# Electrochemical Long-Term Stability of a Choline Chloride-Based Deep Eutectic Solvent for Silver Recycling

Ángel Fernández-Blanco, Benjamin Rotonnelli, Sophie Legeai, Hervé Fontaine, Emmanuel Billy, Romain Duwald

► To cite this version:

Ángel Fernández-Blanco, Benjamin Rotonnelli, Sophie Legeai, Hervé Fontaine, Emmanuel Billy, et al. Electrochemical Long-Term Stability of a Choline Chloride-Based Deep Eutectic Solvent for Silver Recycling. ChemElectroChem, 2025, 10.1002/celc.202500325. hal-05354114

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# Electrochemical Long-Term Stability of a Choline Chloride-Based Deep Eutectic Solvent for Silver Recycling

Ángel Fernández-Blanco, Benjamin Rotonnelli, Sophie Legeai, Hervé Fontaine, Emmanuel Billy, and Romain Duwald*

Deep eutectic solvents (DES) are investigated since 20 years as ionic liquids alternative for electrochemical applications due to their reduced toxicity and cost. Metal electrowinning, is a process which is especially attractive for DES, since it allows the replacement the polluting state-of-the-art hydrometallurgical routes. Though, the solvent interest is conditioned to the DES long term stability under operating conditions. In this article experimental proofs of the long term stability of Propeline 1:3 both in storage and under electrochemical operation conditions, is provided using silver electroleaching-electrodeposition process as a model electrochemical system for the DES aging study. The solvent long term stability is evaluated during prolonged storage without

electrochemical stress, providing a baseline for long-term usability. The electrochemical stability window of Propeline 1:3 is then determined, and its electrochemical degradation is forced for days in order 1) to identify degradation products and 2) to understand degradation mechanism that DES can encounter during extreme operating conditions. Finally, Propeline 1:3 long-term stability under operation conditions is evaluated using different electroleaching-electrodeposition current densities by the monitoring of degradation product formation. The impact of DES aging on the performances of silver electrowinning is assessed through a comparative analysis of fresh and aged electrolyte.

## 1. Introduction

Silver metal has numerous industrial applications, including electrically conductive adhesives for electronics,[1] solar energy technologies,[2] and medical devices.[3] Since 2021, annual global demand for silver has exceeded supply.[4] Although almost 18% of silver supply is provided by recycling, improving this ratio remains essential to securing the future supply. Fortunately, numerous technologies for recovery of silver from primary or secondary sources are available. Traditionally, silver is recovered from secondary sources, such as waste electronic equipment, by pyrometallurgical processes.[5] This method is highly energy intensive and yields silver with a purity up to 99%. Hydrometallurgical technologies allow to recover silver from complex mixtures by dissolving the metal, usually followed by precipitation as silver chloride (AgCl).[6] Nevertheless, hydrometallurgical methods often require

large volumes of concentrated acids or caustic solutions and generate important amounts of effluents.

Since the first report of deep eutectic solvents (DESs) by Abbott et al.,[7] this class of solvent has attracted a significant interest from the scientific community as a costeffective and alternative to ionic liquids. DESs typically consist of a mixture of a quaternary ammonium salt, which acts as a hydrogen bond acceptor (HBA), and a second molecule, which acts as a hydrogen bond donor (HBD). The electrochemical properties of choline chloride (ChCl) based DESs have been widely investigated. One of the most studied eutectic mixture is composed of a mixture of ChCl and ethylene glycol (EG) in a 1:2 molar ratio respectively, so-called Ethaline (Et 1:2).[8] Compared to aqueous systems, DESs offer a wider electrochemical stability window (ESW). Despite their high proton concentration, proton reduction is typically not observed due to strong interactions between protons and chloride anions. In addition to their wide electrochemical window, DESs offer good conductivity and typically high metal ion solubility, and have making them suitable for various applications, including metals leaching, solvent extraction, electroplating, metals electropolishing and electrowinning.[9–11] The mechanism of Ag electrodeposition in a DES made of ChCl:urea in a 1:1 molar ratio was investigated and compared to the corresponding in aqueous media.[12,13] These studies conclude that the silver reduction process is controlled by diffusion, meaning there is a direct link between the DES composition and the silver deposition rate and morphology. Recently, electrochemical properties of propeline 1:3 (Pr 1:3), a DES composed of ChCl and propylene glycol (PG) in a molar ratio of 1:3 respectively, was deeply studied, and the diffusion coefficients of Ag(I), Pd(II) and Au (I) salts were determined.[14] It was shown that the

