Phytic Acid‐Based Deep Eutectic Solvents for Metal Extraction from Lithium Cobalt Oxide and Nickel Manganese Cobalt and the Use of the Resulting Leachates as Electrolytes for 2.0 V Supercapacitors

Abstract

Li‐ion batteries (LIBs) are essential in modern society but raise environmental concerns due to the intensive use of metals in cathodes and the challenges of end‐of‐life disposal. Besides traditional pyrometallurgical and hydrometallurgical processes used for metal recovery, deep eutectic solvents (DESs) have recently emerged as greener alternatives for leaching metals from spent cathodes of LIBs. A key drawback is, however, the unresolved recovery of the DES, whose cost can represent 30‐60% of the leachate, thereby reducing the overall sustainability of the process. Herein, we used the leachates as electrolytes for supercapacitors. DESs based on phytic acid provided leachates with mass loadings of 40 mg of lithium cobalt oxide (LCO) or lithium nickel manganese cobalt oxide (NMC) per gram of DES. The typically poor performance of acidic leachates as electrolytes was addressed through chemical and solvent treatments. Neutralization with tetramethylguanidine expanded the electrochemical window, while dilution with water and/or water–dimethyl sulfoxide mixtures enhanced ionic mobility and rate capability. As a result, the processed leachates delivered energy densities of ≈17.9 Wh kg⁻¹ at 488.35 W kg⁻¹ and 5.77 Wh kg⁻¹ at 4343.72 W kg⁻¹, in the range of those provided by much less cost‐efficient electrolytes such as 21 m LiTFSI.

Leachates with DES based on phytic acid and metals from spent LIBs were used as electrolytes for supercapacitors.

1. Introduction

Li‐ion batteries (LIBs) have become essential in modern society due to the rapid rise in energy storage demand from consumer electronics, such as cellphones, smartphones, and other electronic devices.^{[ 1 ]} However, the growing reliance on LIBs with ever‐increasing capacities, driven by both the electric vehicle (EV) and stationary energy storage markets, poses a significant environmental threat. This is primarily due to 1) the rising consumption of raw materials (i.e., mainly the metals used in the cathodes) and 2) the disposal of a vast number of spent batteries, the latter of which is expected to become an even more pressing issue in the near future. Recycling and reusing cathodes from spent LIBs not only aid in the safe disposal of metals but also help mitigate the shortage caused by their continuous extraction from natural ores.^{[ 1} , ^{2 ]} Furthermore, from a circular economy perspective, recycling helps close the loop by reintegrating materials into the value chain.^{[ 3 ]}

Pyrometallurgical and hydrometallurgical processes have been typically used for cathode recycling.^{[ 4} , ⁵ , ⁶ , ^{7 ]} Despite the latter offers a more sustainable alternative than the former, the neat eco‐friendly features of most hydrometallurgical routes are yet quite questionable.^{[ 1 ]} Deep eutectic solvents (DESs) have recently emerged as a sort of sustainable solvents capable of the efficient leaching of metals from spent cathodes of LIBs.^{[ 8 ]} DESs are formed by hydrogen bond (HB) complexation between HB donors (HBDs) and acceptors (HBAs) composed of organic acids (as they are also suitable extractants by their own) as HBDs and ammonium salts (e.g., choline chloride or betaine hydrochloride) as HBAs.^{[ 9} , ¹⁰ , ^{11 ]} In particular, DESs have been used for metal recovery of different spent cathodes such as lithium cobalt oxide (LiCoO₂, abbreviated as LCO), lithium nickel manganese cobalt oxide (LiNi_xCo_yMn_zO₂ with x + y + z = 1, abbreviated as NMC), lithium nickel cobalt aluminium oxide (LiNi _x Co _y Al _w O₂, with x + y + w = 1, abbreviated as NCA), and lithium manganese oxide (LiMn₂O₄, abbreviated as LMO).^{[ 12} , ¹³ , ¹⁴ , ¹⁵ , ^{16 ]} In all these cases, the increase of the leaching efficiency and/or the capability to deal with high mass loadings (i.e., originally starting with 40–50 mg of cathode per gram of DES,^{[ 17} , ^{18 ]} and more recently increasing to 70 and up to 120 mg of cathode per gram of DES^{[ 19} , ^{20 ]}), the use of mild temperatures (i.e., with first works using temperatures as high as 180 °C,^{[ 8 ]} now decreasing to milder temperatures),^{[ 21} , ²² , ^{23 ]} and/or the reuse of recovered metals to close the recycling loop^{[ 24 ]} have been the challenges most widely explored . However, the recovery of the DES remains an unresolved problem, and its reuse in subsequent dissolution processes of spent cathodes is limited to 3–5 cycles.^{[ 21} , ^{25 ]} This problem is by no means trivial as, opposite to regular hydrometallurgical processes where the metals are the valuable part of the leachate, DES costs may account for over 30% of the leachate cost, reaching up to 60% in some cases.

Besides the intrinsic difficulties to recover the DES (i.e., where solvents and/or reagents added for metal separation and recovery are difficult to remove from the DES and are ultimately contaminating it) and the timely and costly procedures needed for the recovery of the individual metals, the rapid development of next‐generation cathode materials offering improved safety and energy density poses a genuine threat to the continued relevance of LCO and NMC cathodes.^{[ 2 ]} In this scenario, with LCO and NMC cathodes facing a tangible risk of obsolescence, the exploration of novel pathways for the reutilization of these critical metals represents a compelling area of research.^{[ 26} , ²⁷ , ^{28 ]} In this regard, our group has recently reported on the direct use of the leachate as the electrolyte of supercapacitor (SC) cells.^{[ 19} , ^{29 ]} Following this strategy, both the DES and the metals were reutilized without timely and/or costly procedures. Interestingly, these leachate‐based electrolytes exhibited a synergistic effect where the presence of metals, even at low concentrations, contributed to extending the electrochemical stability window (ESW) of the original DES, thereby improving the stored energy density.

