~1 M, and the molar ratios of the cations in the brine are indicated on the Figures. At the top of the Figures are the patterns for the as received LFP powder and the LFP powder after chemical oxidation with hydrogen peroxide to form FP.

[0025] FIGs. 16A-B show example droplet images for determining contact angle of 1 :78 Li+:Na+brine solution with 0.1 M of redox mediator (FIG. 16A) Citrate-Fe2+and (FIG. 16B) EDTA-Fe2+on FP composite electrode.

[0026] FIGs. 17A-B show (FIG. 17A) Li+adsorption capacity and (FIG. 17B) Li+selectivity factor for this work (black stars) and previous reports (blue dots, from Refs. 2-13 of Example 1 Supplementary Information). In most cases, adsorption capacity was converted to the same unit basis from the reported data and Li+selectivity factor was recalculated based on Eq. 1.

[0027] DETAILED DESCRIPTION

[0028] In general, the present disclosure provides for methods and system for processing brine to obtain (e.g., capture and / or extract) lithium, and the like. Additional details are provided herein and in the Examples.

[0029] This disclosure is not limited to particular embodiments described, and as such may, of course, vary. The terminology used herein serves the purpose of describing particular embodiments only, and is not intended to be limiting, since the scope of the present disclosure will be limited only by the appended claims.

[0030] Where a range of values is provided, each intervening value, to the tenth of the unit of the lower limit unless the context clearly dictates otherwise, between the upper and lower limit of that range and any other stated or intervening value in that stated range, is encompassed within the disclosure. The upper and lower limits of these smaller ranges may independently be included in the smaller ranges and are also encompassed within the disclosure, subject to any specifically excluded limit in the stated range. Where the stated range includes one or both of the limits, ranges excluding either or both of those included limits are also included in the disclosure.

[0031] Embodiments of the present disclosure will employ, unless otherwise indicated, techniques of chemistry, material science, electrochemistry and the like, which are within the skill of the art. Such techniques are explained fully in the literature.

[0032] The following examples are put forth so as to provide those of ordinary skill in the art with a complete disclosure and description of how to perform the methods and use the compositions and compounds disclosed and claimed herein. Efforts have been made to ensure accuracy with respect to numbers (e.g., amounts, temperature, etc.), but some errors and deviations should be accounted for. Unless indicated otherwise, parts are parts by weight, temperature is in °C, and pressure is at or near atmospheric. Standard temperature and pressure are defined as 20 °C and 1 atmosphere.

[0033] Before the embodiments of the present disclosure are described in detail, it is to be understood that, unless otherwise indicated, the present disclosure is not limited to particular materials, reagents, reaction materials, manufacturing processes, dimensions, frequency ranges, applications, or the like, as such can vary. It is also to be understood that the terminology used herein is for purposes of describing particular embodiments only, and is not intended to be limiting. It is also possible in the present disclosure that steps can be executed in different sequence, where this is logically possible. It is also possible that the embodiments of the present disclosure can be applied to additional embodiments involving measurements beyond the examples described herein, which are not intended to be limiting. It is furthermore possible that the embodiments of the present disclosure can be combined or integrated with other measurement techniques beyond the examples described herein, which are not intended to be limiting.

[0034] It should be noted that, as used in the specification and the appended claims, the singular forms “a,” “an,” and “the” include plural referents unless the context clearly dictates otherwise. Thus, for example, reference to “a support” includes a plurality of supports. In this specification and in the claims that follow, reference will be made to a number of terms that shall be defined to have the following meanings unless a contrary intention is apparent. Prior to describing the various embodiments, the following definitions are provided and should be used unless otherwise indicated.

[0035] Definitions

[0036] The term "brine" as used in this description refers to fluid that include materials (e.g. salts) such as sodium, potassium, bromine, boron, lithium, iodine, magnesium, calcium, rubidium, iron, zinc, manganese, and sodium chloride (the exact materials depend upon the source of the brine), where the salt(s) (e.g., Li salt) is at a relatively high concentration as compared to seawater. While the example describes brines with relatively high dissolved solids content, such as those that may be found in geothermal fluids or produced water sources, the methods and systems of the present disclosure can be used with brines such as those described herein. Brine is a source of lithium. Brine sources of lithium include the salars in various parts of the world (e.g., Salton Sea Known Geothermal Resource Area, Andes Mountains of South America (lithium triangle) (e.g., the Salar de Atacama, Chile, Salar de Uyuni, Bolivia, and Salar de Rincon, Province of Salta, Argentina). The term "salar" as used in this description is understood as an evaporite from an area that accumulated and accumulates salts of hydrothermal and / or volcanic origin, and that in past geological eras was salt lakes with a high content of salts. Brine (also referred to as a “brine solution”) can be a natural brine, a synthetic brine, or a combination thereof, such as a continental brine, a geothermal brine, an oil field brine, a brine from hard rock lithium mining, or a combination thereof. Brines may be graded in terms of lithium concentration, for example varying between high grades of on the order of 0.3% Li to lower grades on the order of 0.03% Li (300 ppm or 300 mg / L) or even lower, where methods of the present disclosure can treat brines across the entire range of grades.

[0037] Discussion

[0038] The present disclosure provides for methods and systems for processing brine to obtain lithium. In an aspect, the present disclosure provides for methods that leverages intercalation materials and drives lithium capture using a chemical approach to treating brine as opposed an electrochemical approach, which can be complicated and expensive. The present disclosure provides for a chemical redox-driven approach to selectively extract Li+from brine. The present disclosure uses additives (e.g., Fe2+-compound and a complexing agent) to modify the redox potential of soluble iron(ll) compounds to be lower, such that they can drive reduction of an intercalation material (e.g., FePO4(FP)). After chemical reduction of the intercalation material, the Li+is captured by intercalation into the solid intercalated material. Subsequently, the Li+can be later separated from the Li+intercalated material through contact with a chemical oxidizing agent. Example 1 describes additional details. The present disclosure provides for chemically treating a brine solution to selectively obtain (e.g., capture, separate) lithium ions from the brine solutions. The brine solution includes lithium ions along with other cations and chemicals that make capturing the lithium ions complicated. The brine solution can include varying concentrations of lithium ions (and other species) depending upon the source of the brine solution. The brine solution can be obtained from continental brine, a geothermal brine, an oil field brine, a brine from hard rock lithium mining, or a combination thereof.

[0039] In an aspect, the method includes mixing the brine solution with an Fe2+-compound and a complexing agent to form a first mixture that includes the source brine (including Li+) and the Fe2+-complex, wherein the complex is a central metal (the Fe) bonded to one or more ligands (at least one of which is the additives added to the brine solution with and extra Fe salt). In an aspect, the only additives (those acting upon the brine or components in the brine) added to the brine are the Fe2+-compound and the complexing agent. The first mixture is not treated electrochemically. The brine solution and Fe2+-complex can be mixed for about 1 second to 1 hour. The first mixture can be adjusted (if needed) to a pH between 1 .0 to 12.0. In an aspect, the Fe2+-compound can be FeCh, FeSO4, Fe(NO3)2, or another source Fe salt where the Fe has a 2+ oxidation state, and mixtures thereof. The amount of Fe2+-complex can have a molar ratio of Fe2+: additive ranging from 20:1 to 1 :20 in the brine with the complex added. In an aspect, the complexing agent can be citrate, ethylenediaminetetraacetic acid, an agent / ligand that reduces the redox potential of the Fe2+ / Fe3+redox to below the potential of the intercalation material (e.g., FePO4), or a combination thereof. The amount of complexing agent can be about 0.001 to 0.4 moles per liter (M).

[0040] Next, the first mixture (brine solution with Fe2+complex) is contacted with a plurality of particles. In an embodiment, the first mixture can be flowed across or through a bed of the particles, where the first mixture and the particles are in contact for a period of time (e.g., about 1 second to 10 hours, about 1 to 10 minutes, or about 1 to 2 minutes), at a pH (e.g., about 1.0 to 12.0 or about 2 to 8), and at a temperature (e.g., about 10° C to 120° C or about 20° C to 75° C). In another embodiment, the particles can be added to a container that includes the first mixture and allowed to be in contact with one another for a period of time (e.g., about 1 second to 10 hours, about 1 to 10 minutes, or about 1 to 2 minutes ) at a pH (e.g., about 1.0 to 12.0or about 2 to 8), and at a temperature (e.g., about 10° C to 120° C or about 20° C to 75° C). In an aspect, the particles can have a d5o of about 0.5 m to 1 mm, 0.5 pm to 0.5 pm, or 1 pm to 15 pm. In an aspect, the particles can be agglomerated to form agglomerates that that have the dimensions that are a multiple of 2 to 100 the d5o of the particles used to form the agglomerate. The amount of particles present is sufficient to achieve the desired goal of capturing the lithium ions from brine due to reduction of the particles by the Fe2+-complex.

[0041] The particles include an intercalation material. In an aspect, the intercalation material includes a transition metal cation compound (e.g., intercalation Fe3+-solid material). Although not intending to be bound by theory, a portion of the transition metal cations in the intercalation material are reduced and Li+from the brine solution is intercalated into the intercalation material to form Li+intercalated material (and thus the Fe in the solution converts to Fe3+-complex). Other cations can also intercalate in the intercalation material, but these are not desired and the intercalation process is targeted to be selective to Li+. In an embodiment, the intercalation Fe3+-solid material be iron(lll) phosphate, iron pyrophosphate, an iron polyanion solids, or a mixtures thereof. The amount of the transition metal cation compound can be about 10 to 70 %, by volume, of the relevant vessel interior region containing the intercalation particles.