Despite this approach allowed closing the recycling loop for the DESs and the metals, improving the sustainable features of the DES used for metal recovery would yet be of interest. It is worth noting that the use of DESs with one acidic component has typically provided better capabilities for high mass loadings. In this regard, the use of acids of natural origin has been explored in just few examples of the so‐called natural DESs (NADESs, containing compounds such as citric acid^{[ 30 ]} or phytic acid, PA).^{[ 31 ]} Moreover, the study of this sort of leachates as electrolytes in SCs remains fully unexplored.

Herein, we aimed to obtain leachate‐based electrolytes using ethaline, a DES composed of 2 equivalents of ethylene glycol (EG) as HBD and 1 equivalent of choline chloride (ChCl) as HBA, acidified with PA. The cathodes of choice were LCO and NMC. Ethaline (e.g., ChCl:2EG) was chosen because it has been widely used as electrolyte in SC cells, exhibiting good performances both in neat form^{[ 32 ]} and forming ternary mixtures with either urea^{[ 33 ]} or H₂O.^{[ 34 ]} PA (50 wt% in H₂O) was added to improve the Li and Co solvent capabilities of ethaline which, otherwise, needs of high temperatures and long extraction times to perform.^{[ 8 ]} We continued with the evaluation of the ESW of the resulting leachates depending on their acidic nature. It is worth noting that, in DES‐based electrolytes, the presence of components with an acidic nature causes a detrimental effect on the electrochemical performance because of the narrowing of the ESW. Thus, leachates with different acidic natures were obtained upon the addition of tetramethylguanidine (TMG). The evaluation of these leachates as electrolytes was achieved by cyclic voltammetry (CV) in a two‐ and three‐electrode configurations, galvanostatic charge/discharge (GCD), chronoamperometry (CA) in a three‐electrode configuration, and electrochemical impedance spectroscopy (EIS). Experiments in the three‐electrode configuration were particularly suitable to assess the ESWs of the different electrolytes. Cycling experiments were also performed to confirm these ESWs.

2. Results and Discussion

2.1. DES‐Based Leachates: Component Selection and Characterization

As mentioned in the introduction, ethaline‐assisted solid–liquid extractions typically required temperatures as high as 180 °C to work at pseudo‐high loadings.^{[ 8 ]} Actually, the use of high temperatures was a common issue for other EG‐based DESs.^{[ 14} , ^{15 ]} However, the use of high temperatures is not desirable given the thermal decomposition reported for ChCl:2EG at 180 °C^{[ 35 ]} can lead to byproducts that, in our case, could ultimately compromise the electrolyte's electrochemical performance.

Our group recently demonstrated how loadings as high as 70 mg of NMC per gram of DES could be dissolved by ethaline‐based ternary DESs (i.e., formed by addition of 1 equivalent of p‐toluene sulfonic acid in its monohydrated form, pTsOH:1H₂O) at temperatures as low as 60 °C.^{[ 19 ]} In this work, we replaced pTsOH:1H₂O by addition of different amounts of an aqueous solution of PA (50 wt% in H₂O, PA:36.6H₂O in molar ratios). In particular, we used (ChCl:2EG):n(PA:36.6H₂O) mixtures with different molar ratios (e.g., n = 0.1, 0.2, 0.4, 0.8, and 1.6). We first studied the solvent capabilities of these mixtures for LCO. At 85 °C, every mixture was able to dissolve 20 mg of LCO per gram of mixture in 24 h, but only the mixture with n = 0.2 (e.g., ChCl:2EG:0.2PA:7.3H₂O) was able to dissolve 40 mg of LCO per gram of mixture in 24 h (Table S1,Supporting Information).

This solvent capability was confirmed by UV‐Vis spectroscopy. For this purpose, we prepared five different leachates with a mass loading of 1 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O for which the full dissolution of LCO was ensured. Then, we diluted these leachates (with further ChCl:2EG) to contents of 0.011, 0.033, 0.065, 0.086, and 0.1 mg of LCO per gram of liquid mixture to place the intensity of the absorption bands within the 0–1 range where the relationship between absorption intensity and concentration is linear (Figure  1  A). The UV‐Vis spectrum contained bands at 616, 631, 668, and 696 nm that have been typically ascribed to Co complexes (Figure 1 A). In particular, the band at 616 nm is ascribed to octahedral cobalt complexes such as the octahedral [Co(L)₆]²⁺ complex (with L = EG and/or H₂O),^{[ 26 ]} the band at 631 nm to the hexacoordinated [CoCl₂(L)₂]^{[ 36 ]} and/or the tetrahedral [CoCl₃(L)]⁻ complex (with L = EG),^{[ 37 ]} and the bands at 668 and 696 nm to tetrahedral [CoCl₄]²⁻complexes.^{[ 13} , ^{38 ]} We represented the intensity of the band at 696 nm obtained for the different contents mentioned above (e.g., 0.011, 0.033, 0.065, 0.086, and 0.1 mg of LCO per gram of liquid mixture), and we obtained an excellent linear fit (Figure 1B). Finally, we prepared three different leachates with a mass loading of 40 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O. The leachates were first centrifuged; the supernatants were then passed through a 0.45 μm filter (to eliminate any eventual solid residue) and finally diluted with further ChCl:2EG to, theoretically, obtain contents of 0.13, 0.093, and 0.061 mg of LCO per gram of liquid mixture. The intensity of the absorption band measured for these leachates fitted well into the calibration curve depicted in Figure 1B providing contents of 0.14, 0.11, and 0.067 mg of LCO per gram of liquid mixture. It was worth noting the agreement between theoretical and experimental results thus confirming the capability of 1 gram of the ChCl:2EG:0.2PA:7.3H₂O mixture to dissolve 40 mg of LCO (Figure 1C and 1D).

Figure 1.

A) Pictures and UV‐Vis absorption bands of the five different leachates prepared by dilution of 1 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O mixture and subsequent dilution to 0.011, 0.033, 0.065, 0.086, and 0.1 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O. B) Calibration curve obtained by linear fitting between absorption intensity and mg of LCO per gram of liquid mixture. C) Pictures and UV‐Vis absorption bands of the leachates obtained by dilution of 40 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O mixture. These leachates were prepared by triplicate and, after filtration, each one was diluted to obtain 0.13, 0.093 and 0.061 mg per gram of liquid mixture. D) Deviation between theoretical values and real ones obtained from the calibration curve depicted in (B).

We also confirmed the solvent capability of the ChCl:2EG:0.2PA:7.3H₂O mixture by total reflection X‐ray fluorescence (TXRF) and inductively coupled plasma optical emission spectrometry (ICP‐OES) (Figure S1, Supporting Information). In this case, we also prepared a leachate of 40 mg of LCO per gram of ChCl:2EG:0.2PA:7.3H₂O as described above; this is by dissolution and centrifugation. Finally, a small volume of the supernatant was diluted in 1 M HNO₃ and sent to the respective analyses. The theoretical masses of Co and Li per gram of ChCl:2EG:0.2PA:7.3H₂O were 23.2 and 2.7 mg, respectively. The TXRF analysis provided a mass of Co of 24.0 mg while the ICP analysis provided masses of Co and Li of 23.7 and 2.9 mg, respectively.

In addition, we also studied the solvent capability of ChCl:2EG:0.2PA:7.3H₂O for NMC. In particular, we used two different NMCs; the first one (NMC‐1) was acquired from a chemical supplier, and the second was derived from the cathode of a recycled battery. The specific compositions of NMC‐1 and NMC‐2 (x, y, and z in LiNi_xCo_yMn_zO₂ with x + y + z = 1) were determined by ICP‐MS as described elsewhere.^{[ 19 ]} In particular, we found that x was 0.32, y was 0.44, and z was 0.24 in NMC‐1, while x = 0.55 y was 0.26, and z = 0.19 in NMC‐2. The leachates were obtained by addition of 40 mg of both NMC‐1 and NMC‐2 per gram of ChCl:2EG:0.2PA:7.3H₂O. The leachates were treated as described above: first by centrifugation, then by passing the supernatants through a 0.45 μm filter (to eliminate any eventual solid residue), and finally by diluting them with further ChCl:2EG to, theoretically, obtain contents of 0.185 and 0.194 mg of, respectively, NMC‐1 and NMC‐2 per gram of liquid mixture. The UV‐Vis spectra of these leachates is depicted in Figure S2A. Considering the Co content in the respective NMC studied herein, the good agreement found between the absorption intensity experimentally observed at 696 nm and the theoretical one calculated from the linear fit obtained from LCO‐based leachates for [CoCl₄]²⁻ complexes (Figure S2B, Supporting Information, adapted from Figure 1B) demonstrated that the solvent capability of ChCl:2EG:0.2PA:7.3H₂O for NMC was in range to that described above for LCO.

At this stage, we used DSC and ¹ H NMR spectroscopy to determine some specific features of the ChCl:2EG:0.2PA:7.3H₂O mixture. The DSC of ChCl:2EG exhibited the melting of the eutectic composition (T _m) at ≈−28.7 °C followed by a crystallization peak at ≈−43.8 °C (Figure  2  A). We also observed a broad transition at ≈10 °C (Figure 2 A) that, in recent works assigning the eutectic composition to the 16:84 molar fraction (rather than to the typical 33:66 one used to date), has been ascribed to an excess of ChCl .^{[ 39} , ^{40 ]} The lack of any T _m and the decrease of the temperature of glass transition (T _g) in the DSC scan of the mixture resulting by addition of water (e.g., ChCl:2EG:7.3H₂O in Figure 2 A) was indicating the effective participation of water in the HB structure of ChCl:2EG and the formation of new eutectic mixtures containing water within its composition.^{[ 41 ]} This trend was repeated upon the addition of PA as observed for other molecules with the capability to act as HBDs and/or HBAs in HB networks (Figure 2 A).^{[ 42 ] 1} H NMR spectroscopy provided useful insights in this regard. Thus, we first obtained the spectra of ChCl:2EG and ChCl:2EG:7.3H₂O to assess the role played by this specific amount of H₂O in the HB structure formed in the original DES. In this regard, the peaks of protons participating in HBs typically experience some sort of broadening and/or downfield shifting.^{[ 43 ]} This was indeed observed in our case for the protons of OH groups of ChCl and EG (Figure  3  A and  3B). Moreover, in DESs, the addition of H₂O drives the original DES to either the water‐in‐DES regime or the DES‐in‐water one depending, respectively, on whether H₂O participates in the HB structure of the original DES as an additional HBD and/or HBA or not. In this latter case, it is widely assumed that the system becomes a regular aqueous solution in which the original HB structure is broken and the DES components are individually solvated by H₂O molecules. The water content at which the transition from the water‐in‐DES regime to the DES‐in‐water one occurs depends on the nature of the DES components. Interestingly, previous works dealing with ¹ H NMR spectroscopy have shown the coalescence of the signals of exchangeable protons occurs at this water content.^{[ 44 ]} Based on this, the ChCl:2EG:7.3H₂O mixture was in the water‐in‐DES regime as indicated its ¹ H NMR spectrum by the presence of separated signals (without coalescence) for the OH protons of ChCl and EG and the protons of H₂O (Figure 3B). Finally, the significant broadening of every signal in the ¹ H NMR spectrum of the ChCl:2EG:0.2PA:7.3H₂O mixture revealed the active participation of PA in the HB structure formed by ChCl, EG, and H₂O (Figure 3C).

Figure 2.