[0042] Subsequently, the particles and the first mixture can be separated from one another and the particles can be optionally rinsed to remove and excess brine residual from contacting the first mixture. In an aspect, the rinse can be performed using an aqueous solution (e.g., water).

[0043] The particles can then be contacted with an oxidizing solution for a period of time (e.g., about 1 second to 1 hour, or about 1 minute to 30 minutes), at a pH (e.g, about 1 .0 to 12.0 or about 2 to 8), and at a temperature (e.g, about 10° C to 120° C or about 20° C to 75°). While not intending to be bound by theory, upon mixing of the particles with the oxidizing solution, the Li+in the Li+intercalated material is released when the portion of the cations (e.g., Fe2+) are oxidized to reform the intercalation material with the transition metal in the oxidized form (e.g., Fe3+). In an aspect, the oxidizing solution can include a first oxidant and an acid. In another aspect, the oxidizing solution can include a second oxidant. In an aspect, the first oxidant can be a peroxide or an oxidant with a high enough potential to oxidize the intercalation material. In an aspect, the peroxide can be hydrogen peroxide, alkali peroxide, or a mixture thereof. In an aspect, the acid can be an acetic acid, sulfuric acid, hydrochloric acid, oxalic acid, or a combination thereof. In an aspect, the second oxidant can be a different Fe3+-complex, a soluble redox active species with high enough potential to drive oxidation of the intercalation material or a mixture thereof. In an aspect, the second Fe3+compound of the oxidizing solution is selected from the group consisting of: FeCI3, [Fe(CN)6]3', Fe2(SO4)3, Fe(NO3)3, and mixtures thereof. In an aspect, the second oxidant can be renewed and reused via regeneration from an electrochemical oxidation process.

[0044] In an aspect, capture of lithium from the brine has a selectivity factor (oa) of about 10 to 10,000. The selectivity factor is the ratio of the Li to other cations captured relative to the ratio of Li to all other cations in the initial feed brine. In an aspect, the intercalation material has an adsorption capacity of about 0.5 to 8.0 millimoles per gram of intercalation material. The adsorption capacity is the amount of cation (can be on specific cation basis like Li or total cation basis) captured in the intercalation material divided by the mass of the intercalation material.

[0045] The present disclosure also provides for a system to chemically treat the brine solution to capture lithium ions (Li+). The system includes a first structure, a flow system, and a second structure. The first structure is configured to mix a brine solution and Feecomplex constituents to form the first mixture including a Fe2+-complex. The first structure can be a container of an appropriate volume to mix the materials and can be made of stainless steel, glass, ceramics and the like. A flow system connects the first structure with the second structure and can include appropriate valves, flow meters, tubing, and the like. The second structure can be a container that is configured to mix the first mixture with the plurality of particles. The second structure can be a container of appropriate volume to mix the first mixture with the particles. The second structure can be a packed bed fixed reactor in fluidic communication with the first structure and the flow system. After contacting of the first mixture with the particles, the particles can be separated from, or kept separate from, the first mixture. The second structure is also configured to contact the plurality of particles with an oxidizing solution (which can be flowed to the second structure using the flow system and optional the first structure can be used to mix the components of the oxidizing solution). The Li+in the Li+intercalated material is released when the portion of the cations are oxidized to reform the intercalation material and a Li+-solution. The Li+-solution can be separated from the particles.

[0046] While embodiments of the present disclosure are described in connection with the Examples and the corresponding text and figures, there is no intent to limit the disclosure to the embodiments in these descriptions. On the contrary, the intent is to cover all alternatives, modifications, and equivalents included within the spirit and scope of embodiments of the present disclosure.

[0047] EXAMPLE 1

[0048] The worldwide demand for lithium (Li) continues to increase due to the increased use of Li-ion batteries in stationary energy storage and electric vehicles, coupled with their already ubiquitous use for powering electronic devices.1Li-ion batteries have a diversity of materials and compounds used for the cathode, anode, and electrolyte; however, in all combinations Li is required for these batteries.2According to the United States Geological Survey, the worldwide production of Li increased from 82,500 to 107,000 metric tons just between the years 2020 to 2021 , and Li-ion batteries accounted for 80% of Li end use.3New Li production sources and technologies are needed both to more sustainably support Li production and to prevent supply shortages. Supply limitations could cause Li prices to increase and impact the cost and adoption of energy storage technologies such as electric vehicles that are highly reliant on this resource and necessary for broader energy decarbonization.

[0049] Li resources naturally exist in ores, brines and seawater, and major supplies of Li are extracted from Li-bearing ores and brines. Methods of calcination, chlorination, and acid / alkaline chemical treatment are used to digest ores and extract out Li+.45Drawbacks of these methods include substantial consumption of chemical inputs, water resource consumption, and generation of waste. Conventional Li extraction resources and practices have also been associated with environmental conflicts related to the mining.6Another continental source for Li exists in brine solution. More recently, there has been interest in extracting Li+from brines which are less enriched in Li+and / or have complications of many other cations present, in many cases with many-fold excess concentration of other cations relative to the Li+.7The conventional method to extract Li+from brine solution is solar evaporation followed by precipitation, where the water is driven out of the solution and the end products are selectively precipitated compounds.8However, these methods require high energy inputs to drive out the water from the brine, which either needs to be provided by electricity / fuel input costs / processing or more commonly evaporation in solar ponds. The solar evaporation drying process can take from months to years. Other efforts to extract Li+from brines have included solvent extraction,9 10ion exchange adsorption,11electrochemical extraction,12and membrane separation.13Low cost, low environmental impact, and robust operation methods are needed for brine extraction.

[0050] Selective Li+intercalation materials have indicated promise for extraction of Li+due to their ion channels favorable for uptake of Li+.5 14Commonly used materials include lithium titanium oxides (e.g., Li4Ti50i2or Li2TiO3), lithium manganese oxides (e.g., LiMn2O4, Lii @Mni6O4), and lithium iron phosphate (LiFePO4, LFP).8 15These materials have generally been applied via methods of ion exchange adsorption, electrochemical driven separation, and redox mediated electrochemical separation.16In the ion exchange adsorption method, the intercalation materials were pretreated with acid to exchange the Li+with protons (H+). The H+are then substituted by the Li+when contacted with the Li+-containing brine, and the Li+is subsequently released in the next acid solution treatment cycle.17 18Ion exchange has required large amounts of acid as the proton sources. Additionally, the intercalation material needs to be stable in those acids. Manganese oxides in particular have known limitations with regards to dissolution of Mn when in contact with acid,1920thus processing must be controlled to minimize solid material dissolution and / or material modifications must be implemented to improve the structural stability.21 23For electrochemical-driven Li+separation, the cell contained a cathode, anode, and membrane separator, while the brine solution functioned as the electrolyte. During operation, Li+from the brine intercalated into an intercalation material within the anode or cathode depending on cell design, with the Li+later released into a separate recovery solution. The driving force for extraction and release was electrochemical potential, with external electrical energy needed for operation.2425Additional considerations for the process have included side reactions such as water splitting, and the addition of organic solvents to facilitate the electrochemical reactions.2627

[0051] Other processes have used intercalation materials for Li+extraction and have included chemical redox as the driving force for Li+insertion and extraction to the host material. For chemical redox, a potential difference is required between the solid intercalation materials and the soluble redox mediators. The most reported intercalation material for these chemical redox systems has been LFP. LFP is used as a commercial Li- ion battery cathode material with relatively low-cost constituent compounds, robust chemical stability, and reversible redox provided by the Fe. In prior work, the half-wave potential difference between LFP (0.186 V vs. Ag / AgCI) and potassium ferricyanide (K3Fe(CN)6) (0.247 V vs. Ag / AgCI) provided a thermodynamic driving force to oxidize the LFP.28Similarly, a redox mediator with a half wave potential below FP would be desired to thermodynamically drive the reduction of FP (and uptake of Li+). Previously, studies have shown that ligands bonded with iron could shift soluble iron redox potentials lower.29Even though LFP materials have ion channels well suited to Li+uptake and transport, some prior work has reported relatively low selectivity for Li+uptake in brine solutions with other cations when the relative Li+concentration was low, for example in solutions with excess competing sodium and magnesium.1430

[0052] In this example, we developed a chemical redox driven Li+extraction system using dissolved iron redox mediators to extract Li+from brine. The soluble iron was complexed separately with two different additives to the solution, ethylenediaminetetraacetic acid (EDTA) and citrate. The relevant redox couples will thus be referred to as EDTA-Fe2+and Citrate-Fe2+, respectively. The exact structures of the iron-ligand complexes were not investigated. The extraction brine was targeted to simulate solution conditions experienced for brine discharged after geothermal energy processing at the Salton Sea.31 32In the simulated brine, we evaluated the Li+extraction efficiency at high and low Li+concentration. In brine with high Li+concentration (1 M, molar ratio of Li+:M+ / 2+of 1 :1 , M+ / 2+=Na, K, Mg), FP had high selectivity to Li+in both EDTA-Fe2+and Citrate-Fe2+redox systems. The FP reached 94% conversion / reduction, and the Li adsorption capacity was ~6 mmol Li+g1FP for EDTA-Fe2+. In brine with lower Li+concentration (39 mM Li+, molar ratio of Li+:Na+, 1 :78), the extraction rate of the EDTA-Fe2+was faster than of the Citrate-Fe2+system, with the Li+adsorption capacity 3.8 mmol Li+g1FP with redox EDTA-Fe2+and 2.5 mmol Li+g1FP with redox Citrate-Fe2+. The selectivity factor aN''awas 78 for EDTA-Fe2+and 350 for redox Citrate-Fe2+. These chemical redox-driven systems for Li+extraction had high Li selectivity, high adsorption capacity, and no input energy consumption during the Li+capture step. The Li+selectivity difference between the two different redox couples facilitated by different soluble additives suggested control over relevant complexes and solution chemistry provided new routes improve Li+uptake and selectivity during mediated Li+capture with an intercalation material.