DSC scans of A) ChCl:2EG (pink line), ChCl:2EG:7.3H₂O (brown line), ChCl:2EG:0.2PA:7.3H₂O (violet line), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O (green line), and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO (dark grey line) and B) ChCl:2EG:0.2PA:7.3H₂O:LCO (light grey line), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO (blue line), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO (cyan line), ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO (orange line), and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO (dark yellow line).

Figure 3.

¹H NMR spectra of A) ChCl:2EG, B) ChCl:2EG:7.3H₂O, C) ChCl:2EG:0.2PA:7.3H₂O, and D) ChCl:2EG:0.2PA:1.2TMG:7.3H₂O.

We also studied the mixtures by ATR‐FTIR spectroscopy (Figure  4 ). The ATR‐FTIR spectrum of ChCl:2EG has been described in previous works.^{[ 45 ]} Thus, the bands at 1083 cm⁻¹ and at 1032 cm⁻¹ were assigned to, respectively, the symmetrical and asymmetrical stretching of the C—O bond (e.g., ν _s(C—O) and ν _as(C—O) modes) of EG (Figure 4 A). The appearance of both bands revealed the monodentate configuration of EG in every mixture,^{[ 46 ]} even in the presence of the Li and Co extracted from LCO. The bands at 880 and 860 cm⁻¹ were also ascribed to EG, in particular to the bending and stretching vibrations of the C—C bond (e.g., δ(C—C) and ν(C—C) modes). The stretching vibrations of the C—N bond in the (CH₃)₃ N⁺ group of ChCl was observed at 1477 cm⁻¹.^{[ 47 ]} The bands at 955 and 923 cm⁻¹ (this latter of small intensity and appearing as a shoulder) were attributed the ν _s(C—N) and the ν _as(C—N) mode.^{[ 48 ]} The presence of PA in ChCl:2EG:0.2PA:7.3H₂O was reflected in the appearance of two main bands at 1185 and 1074 cm⁻¹, assigned to ν _as(P–O in HPO₄ ⁻).^{[ 49} , ⁵⁰ , ^{51 ]} It is worth noting that, at acidic pHs, PA should be in the form of C₆H₆(HPO₄ ⁻)₆ ion. In our case, the band at 1185 cm⁻¹ in ChCl:2EG:0.2PA:7.3H₂O was red shifted to lower wavenumbers in ChCl:2EG:0.2PA:7.3H₂O:LCO, thus indicating the capability of C₆H₆(HPO₄ ⁻)₆ ions to engage with metals (e.g., Co and/or Li) by electron donation (i.e., as a Lewis base).^{[ 52 ]}

Figure 4.

FTIR spectra of EG (black line in A), PA:36.6H₂O (red line in A), ChCl:2EG (pink line in A), ChCl:2EG:0.2PA:7.3H2O (purple line in A), ChCl:2EG:0.2PA:7.3H2O:LCO (grey line in A), ChCl:2EG:0.2PA:1.2TMG:7.3H2O:LCO (blue lines in A and B), ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO (orange line in B), ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO (dark yellow line in B), and ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO (cyan line in B).

2.2. LCO‐Based Leachates: Electrochemical Study

Preliminary electrochemical studies were conducted using ChCl:2EG:0.2PA:7.3H₂O and ChCl:2EG:0.2PA:7.3H₂O:LCO as electrolytes to assess whether the presence of metals provided some beneficial effects in the energy storage capabilities. The electrochemical response of these electrolytes was performed with Swagelok cells, using two mesoporous carbons as electrodes and a 0.45 μm PVDF filter paper as separator (see the Experimental section for further details). The mesoporous carbon of choice has been described in detail in our previous works^{[ 53} , ⁵⁴ , ^{55 ]} and was selected based on its textural properties, with a high surface area (≈2681 m² g⁻¹) and a large contribution of mesopores to the overall porosity with V _mesopores above 70% of the total pore volume and an external surface area of ≈500–650 m² g⁻¹ (Figure S3, Table S2, Supporting Information).

Previous works using DESs as electrolytes for SCs have revealed that the acidic features of the HBD plays a role in their respective ESWs. Thus, ESWs of 1.4, 2.0 and 2.2 V have been observed in DESs with, respectively, pTsOH,^{[ 19 ]} EG,^{[ 29} , ³² , ^{34 ]} and urea^{[ 56 ]} as HBDs. In this work, we used CVs (Figure S4, Supporting Information) and CAs (Figure  5 ) in a three‐electrode configuration to determine the ESWs. Previous works using CAs in a three‐electrode configuration categorized the so‐called “safe potentials” for current densities below 0.1 A g⁻¹ and/or above −10 μA cm⁻².^{[ 57} , ^{58 ]} In our case, we investigated the samples for which the CA curves remained at current densities below 0.1 A g⁻¹ for both negative and positive potentials. Figure 5 shows the CA curves obtained at potentials ranging from −0.1 to −1.4 V and from 0.1 to 1.2 V. The application of the threshold of 0.1 A g⁻¹ for the “safe” regime resulted in potential limits of [−0.4, 1.0] for ChCl:2EG:0.2PA:7.3H₂O (Figure 5 A and 5B); this is an ESW of ≈1.4 V in agreement with that observed for HBDs with strong acidic features.^{[ 19 ]} As in previous works, the presence of LCO in ChCl:2EG:0.2PA:7.3H₂O:LCO resulted in a widening of the ESW to ≈1.7 V (e.g., [−0.7, 1.0] in Figure 5C and 5D, and Figure  6  A). The comparison between the ESWs obtained for ChCl:2EG:0.2PA:7.3H₂O and ChCl:2EG:0.2PA:7.3H₂O:LCO allowed identifying how the presence of metals prevented the electrolyte decomposition at the negative electrode (i.e., where reduction occurs) more so than at the positive one (i.e., where oxidation occurs). The widening of the cathodic limit has typically been ascribed to the occurrence of certain passivation of the negative electrode that prevents the hydrogen evolution reaction (HER). It seems plausible that the presence of metals favored this sort of processes.

Figure 5.

CA curves obtained in a three‐electrode configuration for ChCl:2EG:0.2PA:7.3H₂O A,B), ChCl:2EG:0.2PA:7.3H₂O:LCO C,D), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO E,F), ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO G,H), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO I,J), and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO K,L). Different potentials were studied for the negative (A, C, E, G, I, and K) and the positive (B, D, F, H, J, and L) electrodes depending on the mixture.

Figure 6.

A) CV curves, B) GCD curves, C) the Nyquist plot, and D) the Ragone plot obtained with ChCl:2EG:0.2PA:7.3H₂O (purple lines), ChCl:2EG:0.2PA:7.3H₂O:LCO (grey lines), and ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO (blue lines) acting as electrolytes in SCs operating at 1.4, 1.7, and 2.0 V, respectively. For comparison, the Ragone plot includes data for 21 m LiTFSI (red line) and TCENMC/G70 (black line; see ref. [19] for nomenclature) obtained in SC cells operating at 2.0 and 1.7 V, respectively.

Unfortunately, this widening of the ESW was yet insufficient for ChCl:2EG:0.2PA:7.3H₂O:LCO to be consider a good electrolyte as revealed by the GCD curves and the Ragone plot depicted in Figure 6B and 6D. Moreover, the increase of the viscosity in ChCl:2EG:0.2PA:7.3H₂O:LCO because of the presence of LCO was detrimental for the rate capability of the SC as reflected in the internal resistances obtained from the IR drop at the initiation of the current discharge for different current densities^{[ 59 ]} and from the intersection with the X‐axis of the vertical line of the Nyquist plot (R _{int_IRdrop} and R _{int_ESR}, respectively, in Table  1 ).^{[ 60 ]}

Table 1.

Densities (ρ, in g mL⁻¹), viscosities (η, in cP) and ionic conductivities (σ, in mS cm⁻¹) measured for the different mixtures. The table also includes the internal resistances in SC cells measured from the IR drop (R _{int_IRdrop}, in Ω) and the Nyquist plot (R _{int_ESR}, in Ω) using the mixtures as electrolytes.

Mixtures	ρ [g mL−1]	η [cP]	σ [mS cm−1]	R int_IRdrop [Ω]	R int_ESR [Ω]
ChCl:2EG	1.11484	47.9	6.51	12.1	10.7
ChCl:2EG:0.2PA:7.3H2O	1.22441	26.3	22.2	4.2	3.5
ChCl:2EG:0.2PA:7.3H2O:LCO	1.27710	75.7	6.97	18	16.5
ChCl:2EG:0.2PA:1.2TMG:7.3H2O:LCO	1.23749	255.1	3.15	30.2	26.9
ChCl:2EG:0.2PA:1.2TMG:14.6H2O:LCO	1.19065	38.9	7.22	8.6	7.9
ChCl:2EG:0.2PA:1.2TMG:14.6H2O:7.3DMSO:LCO	1.16847	22.5	5.80	18.1	15.2
ChCl:2EG:0.2PA:1.2TMG:7.3H2O:7.3DMSO:LCO	1.17592	30.4	3.41	21.1	18.7

Based on this, our aim was to increase the pH of the leachate by addition of a suitable base. The base of choice was TMG because its high solubility in water and its strong basicity allowed rising the pH while maintaining the homogeneity of the mixture. Thus, the pH evolved from ≈1 for ChCl:2EG:0.2PA:7.3H₂O:LCO to ≈10 after the addition of 30 wt% TMG (Figure S5A, Supporting Information). The particular coordination features of ChCl:2EG:0.2PA:1.2TMG:7.3H₂O (20 wt% of TMG content) were investigated by DSC (Figure 2B) and FTIR spectroscopy (Figure 4 A and 4B). The DSC scans of both ChCl:2EG:0.2PA:1.2TMG:7.3H₂O and ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO exhibited none T _m thus indicating the favored integration of, first, TMG within the HB structure of ChCl:2EG:0.2PA:7.3H₂O and, then, LCO within the HB structure of ChCl:2EG:0.2PA:1.2TMG:7.3H₂O. With regard to FTIR spectroscopy and as compared to ChCl:2EG:0.2PA:7.3H₂O or ChCl:2EG:0.2PA:7.3H₂O:LCO, the presence of TMG was reflected in the slight increase of the intensity of the band at 930 cm⁻¹ assigned to the bending vibration of the N–H group (e.g., δ(N–H) mode).^{[ 61 ]} Moreover, the band at 1673 cm⁻¹ was assigned to the wagging vibration of the N–H₂ ⁺ group (e.g., γ(N–H₂ ⁺) mode) and hence indicative of the presence of [TMG]⁺ (i.e., the protonated TMG formed by proton transfer between PA and TMG) that caused the pH rise in the medium.^{[ 62} , ⁶³ , ^{64 ]}

We also used ¹ H NMR spectroscopy to study the mixtures without metals, as the presence of paramagnetic metals poses serious limitations to the achievement of useful information from the mixtures with metals. Interestingly, the value of the integral of the peak at ≈7.69 ppm in the ¹ H NMR spectrum of the ChCl:2EG:0.2PA:1.2TMG:7.3H₂O mixture confirmed the complete proton transfer from PA to TMG and the formation of the [TMG₆]⁶⁺[PA]⁶⁻ salt (Figure 3D). It is worth noting that, as described elsewhere,^{[ 65 ]} the formation of alike ionic liquids (ILs) (e.g., [TMG]⁺[LA]⁻, where LA was lactic acid) is favored when the ΔpK between the acid and the base is above 9.^{[ 66} , ^{67 ]} Moreover, the disappearance of the band at 1185 cm⁻¹ in the FTIR spectrum of the ChCl:2EG:0.2PA:1.2TMG:7.3H₂O mixture (blue line in Figure 4 A) was also in agreement with the occurrence of full proton transfer between PA and TMG because it revealed the lack of remaining [PA]⁶⁻ ions available to engage with metals. This lack of [PA]⁶⁻ ions engaged with metals was actually anticipated by the UV‐Vis spectra depicted in Figure 1 A in which the main complexes were in the form of [Co(L)₆]²⁺ with L = EG and/or H₂O, [CoCl₂(L)₂] and [CoCl₃(L)]⁻ with L = EG, and [CoCl₄]²⁻.