[0053] Materials and Methods

[0054] Brine chemicals

[0055] All salts and chemicals were purchased from Sigma Aldrich and used as received. The chemicals used were: sodium chloride (NaCI), lithium chloride (LiCI), potassium chloride (KCI), sodium hydroxide (NaOH), potassium hydroxide (KOH), lithium hydroxide (LiOH), iron (II) chloride tetrahydrate (FeCI2• 4H2O), iron (III) chloride (FeCI3), glacial acetic acid, magnesium chloride (MgCI2), potassium citrate, ethylenediaminetetraacetic acid (EDTA), N- Methyl-2-pyrrolidone (NMP), boric acid, manganese chloride tetrahydrate (MnCI2• 4H2O), and hydrogen peroxide solution (H2O2, 30%).

[0056] Preparation of FP

[0057] FP powder was obtained via a chemical oxidation method from prior literature.33In brief, 1 g of LFP (TOB, Xiamen) was mixed with a solution containing 5 mL of H2O2(30% in water, Fisher), 5 mL of acetic acid, and 140 mL deionized water. The LFP was added to the solution and allowed to react for two hours at room temperature without stirring. At the conclusion of the reaction process, the FP powder was collected and rinsed with DI water via vacuum filtration and dried in an oven with air at 80 °C overnight. Full delithiation was assessed using powder X-ray diffraction (XRD).28This LFP chemical oxidization method was also used to release the cations captured during the redox mediated extraction for analysis and will be referred to as the “chemical oxidation” method herein.

[0058] Preparation of FP and LFP electrode for electrochemical measurements

[0059] LFP and FP electrochemical electrodes were fabricated by blending LFP or FP powder with Super P carbon black (CB, TIMCAL) and polymer binder polyvinylidene fluoride (PVDF, Alfa Aesar) at mass ratio of 8:1 :1 LFP:CB:PVDF. The powder mixture was blended using mortar and pestle by hand and wetted with NMP solvent. The resulting slurry was further mixed using an Ar- 100 Thinky mixer. For LFP electrode used in cyclic voltammetry measurement, the slurry was pasted on the tip of a 5x1 cm nickel foam strip by hand. For LFP or FP electrodes used in electrochemical cells, the FP or LFP slurry was pasted on titanium rod with diameter of 0.046 mm (McMaster-Carr). For FP electrodes used in contact angle measurements, FP composites were fabricated by casting FP slurry onto an aluminum foil with a doctor blade with gap height of 100 pm. After pasting the FP or LFP slurry onto the metal current collectors, the electrodes were dried at 80 °C overnight in air in an oven to remove any residual NMP solvent.

[0060] Electrochemical characterization of FP and redox mediators

[0061] Cyclic voltammetry (CV) was used to determine the half-wave potential of redox mediators (EDTA-Fe2+& Citrate-Fe2+), LFP / FP, and Fe2+ / Fe3+solution. CV experiments were conducted with a Bio-Logic SP50. Separate platinum disc electrodes with diameters of 1 .6 mm were used as the working and counter electrode. Ag / AgCI with saturated KCI was used as the reference electrode (Pine Research). Solutions were scanned at rates ranging from 10 mV S’1to 200 mV s1. In the testing solutions, in addition to the evaluated redox species, 3.04 M of NaCI and 39 mM of LiCI was also added which was a level intended to mimic the relevant concentrations in Salton Sea brine.32Nitrogen gas was bubbled through the solutions during the entire preparation process to minimize air oxidation of Fe2+species. For the half wave potential of EDTA-Fe2+, EDTA and FeCI2with each 0.1 M were dissolved in 1 M acetic acid solution, and LiOH was added to dissolve the EDTA as well as to adjust the solution pH to 7. The scanned voltage range was from -0.5 V to 0.2 V vs. Ag / AgCI. For Citrate-Fe2+, equimolar of potassium citrate and FeCI2(0.1 M each) was dissolved in DI water, the scanned voltage range was from -0.5 V to 0.4 V vs. Ag / AgCI. For the half wave potential of Fe2+ / Fe3+(from chloride salts), equimolar of Fe2+ / Fe3+(0.05 M each) was prepared in DI water, and the scanned range was from 0 V to 1 .2 V vs. Ag / AgCI. For using CV scans to calculate the diffusion coefficient of Fe solution redox species, the cathodic peak current was plotted against the square root of scan rate, and the slope of a linear least squares fit was extracted for the calculation.34

[0062] For CV evaluation of LFP active material, the working electrode was prepared with procedures described in Section 2.3. The LFP electrode loaded within nickel foam was dipped in 1 M Li2SO4 solution and used as the working electrode and platinum wire (1 mm) was the counter electrode. The scanned voltage range was -0.25 - 0.75 V (vs. Ag / AgCI), and the scan rate was 0.2 mV s1.

[0063] Simulated brine preparation and extraction procedures

[0064] Simulated brine with high Li+concentration, but relatively low competing cations concentration (molar ratio Li+:M+ / 2+, 1 :1), and low Li+concentration, but high competing cations concentration (molar ratio Li+:Na+, 1 :78), were both evaluated for Li+extraction. In all the experiments with simulated brine solutions, the molar ratio of FeCI2and ligands (EDTA, Citrate) was kept at 1 :1 for all evaluated concentrations. To avoid Fe oxidation, FeCI2was added to the brine solution last, and after all other chemicals had been dissolved. The brine solution was also at all times purged with nitrogen gas via bubbling through a sealed flask. All pH values reported were after the step of ligand addition / dissolution. The Li+extracting FP solid (prepared from chemical oxidation of LFP), was contacted with brine solution in a porous pellet form. FP pellets were prepared by loading the FP powder into a 13 mm diameter pellet die and hydraulically pressing at 12,000 psi for 2 mins at room temperature. FP was in pellet form for convenience of collection after extraction.

[0065] Simulated brine preparation with high Li* concentration and relatively low competing cation M* / 2* concentration (molar ratio Li*:M* / 2*, 1 :1)

[0066] The detailed compositions of brine solutions with the two redox mediators can be found in Table 1. The total mass of the brine solution was 100 g, and the mass of each chemical was calculated based on assumed volume of 100 ml_ of brine solution. The procedure for preparing simulated brine solution with redox mediator, EDTA-Fe2+is described in the following: 1) 1 M of acetic acid solution was prepared to mitigate dramatic solution pH swings. 2) 0.1 M of EDTA was added to the acetic acid solution. 3) Subsequently, LiOH was added to the solution mixture to facilitate EDTA dissolution and until the solution was clear. Additional LiOH was added, if needed, until the solution pH reached 7, and the corresponding concentration of Li* was ~1 .3 M. 4) Target cation chloride salt (NaCI, KOI, MgCI2) with the same molar concentration as the Li+were added and dissolved in the solution. 5) 0.1 M of FeCI2• 4H2O was added last and the solution was stirred until all material dissolved. Nitrogen gas was bubbled in the solution throughout the process. 6) 0.5 g of FP pellet was directly added to the solution. For preparing brine solution with redox mediator Citrate-Fe2+solution, the experimental procedures were similar to the EDTA-Fe2+solutions, except that 1) No acetic acid was added to the solution due to its chelation competition with citrate ligand, 2) There was no need for solution pH adjustment as the citrate ligand had high solubility in water without pH moderation, 3) LiCI was added as the Li source. Experiments with high concentration of Li+were conducted for 24 hours at 23 °C and 75 °C. Duplicate experiments were completed for each brine solution, and average adsorption capacity were reported for each redox mediator.

[0067] Studies of simulated brine with low Li* concentration and high competing cation Na* concentration (molar ratio Li*:Na*, 1 :78)

[0068] A simulated brine with a Li*:Na* molar ratio of 1 :78 was studied to mimic the concentrations of those two cations in real brine from Salton Sea.31 32The procedures to prepare the brine solution were similar to those described above, except that the competing cation was specifically Na*, and the concentration of Li* and Na* were adjusted to 39 mM and 3.04 M. All salts appeared dissolved based on visual inspection, without residual solids present. The detailed composition of brine solutions with low Li* concentration and high Na* concentration (Li*:Na* molar ratio of 1 :78) can be found in Table 2. Experiments were also conducted where the redox mediator concentrations were varied for solutions that were contacted with the FP. The concentrations of EDTA-Fe2+were 15, 30, 50, 100, 200, and 300 mM. The same concentrations were also evaluated with Citrate-Fe2+, except 100 mM was the highest concentration evaluated for Citrate-Fe2+due to solubility limitations. The Li+extraction process for each redox concentration was also assessed at three different temperatures, 23°C, 45 °C, and 75 °C. For experiments at elevated temperatures (45 °C and 75 °C), the FP was added after the brine solution reached the target temperature via heating in an oil bath. The extraction reactions all proceeded for 24 h.