The CV curves obtained in a two‐electrode configuration showed a widening of the ESW to ≈2.0 V for LCO‐based mixtures with 15, 20, and 30 wt% TMG (Figure S5B, Supporting Information). The capacitance retention after 10 000 cycles increased along with the increase of the pH, from ≈77% for the mixture with 15 wt% TMG to 84% for the mixture with 20 wt% TMG up to 86% for the mixture with 30 wt% TMG (Figure S4D, Supporting Information). However, it is wort noting that the rise in TMG concentration was accompanied by a reduction in specific capacitance (Figure S6, Supporting Information). In this regard, R _{int_ESR} also experienced a dramatic increase along with the TMG content that revealed the detrimental effect caused by TMG in the rate capability of the SC (Figure S4C, Supporting Information, Table 1). Actually, this detrimental effect was also evident in the CVs, where the distortion of the ideal rectangular shape was more pronounced for metal‐free‐TMG‐based mixtures than for their respective counterparts with metals (Figure S7, Supporting Information). Interestingly, cyclability also suffered from the sluggish ion transport, increased internal resistance, and/or irreversible and inefficient charge–discharge processes that caused the distortion of the ideal rectangular shape of the CV curves (Figure S8, Supporting Information).

Based on this, we decided to continue the work with the mixture with 20 wt% TMG (e.g., ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO) as, even though none of its individual properties was the best, they were the best overall. In particular, the use of CVs (Figure S4, Supporting Information) and CAs in a three‐electrode configuration obtained for ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO resulted in potential limits of [−1.2, 0.8] (Figure 5 E and 5 F); this is a widening of the ESW to ≈2.0 V (Figure 6 A). Indeed, the GCD curves and the Ragone plot depicted in Figure 6B and 6D revealed the significant improvement experienced by the ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO mixture as compared to ChCl:2EG:0.2PA:7.3H₂O:LCO, albeit not yet sufficient (mainly in terms of rate capability) as compared to 21 m LiTFSI. Moreover, the capacitance retention after 10 000 cycles at 2.0 V was also remarkable as compared to ChCl:2EG:0.2PA:7.3H₂O:LCO (Figure  7 A and Figure S9, Supporting Information).

Figure 7.

Evolution of the capacitance retention (in%) and the coulombic efficiency during 10 000 cycles carried out 4 A g⁻¹ for A) ChCl:2EG:0.2PA:7.3H₂O:LCO (grey line and symbols) and ChCl:2EG:PA:1.2TMG:7.3H₂O:LCO (blue line and symbols) and B) ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO (orange line and symbols), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO (cyan line and symbols) and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO (dark yellow line and symbols), acting as electrolytes in SCs operating at 2.0 V.

Next efforts when using the ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:LCO mixture as the electrolyte were focused on improving the rate capabilities of the SC (Figure 6D and Figure S10, Supporting Information, Table 1) while maintaining the energy density and the cyclability. For this purpose, we explored the effect that dilution with H₂O, DMSO and a mixture of both had on the electrochemical performance. In a first approach, we prepared three new samples in which the solvent equivalents were doubled (e.g., ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO, ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO) and tripled (e.g., ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO). The FTIR spectra of these mixtures revealed that neither the addition of further H₂O nor of DMSO modified the above described interactions assigned to the ethaline and the [TMG₆]^{6 +}[PA]⁶⁻ salt (Figure 4B). Interestingly, the addition of DMSO caused the appearance of a strong band at ≈1020 cm⁻¹. This band was assigned to the stretching vibration of the S=O bond (e.g., ν(S=O) mode) that typically appears at ≈1050 cm⁻¹ for pure DMSO (free S=O) and is red shifted to ≈1020 cm⁻¹ in certain H₂O:DMSO mixtures because of the formation of HBs.^{[ 68 ]} The DSC scans of the mixtures with DMSO revealed a better DMSO integration for mixtures with 2H₂O:DMSO in which none T _m was observed (Figure 2B). This was quite interesting as the use of 2H₂O:DMSO mixtures has already proved effective to enlarge the operational potential window of aqueous electrolytes.^{[ 53 ]}

In this case, solvent addition resulted in a significant decrease of both the R_{int_ESR} and the R_{int_IRdrop} along with the H₂O content: ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO < ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO < ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO) (Figure S10, Supporting Information, and Figure  8A and 8B, Table 1). Meanwhile, the potential limits obtained from the CV curves in both a two‐ and a three‐electrode configuration (Figure S3, Supporting Information , Figure 8C) and the CA in a three‐electrode configuration followed the opposite trend, with ESWs that ranged from 1.8 V for ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO (Figure 5G and 5H) to 2.0 V for ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO (Figure 5I and 5J) and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO (Figure 5K and 5L). These potential limits were actually reflected in the Ragone plot, with the ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO mixture reaching energy and power densities in range to those exhibited by typical electrolytes such as 21 m LiTFSI^{[ 53 ]} (Figure 8D) and with enhanced cost efficiency (Figure  9 ).^{[ 69} , ⁷⁰ , ^{71 ]} Interestingly, the capacitance retention after 10 000 cycles at 2.0 V was also above 80% for the mixtures containing DMSO (e.g., ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO and ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO; see Figure 7B).

Figure 8.

A) The Nyquist plot, B) the GCD curves, C) the CV curves, and D) the Ragone plot obtained with ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:LCO (orange line and symbols), ChCl:2EG:0.2PA:1.2TMG:7.3H₂O:7.3DMSO:LCO (dark yellow line and symbols), and ChCl:2EG:0.2PA:1.2TMG:7.3DMSO:LCO (cyan line and symbols) acting as electrolytes in SCs operating at 2.0 V. For comparison, the Ragone plot also includes data for 21 m LiTFSI (red line and symbols) obtained in SC cells operating with our electrodes at 2.0 V.