[0069] Experiments were also conducted to assess the time progression of the Li+extraction process. To compare the two mediators and because of citrate solubility limitations at higher concentrations, the condition chosen for comparison was 100 mM redox mediator (EDTA- Fe2+or Citrate-Fe2+) and 39 mM of Li+and 3.04 M of Na+(consistent the ratio of those cations at the Salton Sea). The temperature was 45 °C, and the reaction was allowed to proceed in independent experiments for intervals increasing in timescale of 2 h up to 24 h. Each of these experiments had nominally identical solution and FP preparation, where only the reaction timescale was varied. After the chemical reduction of FP by the solution had concluded, the FP powder was washed several times with deionized water and dried at 80 °C in an oven in air overnight.

[0070] Experiments were also conducted to evaluate multiple Li+capture and release cycles for the FP / LFP solid. FP was reduced by simulated brine solution (39 mM of Li+and 3.04 M of Na+) with 100 mM added redox mediator. The reaction proceeded at 45 °C for 24 h for each capture cycle. After the reduction and cation capture step, the cations were released via the chemical oxidation method (described earlier). The solution supernatant was collected for ion chromatography (IC) analysis, and the FP powder was collected and rinsed with DI water before the next round of reduction via contact with fresh brine solution. The capture and release process were repeated for 5 complete cycles. Both Citrate-Fe2+with EDTA-Fe2+redox mediators were evaluated. Capture / release cycling experiments for each redox mediator were repeated twice and the average values were reported. The FP powder collected after 5 cycles of extraction / release of both redox mediators, and the initial FP prepared by chemical oxidation, imaged using scanning electron microscopy (SEM) imaging. The SEM was a FEI Quantum 650.

[0071] Lithium extraction from brine with iron only

[0072] A control experiment of lithium extraction with only iron (no ligands added) was conducted. 0.1 M of FeCh and 0.5 g of FP pellet was added to brine solution containing 1 M LiCI. The extraction reaction continued for 24 hours at 75 °C. At the end of reduction, the cations captured in FP were released via the chemical oxidation method and the concentration of the cations was quantified by IC.

[0073] Brine solution with composition more representative of Salton Sea extraction

[0074] A brine solution with composition more representative of the Salton Sea brine was used for Li+extraction using FP with redox mediators, EDTA-Fe2+or Citrate-Fe2+. The detailed brine compositions can be found in Table 3. Due to solution stability limitations, Ca2+was omitted from the original formulation during the brine preparation. Based on previous experimental optimizations of the reaction temperature and redox mediator concentration, the reduction reactions were continued for 24 h at 45 °C using 0.1 M of redox mediators. The captured cations were then released by chemical oxidation for IC analysis. Electrochemical Li+uptake from brine

[0075] Electrochemical Li+extraction from brine was achieved using a three-electrode system, where the working and counter electrode were FP and LFP, respectively, and Ag / AgCI with saturated KCI was the reference electrode. The preparation of LFP / FP electrodes were mentioned in “Preparation of FP and LFP electrode for electrochemical measurements” (above), and the mass ratio of working (FP) to counter (LFP) electrode was at 1 :1.5 (e.g., there was excess capacity and Li+in the LFP electrode). The three electrodes were submerged in a salt solution with a Li+:Na+molar ratio of 1 :78 dissolved in DI water. The voltage was kept constant throughout the lithiation process, with the potential below that necessary for intercalation of Li+into the FP (0.242 vs. Ag / AgCI). The constant voltage used for lithiation of the FP ranged from -0.3 V to 0 (vs. Ag / AgCI), with increments of 0.1 V. For comparison with chemical lithiation, the electrochemical lithiation was stopped after achieving 50% extent of lithiation. The extent of lithiation was calculated based on the mass of active material (FP) and assumed theoretical capacity of 165 mAh g-1FP. Upon completion of the extraction process, the working electrode (originally FP) was rinsed with 200 mL of DI water to remove residual salt solution. The captured cations from the solid electroactive material in the working electrode were released using the same chemical oxidizing method as for FP preparation.33The supernatant of the solution after oxidation of the powder was collected for IC composition analysis. Duplicate experiments were conducted, and the average value was reported. Quantification of Li+uptake

[0076] Progression of FP reduction was quantified using analysis of powder XRD patterns. XRD measurements were collected using Cu Ka radiation (wavelength=1 .5406 A) with a PANalytical X'pert ProMPD. To quantify the compositions of cations intercalated into the FP during reduction, the powder sample was chemically oxidized to release the cations to the solution phase. The chemical oxidation method was the same used to prepare the initial FP powder from the purchased LFP, and followed methods published previously.33After two hours of oxidation, the solution supernatant was collected and filtered to separate the solid powder from the suspension. Solution composition was then determined using IC. Samples for IC analysis were all 5 mL, and these solutions were used to flush and fill a 25 pL injection loop, which was used to inject the sample for analysis using a Thermo Scientific Dionex ICS- 2100. IC calibration standards were prepared by using known concentrations of metal chloride salts ranging from 0.01 mM to 1 mM total cation concentration with 1 :1 molar ratio of Li+and target cation M+ / 2+(M: Na+, K+, Mg2+). Linear calibration curves of peak areas as a function of concentrations were obtained, which was later used to determine the Li+and M+ / 2+concentrations in evaluated samples. Cations in collected samples were diluted several folds using DI water to attain concentrations within the range of the calibration standards. Example calibration curves can be found in FIGs. 11A-D.

[0077] Contact angle between FP composite electrode and brine solutions with redox mediator The contact angle between FP composite electrode and brine solutions was measured with a Goniometer (Rame-Hart Instrument). The brine solution was 39 mM of Li+and 3.04 M of Na+, with 100 mM added redox mediator. A 12 ul size brine droplet was dropped onto a FP composite electrode and the contact angle between the brine droplet and substrate was measured for 20 times consecutively with 0.5 s intervals, and the average angle of those measurements was calculated. The contact angle was extracted through image analysis provided with the equipment software. Fabrication of FP composite electrode was described in “Preparation of FP and LFP electrode for electrochemical measurements” (above). Three different regions of the FP composite electrode were used as the substrate, and a total of 10 droplets of each brine solution were evaluated for their contact angles. The reported contact angle was the average of these 10 droplets, with the reported uncertainty the standard deviation of the ten droplet measurements for each brine solution.

[0078] Results and Discussion

[0079] Electrochemical characterization of iron redox mediators

[0080] The intercalation material used to uptake Li+in this study was FP, which has a halfwave potential of 0.242 V vs. Ag / AgCI, as measured from CV (CV scans can be found in FIG. 12). The half-wave potential was obtained by taking the average of the cathodic peak potential and the anodic peak potential.35This half wave potential was equivalent to 3.469 vs. Li7Li, and was consistent with the reported half-wave potential of FP.28To drive Li+from the brine solution into the FP, the Fe3+in the FP must be reduced to Fe2+. To at minimum provide sufficient thermodynamic driving force towards FP reduction, the potential of the redox mediators in the surrounding electrolyte solution must be below the redox potential of FP. Without added ligands to the chloride brine solution, the half-wave potential of Fe2Fe3+ was 0.504 V vs. Ag / AgCI from CV scans (FIG. 1). This value was above that observed for FP / LFP, and thus would not drive FP reduction, at least within a reasonable range of relative state of oxidation / reduction of the soluble Fe. For example, when the solution was 99% Fe2+, the estimated Nernst potential was 0.385 V and still insufficient to reduce FP. Some prior studies have reported that when aqueous iron cations were complexed with specific ligands, the redox potential shifted much lower. For example, the redox potential of Fe was shifted to -1 .04 V vs. Ag / AgCI,360.146 V vs. Ag / AgCI,37and 0.865 V vs. Ag / AgCI3839when bonded to ligands from triethanolamine (TEA), sodium oxalate, and phenanthroline, respectively. Introduction of the ligands needs to also consider the chemical compatibility with the intercalation materials and other components in the brine solution. For example, the reported TEA-Fe solution was highly basic (pH >10), and LFP / FP has previously been reported to have stability limitations when exposed to high pH solutions.5Sodium oxalate has relatively low solubility in water, and even lower in relevant brine solutions - which makes achieving necessary redox mediator concentrations with the ligands to keep the reaction rate high a challenge. In this study, to shift redox potential of Fe27 Fe3+to be lower than FP, two additives were introduced into the brine solution, EDTA and Citrate. The half-wave potentials of redox mediators (Citrate / EDTA-Fe2+) were determined using CV (FIG. 1). From CV it was determined that the half wave potentials for the redox mediators EDTA-Fe2+and Citrate-Fe2+were -0.07 V and -0.1 1 V vs. Ag / AgCI, respectively, in salt solution with 39 mM of Li+and 3.04 M of Na+. The reversibility of EDTA-Fe2+was higher than Citrate-Fe2+based on the more symmetric CV curves.37 40The potential difference (AE) between the cathodic / anodic peak of the two redox mediators (EDTA-Fe2+, Citrate-Fe2+) were 0.12 V and 0.67 V. The higher the (AE) suggested the redox reaction of Citrate-Fe2+was slower.37Li+extraction from simulated brine with high Li+concentration

[0081] Before using redox mediators EDTA-Fe2+and Citrate-Fe2+to extract Li, a control was conducted using Fe2+to extract Li from solution. Assuming the Fe solution potential follows behavior consistent with the Nernst equation, at least a small amount of reduction of FP would be expected before there would be insufficient thermodynamic driving force to drive the redox forward. FP contacted a brine solution that had 1 M of LiCI, and was reduced by 0.1 M of Fe2+at 75 °C. After 24 hours reduction, the lithium uptake by FP was about 0.4 mmol g1FP. This relatively low uptake was also reflected in the XRD measurements of the powder immediately after the control experiment. There was not a large peak at 17.4°, which would be expected if the powder had converted to LFP (FIG. 13).

[0082] EDTA-Fe2+and Citrate-Fe2+had half-wave potentials below FP with magnitudes of 0.312 V and 0.352 V, respectively, which would thermodynamically drive FP reduction and Li+insertion into the solid crystal structure. Chemical redox mediators were first evaluated to extract Li+from simulated brine that contained only single monovalent cations (Li+or Na+), and then Li+with competing cations (Na+, K+, Mg2+) at 1 :1 molar ratio. In the simulated brine, each ion concentration was ~1 M, and each extraction continued for 24 h. Duplicate experiments were conducted with average values reported in FIGs. 2A-B.

[0083] In the EDTA-Fe2+chemical redox extraction system, when the brine had Li+only, the adsorption capacity of Li+was 5 mmol Li+g-1FP at room temperature, and 6 mmol Li+g-1FP at 75 °C (FIG. 2A). The intercalation material achieved 94% conversion by Li+uptake based on a theoretical full conversion of 6.4 mmol Li+g1FP (assuming one Li+per Fe in FP). In the brine with Na+only, there was slight Na+uptake at room temperature, which may have been due to the relatively higher activation energy barrier for intercalation of Na+.21However, at elevated temperature, and without Li+in the solution, about 4 mmol Na+g1FP was inserted into FP. This suggested the higher temperature aided Na in overcoming the activation barrier for insertion into FP. Next, the redox mediator was evaluated in brine solutions that had equimolar of Li+with competing cations, Li / M+ / 2+(M+ / 2+= Na+, K+, Mg2+). In the case of equimolar of Li+to K+, at both room and elevated temperatures minimal K+intercalated into the FP, and the K+also did not interfere with Li+adsorption as the Li+adsorption capacity was similar to the adsorption capacity in brine only containing Li+cations. Hence, K+was considered as a spectator ion and was consistent with reported literature.41For the brine solution containing Li+and Na+, as Na+competed with Li+, the Li+adsorption capacity was slightly lower than the adsorption capacity in the brine only containing Li+cations at 23 °C. There was low Na+insertion at both 23 and 75 °C. This outcome may have been due to the high mobility of Li+at high temperature as well as the high accessibility of Li+near the FP interface due to the high Li+concentration. In addition, the activation energy barrier for Li+was expected to be lower than Na+.21In brine solution with equimolar of Li+ / Mg2+, the Li+adsorption capacity was significantly reduced to 2.3 mmol g1FP at room temperature. The reduced Li+uptake might be caused by the formation of rigid network of Mg2+at the FP particle surface that possibly negatively influenced Li+access and mobility to the storage sites in FP.41At 75 °C, Li+adsorption capacity was similar to the capacity achieved in brine only containing Li+cations. This may have been due to the improved Li+mobility at the elevated temperature coupled with the higher solid phase diffusion barriers for the higher charge density Mg2+.

[0084] In Li+extraction with redox mediator, Citrate-Fe2+, the Li+adsorption capacity was ~5 mmol g1FP at both 23 °C and 75 °C in brines only containing added Li+cations, which was 78% conversion of the FP material based on the theoretical adsorption capacity by Li+uptake (FIG. 2B). In the Na+only brine solution, the Na+insertion was less than 0.2 mmol Na+g1FP at 23 °C but increased to ~1 mmol Na+g1FP at 75 °C. In brines with 1 :1 molar ratio of Li+to competing ions (K+, Mg2+), the Li+adsorption capacity was 4 mmol Li+g1FP at 23 °C, and 0.5 mmol Li+g1FP higher at 75 °C. The detected cations concentrations of K+, Mg2+were both lower than 0.3 mmol g1FP and thus were not inserted into FP with high uptake. In brines with competing Na+, the Li+adsorption capacity reduced to 2.8 mmol Li+g-1FP at room temperature, but achieved a similar adsorption capacity as it obtained in other examined brines at elevated temperature. The elevated temperature and high Li+accessibility promoted Li+adsorption.

[0085] Overall, in the synthetic brine solution with high Li+concentration (~1 M) and with comparable concentration of competing cations (Li+:M+ / 2+, 1 :1), both redox mediators had high uptake and selectivity to Li+as opposed to the competing cations M+ / 2+(M+ / 2+=Na+, K+, Mg2+) at both room and elevated temperatures. The Li+adsorption capacity for EDTA-Fe2+mediator was slightly higher (~1 .5 mmol Li+g1FP) than for Citrate-Fe2+at 75 °C. This might be due to the faster electron transfer rate of the EDTA-Fe2+complex. A previous study reported that the electron transfer rate constant of Fe2Fe3+with citrate ligand was 9.33 X 105cm S’1while it was 3.4 X 102cm s1with EDTA ligands, although the surrounding solution was not equivalent to this work.2942The rate constant of EDTA-Fe27Fe3+was three orders of magnitude higher than for Citrate-Fe27Fe3+. K+was consistent with a role of a spectator ion in both chemical redox brine systems because it neither competed with Li+for storage sites nor interfered with the total Li+intercalation at both 23 °C and 75 °C, and that was consistent with a previous report that used an electrochemical driving force to extract Li+from brine.41However, Mg2+lowered the Li+intercalation capacity significantly at room temperature for redox mediator EDTA-Fe2+, even though the amount of Mg2+detected in FP was minimal. The lowered capacity may have resulted from slower charge transfer of Li+induced by Mg2+. Mg2+has been reported to form a rigid network in the electrical double layer at the surface of FP particles that can impede Li+charge transfer.41However, the reduction of Li+uptake in the presence of Mg2+was not observed for the same brine compositions with redox mediator Citrate-Fe2+. We speculated that the slower diffusion of redox complex and electron transfer rate allowed sufficient time for Li+concentration to replenish near the particle interface, which outweighed the impact of slower Li+charge transfer caused by Mg2+. In addition, this higher Li adsorption capacity might be attributed to the chelation between Mg2+and citrate that restricted the interference of Mg2+to Li+intercalation. Improvement of Li adsorption was not observed in redox mediator EDTA-Fe2+system, which suggested the chelation of EDTA-Mg2+was relatively low. In addition, the elevated temperature only resulted in slightly increasing Li+adsorption for most of the brine solutions for Citrate-Fe2+, which was in contrast with observations for redox mediator EDTA- Fe2+. Even though temperature increased the mobility of cations and would have been expected to facilitate mitigating mass transport limitations for cations near the surrounding particle surface, the relative impacts of the temperature on other factors such as the redox reaction kinetics would also be impacted by temperature. A more detailed assessment of temperature effects and the influence on relevant steps in the Li capture process will be the subject of future investigations.

[0086] Li+intercalation and conversion of FP to LFP was also consistent with XRD patterns of the powders after contact with brine solutions containing the redox mediators (patterns can be found in FIGs. 14A-B and FIGs. 15A-B). The XRD peak associated with the

[0200] index at 17.4° indicated the formation of LFP. There was also no observed broadening of the peaks in the brine solution with mixed cations relative to the brines only containing Li+, which suggested no or minimal impurity cations insertion (i.e. , other than Li+). These XRD results were consistent with the IC analysis of low amounts of competing cation insertion. The correlation of peak broadening and impure cations insertion were also corroborated with samples reduced in Na+only solution, where the XRD peaks appeared broader and shifted. Li+extraction from simulated brine with low LF concentration and high Na+concentration

[0087] After assessing the selectivity for Li+uptake from brines with relatively high Li+concentration, next the Li selectivity during chemical reduction of FP in the presence of solutions with low Li+concentration (both absolute and relative) but high concentration of competing cations was evaluated. The Li+concentration in the brine was 39 mM, while the Na+concentration was 3.04 M, resulting in a molar ratio of Li+:Na+of 1 :78. These absolute and relative Li+and Na+concentrations were intended to match reports for discharged brine used to produce geothermal power at the Salton Sea.32 43Since K+was a spectator ion in brine solution, potassium-based salts were used to aid dissolution of EDTA in preparing redox mediator EDTA-Fe2+and potassium citrate was used as the ligand source for redox mediator Citrate-Fe2+. For both redox mediators, we evaluated the effect of mediator concentration and reaction temperature on Li+uptake into the FP (adsorption capacity) and selectivity to Li+.

[0088] For the EDTA-Fe2+system at 23 °C, the Li+adsorption capacity increased with increasing redox EDTA-Fe2+concentration from 15 mM to 30 mM (FIG. 3A). When the redox concentration was above 30 mM, more redox mediators were expected to provide more driving force that pushed more cations into the FP crystal structure. However, it resulted in more Na+insertion rather than Li+at concentration range from 50 mM to 300 mM. This outcome may have been due to the high accessbility of Na+surrounding the FP particle, such that more available Na+was pushed into the FP crystal structure to accomdate the charge state change of EDTA-Fe2+. As more Na+occupied the crystal lattice of FP, the amount of Li+intercalation reduced, and the corresponding Li selectivity decreased. Selectivity factors aa, were calcuated to quantify the extent of selectivity for Li+to Na+extraction from the brine into the FP solid during the chemical redox. The equation to calculate aais shown in Eq. (1).544 (1) where, CLiand CNaare the concentration of Li+and impurity cations (Na+) in the solid materials; and CL'iand CN'aare the concentration of Li+and impurity cations (Na+) in the brine source. Higher values of a^amean improved separation of Li+and Na+from brine. The calculated for various concentrations of EDTA-Fe2+can be found in FIG. 3B. The aadecreased from 1200 to 19 at the corresponding concentrations from 15 mM to 300 mM. The lower redox concentration might result in slower transport of redox mediator near the particle surface, which allowed sufficient time for Li+to distribute around the FP particle surface, thereby enhancing Li+selectivity. Thus, increasing redox EDTA-Fe2+concentration to above 50 mM was not necesssary for improving both the Li extraction capacity and Li selectivity. Similar trends for Li+and Na+adsorption were also observed at higher reaction temperatures of 45 °C and 75 °C, and higher reaction temperatures resulted in more intercalation of both Li+and Na+. At 75 °C, the maximum amount of Li+intercalation was 1 mmol Li+g1FP more than that at room temperature. The Na+insertion was also increased, such that the highest amount of Na+insertion was ~2.3 times that observed at room temperature. The a^laat elevated temperatures were lower than at room temperature because higher temperature increased the mobility of both cations.