Figure 9.

Total cost (in euros) of the electrolytes used in this work a) in comparison with some other electrolytes described in previous papers; b) 21 M LiTFSI in H₂O, c) 1.6 M TEABF₄ in CH₃CN, d) 1.6 M TEABF₄ in propylene carbonate, and e) 1 M LiPF₆ in a mixture of ethylene carbonate and dimethyl carbonate (see refs. [53,69,70,71] in main text). Data was taken from https://www.sigmaaldrich.com/ES/es on September 2023.

Finally, we focused on the study of the mixture that best performed with LCO and compared it with the analogues mixtures obtained with NMC‐1 and NMC‐2; these were ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO, ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:NMC‐1, and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:NMC‐2. Figure  10 shows the CV curves in the two‐electrode configuration, the GCD curves, the Nyquist plot, and the Ragone plot, as well as the evolution of the capacitance retention (see also Figure S11, Supporting Information) and the coulombic efficiency during 10 000 cycles carried out at 4 A g⁻¹. We found quite similar performances for LCO, NMC‐1, and NMC‐2, in agreement with our previous work also studying leachates obtained from both LCO and NMC.^{[ 29 ]}

Figure 10.

A) The CV curves, B) the GCD curves, C) the Nyquist plot, D) the Ragone plot, and E) the evolution of the capacitance retention (in%) and the coulombic efficiency during 10 000 cycles carried out 4 A g⁻¹ obtained with ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:LCO (orange line and symbols), ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:NMC‐1 (black line and symbols), and ChCl:2EG:0.2PA:1.2TMG:14.6H₂O:7.3DMSO:NMC‐2 (red line and symbols) acting as electrolytes in SCs operating at 2.0 V.

3. Conclusions

This work provided a novel approach to the treatment of leachates coming from spent LCO cathodes of LIBs. In particular, we used these leachates as electrolytes in SCs, thus helping to overcome key barriers to real‐world adoption of DES‐based metal recovery technologies in the field of LIBs.^{[ 72 ]} Moreover, we avoided the separation processes that typically follow metal extraction and are not only tedious and economically inefficient but also jeopardize the recovery of the DES (and, hence, its reuse in subsequent extractions). The novelty of this work was the use of TMG for neutralization of the acid leachate resulting from PA‐based DES. It is worth noting that acidic DESs typically provide narrow ESW, and neutralization had never been explored as an approach for the ESW enlargement. It is worth noting that the addition of TMG caused a detrimental effect in terms of rate capability. Interestingly, this problem was circumvented in the presence of metals and upon the addition of a certain cosolvent, following the common strategy to improve the rate capability of many electrochemical energy storage systems operating with water‐in‐salt electrolytes.^{[ 73} , ^{74 ]} Thus, after neutralization and dilution with H₂O and/or DMSO, the resulting electrolytes were capable to operate at ESW of 2.0 V providing energy densities of up to 17.9 W h kg⁻¹. Interestingly, the electrochemical stability of the electrolyte operating at these potentials was reflected in the good capacitance retentions (i.e., of ≈85%) obtained after 10 000 charge‐discharge cycles at 2.0 V. Moreover, the DES composition used in this work was particularly well suited for Co recovery from LCO and NMC cathodes, with efficiencies of ≈100% for extractions carried out at mild experimental conditions (e.g., 60 °C) and loadings as high as 40 mg_LCO/NMC/g_DES. Based on this, this work contributes to the circular economy in the energy storage field by adding a new path for the direct utilization of the leachate.

4. Experimental Section

4.1.

4.1.1.

Materials

Ethylene glycol (EG, >99.5%) was obtained from Supelco. Phytic acid solution (PA, 50 wt% in water), dimethyl sulfoxide (DMSO), tetramethylguanidine (TMG), lithium cobalt oxide (LCO), and PTFE emulsion (65 wt%) were purchased from Sigma‐Aldrich. Isopropanol was ordered from Scharlau. Choline chloride (ChCl, >98%) was obtained from VWR Chemicals and subjected to a drying process at 90 °C for 24 h prior to the preparation of DESs. The 0.45 μm PVDF membrane was purchased from Durapore. Black carbon was obtained from CABOT Corporation. Lithium nickel manganese cobalt oxide (NMC) was either purchased from TCI or obtained by precise dismantling of a 3.7 V LIB model 18 650 obtained from Pilanet Inc.

DES Preparation

The synthesis was accomplished by physical mixing of ChCl and EG in a 1:2 molar ratio. The mixture was heated to 60 °C for 2 h to obtain a transparent liquid without further treatment.

Preparation of the Electrolyte Without Metals

The liquid binary mixtures used as electrolyte were prepared by simple addition of the DES to the aqueous solution of PA in different molar ratios (Table S1, Supporting Information).

Preparation of the Electrolyte with Metals (Leachates)

Extraction of metal ions from LCO and NMC was conducted with the chosen mixture (≈5 g) under magnetic stirring (500 rpm) in a 25 mL closed glass vial. The LCO/NMC oxide was added to the DES‐based mixture at room temperature under agitation and then heated to 85 °C over 16–24 h. The resulting leachate was centrifuged, and the liquid phase was filtrated (with regenerated cellulose syringe filters with a pore size of 0.45 μm and a diameter of 25 mm, from Filter‐Lab) to remove any residual non‐dissolved oxide prior to UV‐Vis analysis and/or use as an electrolyte.