[0089] In the extraction system with redox mediator Citrate-Fe2+, the Li+adsorption capacity increased as a function of mediator concentration for each of the temperatures evaluated (FIG. 3C). Due to solution stability / solubility limitations, the highest mediator concentration evaluated was 100 mM. The amount of Na+insertion was less than 0.5 mmol g-1FP across all conditions assessed. The a^lastabilized around 300 (FIG. 3D), except for the highest a^laat reduction with 15 mM Citrate-Fe2+at 23 °C. We speculated that this high Li+selectivity was a combined outcome of the slow redox reaction, slow diffusion coefficient of the redox mediators, and low reaction temperature in the brine solution. The slow reaction enabled enough time for the cations around the intercalation material surface to be redistributed for Li+to have more accessbility to the ion channels of the FP. Both low temperature and low redox concentration are beneficial for this slow redox reaction, and thus the aareached its highest value at 800. Although at high redox concentration the Li selectivity for Citrate-Fe2+was higher than for EDTA-Fe2+, the total Li+uptake did not achieve the highest values observed for the EDTA-Fe2+, where the highest uptake achieved was still lower than EDTA- Fe2+by ~1 mmol Li+g1FP. This outcome may have been due to the lower electron transfer rate constant of redox mediator Citrate-Fe2+. Hence, when redox concentration was lower than 50 mM, reduction with mediator EDTA-Fe2+had advantages of high Li adsorption capacity and Li selectivity, where for other solution conditions Citrate-Fe2+was more preferred.

[0090] For Li+capture experiments with large amounts of competing Na+insertion, the Na+insertion into the FP influenced the resulting XRD patterns for the powders after contact with the brine (FIG. 4). From the IC results (FIGs. 3A-D), Na+insertion into FP samples was greater for FP reduction conditions of high mediator concentration and high temperature. For comparison between the two mediators with identical temperature and mediator concentration during reaction, reacted FP powder after contact with brine at 45 °C and with mediator concentration of 0.1 M was evaluated using XRD. The XRD patterns were also evaluated in the context of the as-purchased LFP and the FP after chemical oxidation of the initial LFP powder. The

[0200] peak has previously been used to indicate the conversion between LFP / FP, as for LFP this peak is at 17.4° and for FP this peak is at 18.3°.28A partial reduction of FP will result in the presence of both the LFP and FP

[0200] peaks, where the relative peak magnitude can be correlated to the relative amounts of each phase present.28The XRD peaks suggested that both redox mediators EDTA-Fe2+and Citrate-Fe2+drove nearly full conversion of FP, where the Citrate-Fe2+had a slight FP peak remaining and the EDTA-Fe2+did not have a discernable FP peak remaining (FIG. 4). The two different mediators had noticeable differences in the peak broadening for the resulting powders. There was a greater extent of peak broadening for the powder reduced using EDTA-Fe2+. The EDTA-Fe2+processed powder also had more Na+intercalated into the material structure. Thus, the XRD was consistent with the lower selectivity when the EDTA-Fe2+mediator was used, as the larger Na+can introduce more strain and peak broadening in the crystal structure.14Both the XRD and IC results were consistent with high conversion / reduction of the FP when in contact with both redox mediators, and higher selectivity for Li+over Na+for the Citrate-Fe2+mediator. Li+adsorption time evolution and uptake / release cycling

[0091] As discussed above, the Li+uptake and selectivity were different depending on the redox mediator used to reduce the FP. However, the analysis above was only conducted on solid material after the conclusion of a chemical redox process, typically after 24 hours. As Li+adsoprtion rate needed to be different to achieve the final uptake values, experiments were also conducted to assess the Li+adsorption over time. Cation capture via chemical redox was conducted where the FP / LFP powder was collected after fixed times of reaction. Fresh synthetic brine solutions with 1 :78 Li+:Na+and both 0.1 M of EDTA-Fe2+and Citrate- Fe2+mediators were prepapred for each set of experiments to assess how uptake and selectivity varied with FP contact time with the brine at 45 °C. The contact intervals for the FP were in 2 hour increments up to a maximum of 24 hours. The resulting Li+uptake as a function of reaction time for each redox mediator can be found in FIGs. 5A-C. The total amount of Li+extracted from brine was similar for both mediators at the conclusion of 24 hours reaction, at about 2.5 mmol Li+g-1FP. The Li+adsorption into the FP plateaued at a shorter reaction time for the EDTA-Fe2+system, after roughly 10 hours. However, the Citrate-Fe2+mediated reaction took about 15 hours of reaction time to plateau. Even though Li+uptake rate was higher at the initiation of the chemical redox process for the EDTA-Fe2+mediator, Li+selectivity was relatively lower compared to Citrate-Fe2+mediator, with Li+and Na+inserted at nearly 1 :1 molar ratio. For EDTA-Fe2+mediator, at later times more Na+was inserted compared to Li+(FIG. 5A). For chemical redox using Citrate-Fe2+mediator, the total Na+uptake was fairly constant with reaction time. After initial insertion of approximately 0.5 mmol Na+g1FP, the intercalation of Na+appeared to stop, and only Li+uptake increased with reaction time until plateauing at about 15 hours (FIG. 5B).

[0092] Selectivity factors were calculated for these two extraction systems using Eq. 1 to compare the separation of Li+and Na+from brine. As shown in FIG. 5C, over the initial 10 h of the Li+extraction from the brine into the FP, the aausing Citrate-Fe2+was slightly higher than for the EDTA-Fe2+. Beyond 10 hours extraction, the a^afor the Citrate-Fe2+greatly exceeded the EDTA-Fe2+, with the relative improvement being approximately a factor of 6. This outcome suggested that the choice of redox mediator provided a method to control and improve the selectivity for Li+uptake into the FP. In previous reports, electrochemically- driven extraction of Li+from brine into FP introduced a controlled pulsing mode to allow the cations to redistribute around the FP particle surface, which increased Li+accessbility to the FP ion channels.30In another report, researchers showed that a polymer coating on FP can modify Li+selectivity, and attributed the selectivity improvement to modifications in particle surface hydrophilicity.14The surface hydrophilicity of intercalation material can be assessed in the context of contact angle measurements.14A more hydrophilic surface tends to have lower contact angle with water droplets, and in prior reports more hydrophilic surfaces have resulted in more adsorption capacity,23while a higher contact angle increased the selectivity of Li+to Na+.14In this work, 12 pl drops of brine solution (Li+:Na+, 1 :78) with either EDTA-Fe2+or Citrate-Fe2+mediators were placed onto FP composite electrodes, and optical images of the brine drops were recorded (example contact angle images can be found in FIGs. 16A-B). The contact angle of brine droplets with redox mediator Citrate-Fe2+and EDTA-Fe2+was 119° ± 2° and 120° ± 2°, respectively, with the uncertainty being the standard deviation of the average contact angle for 10 replicate droplet measurements. The almost identical contact angles for brine solutions with both mediators and the FP electrode indicated the surface of FP had a similar hydrophilicity with both solutions, and this was not a factor which influenced Li+selectivity differences for these two brine solutions. It is noted that relatively large fluctuations were observed with regards to uptake and selectivity in the results in FIGs. 5A-C. We speculate some of the fluctuations may have been dependent on relatively small differences in the initial Na uptake into the FP materials at the early stages of each experiment. The larger Na+compared to Li+would be expected to impact the total accessible locations within the FP material and the structural stability of the FP, even for relatively small differences in initial Na+uptake. Future directions for investigating the instability could include investigating the influence of Na+uptake on the FP structure and total cation uptake.

[0093] In a real Li+extraction application, the FP solid material would run through multiple cycles of Li+uptake / release, and maintaining selectivity over many cycles will reduce process downtime and / or cost for FP material replacement. The same FP material was used for five extraction and release cycles (FIGs. 6A-B). The FP was lithiated via chemical redox in brine of 1 :78 Li+:Na+with EDTA-Fe2+or Citrate-Fe2+mediators. The extraction stage was always at 45 °C and allowed to proceed for 24 h. Delithiation / release was done using the chemical oxidation method described earlier. The FP material was able to uptake Li+over the 5 cycles, although there was some variation in the specific amount of Li+uptake. For chemical redox with EDTA-Fe2+, the Li+adsorption capacity for the first two cycles was ~2.5 mmol Li+g1FP. Cycles 3-5 were lower at ~2.2 mmol Li+g1FP (FIG. 6A). The reduced uptake in the later cycles may have been caused by partial insertion of Na+, which could distort the crystal structure and block Li+transport within the FP.14For the Citrate-Fe2+mediated FP reduction, the Li+adsorption was ~2 mmol Li+g1FP (FIG. 6B). For cycles 3 and 5, the adsorption capacity was slightly higher. This preliminary Li+uptake / release demonstrated cycling of the FP material to extract Li+from brine. Detailed assessment of changes in uptake and selectivity with extended cycling will be explored in future investigations of these materials.