Preparation of Carbons

Carbon materials were obtained as described elsewhere.^{[ 53} , ⁵⁴ , ^{55 ]} Briefly, 4 g of sodium citrate dihydrate and 4 g of urea were dissolved in 30 mL of deionized water. After stirring for 2 h, the homogeneous aqueous solution was frozen at −80 °C and freeze‐dried using a Thermo Savant (Micromodulyo–230). The resultant white powder was thermal treated at 800 °C for 2 h under a nitrogen (N₂) atmosphere, with a heating rate of 3 °C min⁻¹. The carbon obtained after this treatment was washed multiple times, first with a diluted solution of hydrochloric acid (1 M) and then with deionized water to eliminate any remaining inorganic salts. Finally, the carbon was dried at 80 °C overnight in a vacuum oven.

Preparation of Electrodes

Carbon powder (40 mg) was blended in a glass vial with 5 mg of carbon black and 20 mL isopropanol. Then, a certain amount of PTFE emulsion (≈6.2 mg) was also added to this mixture. The suspension was later subjected to 15 min of ultrasonication (BANDELIN SONOREX) to ensure homogeneity. Subsequently, part of the isopropanol was evaporated at 90 °C to obtain a slurry. This slurry was drop‐cast onto carbon fiber paper and dried overnight at 80 °C in a vacuum oven. The resulting electrodes were cut into circular disks (≈0.8 cm diameter) with a total mass of ≈4.4 mg (including the 3.1 mg carbon fiber paper substrate) and thickness of ≈0.1 mm.

Electrolytes Characterization

Density and viscosity of all mixtures were measured at 25 °C in a DSA 5000 M densimeter coupled with a LOVIS 2000 ME module from Anton Paar. The equipment was calibrated by measuring the densities of water and the viscosity of the standards provided by Anton Paar (e.g., APS3, APN26, and APN415). Differential scanning calorimetry (DSC, TA Q2000) analysis was also used to investigate thermal behaviors of mixtures under N₂ atmosphere. During the analysis, samples were securely sealed within aluminum pans and placed inside the calorimeter furnace. For data acquisition, samples were initially cooled from room temperature to −150 °C at a scan rate of 5 °C min⁻¹. Then, two heating/cooling cycles consisting of heating the sample to 30 °C and subsequently cooling it back to −150 °C were accomplished always at a scan rate of 5 °C min⁻¹. ¹H NMR analysis of all samples were conducted using a Bruker Avance 500 spectrometer operating at 500 MHz with a broad‐band‐fluorine observe (BBFO) probe capable of producing z‐axis gradient pulses. The ¹H NMR parameters used included a 30° pulse, an acquisition time of 3.1719 s, a relaxation delay of 1 s, and a total of 16 to 32 scans. All samples were placed in 5 mm diameter NMR tubes and analyzed using deuterated chloroform (CDCl₃) placed in a capillary tube as the external reference. The spectra were acquired after setting the temperature to 25 °C using a Bruker Variable Temperature BVT 3000. The peaks were identified and spectra were processed using the software MestReNova. The UV‐Vis spectra of the leachates were recorded from 300 nm to 800 nm using a Cary Varian 4000 spectrometer. Qualitative and quantitative TXRF analyses were performed with a benchtop S2 PicoFox TXRF spectrometer from Bruker Nano (Germany). TXRF system was equipped with a Mo X‐ray source working at 50 kV and 600 μA, a multilayer monochromator with 80% of reflectivity at 17.5 keV (Mo Kα), a XFlash SDD detector with an effective area of 30 mm², and an energy resolution better than 150 eV for 5.9 keV (Mn Kα). For deconvolution and integration, the commercial Spectra v. 7.5.3 software package from Bruker was used. Li and Co concentrations in the leachates were obtained using an inductively coupled plasma atomic emission spectrometer (ICP‐OES Spectro ARCOS III) after dilution of a small volume of each leachate in 50 mL of 1M HNO₃.

Electrochemical Characterization

Ionic conductivity measurements were measured in a Mettler Toledo SevenExcellence multiparameter with an InLab 731‐ISM sensor (0.01–1000 mS cm⁻¹).

The electrochemical properties of the SCs were evaluated by CV, galvanostatic charge‐discharge tests (GCD), and EIS using a VMP‐3e Biologic electrochemical working station. The cells were preconditioned at 1 V by performing 5 consecutive CVs at a scan rate of 20 mV s⁻¹. Subsequently, CVs and GCDs were performed. All figures represented the 3rd cycle obtained at the ESW of interest. The cycling life of SCs was assessed by performing charge/discharge cycling in the selected voltage under a current density of 4 A g⁻¹ using a LBT21084 battery test system (Arbin Instruments). CVs (at 5 mV s⁻¹) and chronoamperometry (CA, holding each potential over 30 min) were also performed in a three‐electrode configuration with Ag/AgCl (1 M KCl) as the reference electrode. The figures in the three‐electrode CV test represented the 5th cycle. All the experiments were carried out at 20 °C.

The gravimetric capacitance of each electrode based on the GCD data was calculated according to Equation (1)

$$C=2I/(\Delta V_d/\Delta t_d)/m$$ (1)

where C is the specific capacitance of each electrode (F/g), I is the applied current (A), Δt _d is discharge time (s), m is the mass of each electrode (g), and ΔV _d is the discharge voltage (i.e., V_max – IR drop) (V).

The equivalent series resistance (R _{int_IRdrop}, Ω) was calculated by Equation (2)

$$ESR=IRdrop/2I$$ (2)

The total energy density (E, Wh kg⁻¹) and power density (P, W kg⁻¹) of SCs were calculated according to Equation (3) and (4)

$$E=1/2\times C/4\times \Delta V_d^2\times 1/3.6$$ (3)

$$P=3600\times E/\Delta t_d$$ (4)

Conflict of Interest

The authors declare no conflict of interest.

Supporting information

Supplementary Material

Xu Boren, Ferrer María L., del Monte Francisco, Gutiérrez María C.. ChemSusChem. 2025; 18, e202501306. 10.1002/cssc.202501306

Data Availability Statement

The data that support the findings of this study are available from the corresponding author upon reasonable request.