[0094] The FP powder after 5 cycles of extraction and release were imaged by SEM, and for comparison the intial FP powder was also imaged (FIGs. 7A-F). The initial FP particles collected after chemical oxidation of LFP can be seen in FIGs. 7A and 7B. The FP particles all had primary particles sizes less than 1 mm. Larger aggregate of the smaller particles of a few mm such as seen in FIG. 7A were also observed, but these larger aggregates did not have a consistent secondary particle morphology. FP powder after 5 cycles of extraction / release by redox mediator EDTA-Fe2+can be found in FIGs. 7C and 7D. The primary particles appeared slightly smaller, which may have resulted from particle fracture. It is noted that particles likely experienced increased stress depending on the amount of sodium insertion. A previous study reported that the volume expansion of first sodiation was four times that of the first lithiation for FP materials.45As for the FP after 5 cycles of extraction / release by redox mediator Citrate-Fe2+, the morphorlogy of the FP particles was more similar to the intiial FP material, with relatively larger primary particles compared to the EDTA-Fe2+cycling experiments. The SEM images suggested the primary particles and morphology of the FP was less impacted by extraction / release of Li+Citrate-Fe2+as opposed to EDTA-Fe2+, suggesting the redox mediator with the relatively lower Na+uptake may have greater structural resilience to cycling during Li+extraction driven by redox mediators. Diffusion coefficient of redox mediators in simulated brine with Li+:Na+molar ratio of 1 :78

[0095] To provide insights into brine solution properties that may influence the Li+uptake and selectivity, electrochemical analysis was conducted on the brine to determine effective diffusion coefficients for the redox mediators in the brine solution. To estimate the effective diffusion coefficient of the redox mediators, an equimolar solution of Fe2+and Fe3+was dissolved in simulated brine solution that had the same Li+:Na+molar ratios along with the EDTA and Citrate additives at room temperature. CV was performed at increasing scan rates using a three-electrode system of platinum disk counter and working electrodes and Ag / AgCI (in saturated KCI) as reference electrode. As the scan rate increased, the potential of the cathodic / anodic peaks gradually shifted to higher / lower potentials, and peak currents also increased with increasing scan rates (FIGs. 8A, 8C, and 8E). Cathodic peak currents were plotted against the square root of scan rate, and the linear relationship of these two variables indicated the reaction was diffusion limited (FIGs. 8B, 8D, and 8F). The slope of these plots was used to calculate the effective diffusion coefficient of the redox mediators using the Randles-Sevcik equation,34

[0096] Ip= (2.687 X 105')n1'5SC(Dv')°'5(2) where Ipis the peak current (A), n is the number of electrons transferred during the redox reaction, S is the geometric surface area of the working electrode (cm2), C is iron molar concentration (mol cm3), v is the scan rate (V s1), and D is the effective diffusion coefficient of mediator in solution. The calculated diffusion coefficients of EDTA-Fe2+(PEDTA-Fe+), and Citrate-Fe2+(Dcitrate-Pez+) were 2.20 x 106cm2s1, and 1.27 x 10-7cm2s1, respectively. The diffusion coefficient for the soluble Fe (DFe2+fFe3+) in brine solution in the absence of the EDTA or Citrate ligands was 5.5 x 107cm2s1. It is noted that the soluble Fe in the brine solution would be expected to form a complex with multiple Ch species and water.46The £>Fe2+ / fe3+ of plain Fe was higher than the Dcitrate-Fe2+ of redox mediator Citrate-Fe2+, but lower than DEDTA-Fe2+ of redox mediator EDTA-Fe2+. The lower diffusion coefficient would likely result in a slower overall reaction rate when the FP pellet contacted the brine, as the porous Fe aggregate likely had solution transport restrictions through the porous pellet.28It was speculated that the slower measured diffusion coefficient for the Citrate-Fe2+may have reduced the reaction rate for reduction of the FP, which may have allowed additional time for the much lower Li+concentration to access the FP surface region in the pellet interior. In contrast, faster diffusion of the soluble redox meadiator and increased reaction would favor the reaction proceeding with less time for additional Li+to diffuse near the FP interface, and potentially drive other cations such as Na+into the FP due to the abundance of those cations in the brine.

[0097] Comparison of electrochemical and chemical redox Li+extraction processing

[0098] The results above demonstrated selective extraction of Li+from brine using redox mediators to drive FP reduction and concurrent cation insertion. An alternative method for capture of Li+from brine using intercalation materials that has been reported has involved application of electrochemical rather than chemical redox to drive reduction of the solid material.12For comparison, selectivity for Li+uptake was assessed using the same FP material used for the previous chemical redox Li+extraction experiments. A salt solution containing LiCI and NaCI with a molar ratio of 1 :78 that had the same target salts and concentration as the prior brine solution was prepared. FP and LFP were processed into composite electrodes coated on Ti wire for working and counter electrodes, respectively. The three-electrode system was completed with an Ag / AgCI (saturated KCI) reference electrode. The mass ratio of LFP:FP was kept at 1 .5:1 to ensure the working electrode was capacity limiting. The two electrodes were kept separate from each other by multiple cm, and the amount of Li+released from counter electrode LFP accounted for less than 0.5 % of the total Li+concentration in the brine solution. Hence, the influence of Li+from the counter electrode was assumed negligible. Due to the low Li+concentration, the polarization was large and the discharge plateau of FP in the brine solution shifted to a lower voltage.547A constant voltage several hundred millivolts below the discharge plateau potential for FP in brine (0.242 V vs. Ag / AgCI) was applied to the three electrodes system to lithiate the working electrode (FP). Various voltages in the range of -0.3 V to 0.0 V (in intervals of 0.1 V vs. Ag / AgCI) were evaluated to determine the impact of voltage selected on the extent of Li+extraction. The electrochemical lithiation was stopped after 50% of extent of discharge, determined based on the mass of the loaded FP with assumed theoretical capacity of 165 mAh g1. The composition of cations intercalated into FP after 50% extent of discharge can be found in FIG. 9A. As the fixed potential applied during reduction of the FP increased, the amount of Li+intercalated was reduced. The more negative applied potentials (greater driving for reduction of FP and cation insertion) resulted in increased Na+insertion into the FP. The maximum Li adsorption capacity was achieved at 0.0 V and was 3.6 mmol Li+g1FP, which was similar to the adsorption capacity achieved with both chemical redox mediators. As more Na+inserted into FP with more negative fixed potential at the working electrode, the selectivity of the intercalation material to Li+over Na+decreased from 1300 to 28 (FIG. 9B). When the applied voltage was lower than 0.0 V, using redox mediators for extraction of Li+had advantages over electrochemical extraction from the perspective of higher adsorption capacity and Li+selectivity.

[0099] Redox mediators to extract Li+from more representative Salton Sea Brine

[0100] As a final analysis, brine more representative of geothermal power discharge brine at the Salton Sea was used as the source brine for Li+capture. The Li+extraction conditions were kept the same as those used in the synthetic brine in that the redox mediator concentration was 0.1 M of EDTA-Fe2+or Citrate-Fe2+, and the reaction was allowed to proceed for 24 hours. The detailed composition of the brine can be found in Table 3.32Due to solution stability considerations, the Ca2+was omitted. The Li+extraction was carried out at three different temperatures, 23 °C, 45 °C, and 75 °C. The Li+adsorption capacity for all experiments ranged from 1.52-2.02 mmol Li+g1FP (FIG. 10A). Similar to the results discussed above with brine only containing Li+and Na+added cations, the EDTA-Fe2+mediator was less selective for Li+uptake compared to the Citrate-Fe2+, and as the temperature increased the EDTA-Fe2+was even less selective for Li+uptake. For uptake driven by EDTA-Fe2+mediator, the FP conversion attributed by the uptake of monovalent cations (Li+and Na+) increased with increasing temperature at 58%, 70%, and 83% for 23 °C, 45 °C, and 75 °C, respectively (percentages based on assumpution of complete conversion of FP requiring 6.4 mmol Li+g1FP). The increased conversion coincided with increased uptake of Na+, thus suggesting increasing temperature drove greater conversion of FP but decreased selectivity. For Citrate-Fe2+, the Li+adsorption capacity did have an increase with increasing temperature from 23 °C to 45 °C, but did not increase again for the highest temperature of 75 °C (FIG. 10A). The uptake of Li+and Na+accounted for 30% and 38% conversion of FP from 23 °C to 45 °C, respectively, and increased Li+adsorption accounted for the majority of FP conversion. In the brine solution with additional added cations, the Li+selectivity was higher for Citrate-Fe2+relative to the EDTA-Fe2+for all temperatures evaluated. The selectivity factor (ajya) was ~400 for the reaction using Citrate- Fe2+and 60 for EDTA-Fe2+(FIG. 10B).

[0101] For reference, a comprison of the outcomes of this work with other Li+extraction uptake and selectivity reported in the literature can be found in Table 4. It is noted that the prior reports were all electrochemical driven extraction. It is challenging to directly compare the reported works due to the varitiations of experimental conditions and feed brine compositions among the experiments. However, the reported Li+uptake and selectivity factor could still be considered as benchmarks for future studies. The prior literature generally had Li+uptake above 10 mg Li+g1FP, and the Li+adsorption capacity herein for reduction using both redox mediators was 14 mg Li+g1FP. Note that the Na+excess for the work herein was relatively high, being higher than all but one of the prior reported values aggregated (detailed data can be found in FIG. 17A). Lower amounts of Li+relative to Na+in the brine increased the difficulty of selective Li+separation. Some reported adsorption capacities were higher than our work. However, generally for those reports the brine solution was relatively less complex (containing only Li+and Na+) and / or the molar excess for Na+relative to Li+was much lower. Both these factors make Li+extraction easier. Meanwhile, maintaining high Li+selectivity in brine solution with great molar excess of Na+is challenging. For the reported aavalue above 350, Li+extraction with redox mediator Citrate-Fe2+shows a promising Li+selectivity factor (FIG. 17B). In addition, using redox molecules to facilite Li+extraction has potential processing advantages with regards to not requiring input electircal energy to drive the cation insertion and avoiding contact between the brine and expensive membrane materials. Furthermore, redox mediated reduction may have scale up and operational advantages compared to electrochemical extraction methods.

[0102] Conclusion

[0103] Chemical mediator-driven Li+extraction from brine into FP solid material using soluble complexes formed within the brine solution was investigated. Two redox mediators, EDTA-Fe2+and Citrate-Fe2+, were assessed, where each had a redox potential below that of the FP solid material to drive reduction of the FP and insertion of cations into the material structure. These two mediators drove chemical redox processes that selectivley extracted Li+from a variety of different brine solutions and processing conditions, including using brine with low and high Li+concentration. For brines with low concentration of competing cations (Li+:Na+, 1 :1) , the FP had high Li selectivity using both mediators, and the dominantly competing cations in brine, Na+, K+, and Mg2+, were minimally intercalated. The FP material achieved 94% of conversion with one round of extraction compared to the theoretical adsorption capacity assuming one Li+per FP. In the synthetic brine with higher concentration of competing cations (Li+:Na+, 1 :78), both redox mediators extracted 2.5 mmol Li+g1FP. The redox mediator Citrate-Fe2+had higher Li+selectivity than the redox mediator EDTA- Fe2+. This selectivity advantage still remained for brines with additional cations more representative of Salton Sea geothermal discharge brine. Applying redox mediators for Li+extraction provided new insights to Li+extraction from aqueous solutions, and a route to improve Li+selectivity. The results suggested that at least for these two redox mediators evaluated that reducing the reaction rate favored incresaing the selectivity for Li+uptake from the brines. Redox mediated extraction was demonstrated to have encouraging results as a method to selectively extract Li+from complex brine solutions. This scalable extraction setup provided a promising route towards Li extraction from complex brine solutions.

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[0159] EXAMPLE 1 SUPPLEMENTARY INFORMATION Table 1. Simulated brine composition with high Li+concentration and 1 :1 molar ratio of target Li+:M+ / 2+, cations (M = K+, Na+, Mg2+). Chemicals listed in their order of addition.

[0160] “X” and “Y” were part of the baseline solution during preparation of the brine. These amounts were determined after subtracting the mass intended for all the other chemicals, and the mass of each chemical was calculated based on a targeted volume of 0.1 L of brine. The total mass of brine was 100 g. Due to the different mass of competing cations (M), the mass of the baseline solution (X or Y) slightly varied. The mass ranges for X and Y in the prepared solutions was 78-90 g and 80-90 g, respectively.

[0161] Table 2. Simulated brine composition to match the Li+:Na+molar ratio representative of the Salton Sea (1 :78).1Chemicals listed in their order of addition.

[0162] “X” and “Y” were part of the baseline solution during preparation of the brine. These amounts were determined after subtracting the mass intended for all the other chemicals, and the mass of each chemical was calculated based on a targeted volume of 0.1 L of brine. The total mass of brine was 100 g. Due to the different mass of the redox mediator, EDTA-Fe2+and Citrate- Fe2+, the mass of the baseline solution (X or Y) slightly varied. The mass ranges for X and Y in the prepared solutions was 70-80 g and 70-84 g, respectively.

[0163] Table 3. Composition for brine prepared to be more representative of discharge from geothermal power plants at the Salton Sea.1

[0164] Table 4. Comparison of Li+adsorption capacity, selectivity, and energy consumption with reported literature using electrochemical methods for Li+extraction from brine.

[0165]

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[0179] It should be noted that ratios, concentrations, amounts, and other numerical data may be expressed herein in a range format. It is to be understood that such a range format is used for convenience and brevity, and thus, should be interpreted in a flexible manner to include not only the numerical values explicitly recited as the limits of the range, but also to include all the individual numerical values or sub-ranges encompassed within that range as if each numerical value and sub-range is explicitly recited. To illustrate, a concentration range of “about 0.1 percent to about 5 percent” should be interpreted to include not only the explicitly recited concentration of about 0.1 weight percent to about 5 weight percent but also include individual concentrations (e.g., 1 percent, 2 percent, 3 percent, and 4 percent) and the sub-ranges (e.g., 0.5 percent, 1.1 percent, 2.2 percent, 3.3 percent, and 4.4 percent) within the indicated range. The term “about” can include traditional rounding according to significant figures of the numerical value. In addition, the phrase “about ‘x’ to ‘y’” includes “about x’ to about ‘y’”.

[0180] Many variations and modifications may be made to the above-described aspects. All such modifications and variations are intended to be included herein within the scope of this disclosure and protected by the following claims.

Claims

CLAIMSWhat is claimed:

1. A method of processing a brine solution to capture lithium ions (Li+), comprising: mixing the brine solution with an Fe2+-compound and a complexing agent to form a first mixture that includes a dissolved Fe2+-complex including the Li+; exposing the first mixture to a plurality of particles, wherein the particles comprise an intercalation material, wherein the intercalation material includes a transition metal cation compound, wherein a portion of the transition metal cations in the intercalation material are reduced and Li+is intercalated into the intercalation material; and exposing the plurality of particles to an oxidizing solution, wherein the Li+in the intercalation material is released when the portion of the transition metal cations in the intercalation material are oxidized to reform the intercalation material and a Li+-solution.

2. The method of claim 1 , wherein the Fe2+compound is selected from the group consisting of: FeCI2, FeSO4, Fe(NO3)2, or a mixture thereof.

3. The method of claims 1 or 2, wherein the complexing agent is selected from the group consisting of: citrate, ethylenediaminetetraacetic acid, and a mixture thereof.

4. The method of claims 1 , 2, or 3, wherein the intercalation material contains a Fe3+transition metal and is an intercalation Fe3+-solid material.

5. The method of claim 4, wherein the intercalation Fe3+-solid material contains a compound selected from the group consisting of: iron(lll) phosphate, iron pyrophosphate and a mixture thereof.

6. The method of claim 1 , wherein the particles have a dso of about 0.5 pm to 1 mm.

7. The method of claim 1 , wherein the particles have a dso of about 0.5 pm to 30 pm.

8. The method of claim 1 , wherein the particles are agglomerated into larger aggregates.

9. The method of claims 1 , 2, 3, or 4, wherein the oxidizing solution includes a first oxidant and an acid.

10. The method of claim 9, wherein the first oxidant is a peroxide.11 . The method of claim 10, wherein the peroxide is selected from the group consisting of: hydrogen peroxide, alkali peroxide, and a mixture thereof.

12. The method of claim 9, wherein the acid is selected from the group consisting of: acetic acid, sulfuric acid, hydrochloric acid, oxalic acid, or a combination thereof13. The method of claims 1 , 2, 3, or 4, wherein the oxidizing solution includes a second oxidant.

14. The method of claim 13, wherein the second oxidant is a Fe3+-compound.

15. The method of claim 14, wherein the Fe3+-compound of the oxidizing solution is selected from the group consisting of: FeCh, [Fe(CN)6]3', Fe2(SO4)3, Fe(NO3)3, and a mixture thereof.

16. The method of claim 1 , wherein after exposing the first mixture to the plurality of particles, the plurality of particles is exposed to a rinse solution.

17. The method of claims 1 , 2, 3, or 4, wherein the brine solution is selected from the group consisting of: a natural brine, a synthetic brine, or a combination thereof.

18. The method of claims 1 , 2, 3, or 4, wherein the brine solution is selected from the group consisting of: a geothermal brine, an oil field brine, a brine from hard rock lithium mining, or a combination thereof.

19. The method of claim 1 , wherein the first solution has a pH of 2 to 8, wherein exposing the first mixture to the plurality of particles is performed a temperature of about 10° C to 120° C at a pH of 2 to 8, and wherein exposing the plurality of particles to the oxidizing solution is performed a temperature of about 10° C to 120° C at a pH of 2 to 8.

20. The method of claim 1 , wherein capture of lithium from the brine has a selectivity factor (a a) of about 10 to 10,000.

21. The method of claim 1 wherein the intercalation material has an adsorption capacity of about 0.5 to 8.0 millimoles per gram of intercalation material.

22. A system for chemically treating a brine solution to capture lithium ions (Li+), comprising: a first structure for mixing a brine solution with an Fe2+-compound and a complexing agent to form a first mixture including a Fe2+-complex; a flow system to flow the first mixture; and a second structure that receives the first mixture, wherein the second structure functions to contact a first mixture with a plurality of particles, wherein the first mixture includes a Fe2+-complex, wherein the particles comprise an intercalation material, wherein the intercalation material includes a transition metal cation compound, wherein a portion of the transition metal cations in the intercalation material are reduced and Li+is intercalated into the intercalation material to form Li+intercalated intercalation material; wherein the second structure also functions to mix the plurality of particles with an oxidizing solution after the first mixture is removed from the second structure, wherein the Li+in the Li+intercalated material is released when the portion of the transition metal cations in the intercalation material are oxidized to reform the intercalation material and a Li+-solution.

23. The system of claim 22, wherein the second structure is a fixed packed bed reactor, and agitated tank, or a fluidized bed.

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