Á. Fernández-Blanco, H. Fontaine, E. Billy, R. Duwald

Université Grenoble Alpes

CEA, Liten, DTNM

38000 Grenoble, France

E-mail: romain.duwald@cea.fr

B. Rotonnelli, S. Legeai

Université de Lorraine

CNRS, IJL

F-57000 Metz, France

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Supporting information for this article is available on the WWW under https://doi.org/10.1002/celc.202500325

<figure><img src="image_4.png" /></figure>

© 2025 The Author(s). ChemElectroChem published by Wiley-VCH GmbH. This is an open access article under the terms of the Creative Commons Attribution License, which permits use, distribution and reproduction in any medium, provided the original work is properly cited.
<figure><img src="image_5.png" /></figure>

diffusion coefficient of the metal complexes was similar to the one measured in Et 1:2. The lower toxicity of PG compared to EG is an advantage in the development of the process at a higher scale.

Despite the fact that DES have been the subject of extensive study in academic research, there have been no industrial breakthroughs to date. The main limitations hindering the use of DES in the metallurgical industry include their limited chemical stability under metallurgical process conditions and the difficulties associated with recycling and reuse.[15–17]

In this context, we performed a study to evaluate the stability of Pr 1:3 under various conditions. First, we evaluate the thermal stability of Pr 1:3 during prolonged storage in the absence of electrochemical stress in order to provide a baseline of the solvent aging. This is followed by an electrochemical characterization of the system to define the ESW of the DES. Electrochemical degradation is then investigated in the absence of any redox-active species by applying a constant current density to simulate extreme operation conditions. Degradation compounds in both, gas and liquid phase, were identified to give an understanding of the degradation mechanisms. Afterwards, the long-term stability of Pr 1:3 is studied during prolonged silver electrodeaching-electrodeposition (EL-ED) cycles, representative of reasonable long-term operation conditions. The last section is dedicated to quantify the change of physico-chemical properties such as viscosity, density and electrical conductivity to evaluate the overall impact of degradation.

## 2. Results and Discussion

### 2.1. Thermal Degradation

The first stage of this study focused on the stability of Pr 1:3 under thermal conditions in the absence of any electrochemical influence. To this end, a sample was stored at 50 °C in a closed plastic bottle for 7 months. The GC-MS analysis comparing the aged sample to freshly synthesized Pr 1:3 is shown in **Figure 1**. Several degradation products were already detected in the freshly prepared Pr 1:3, including 2-chloroethanol 5 at Rt = 11.31 min, 1,1'-oxybis(propan-ol) 14 at Rt = 15.25 min and ethylene glycol 12 at Rt = 14.09 min. These compounds are likely formed during the synthesis process itself, which involves heating at 80 °C. After 7 months of storage, the only additional degradation compound identified was acetic acid 7 at Rt = 12.06 min. The formation of these degradation products arise from multiple degradation pathways involving choline chloride and propylene glycol. The precise mechanisms leading to their formation are further discussed in the Electrochemical degradation section. Based on the absence of additional degradation products after 7 months, it can be concluded that Pr 1:3 exhibits a good thermal stability at 50 °C.

Additionally, a qNMR study determined that the concentration of ChCl remained stable at 2.8 M, while the PG concentration decreased from 8.7 to 8.4 M. This loss represents ≈4% of the initial mass of propylene glycol and can be mostly attributed to PG

<figure><img src="image_6.png" /><figcaption>**Figure 1.** Normalized GC-MS chromatograms of freshly synthesize Pr 1:3 (orange) and after 7 months of storage at 50 °C (blue).</figcaption></figure>

condensation on the flask wall, due to the difference of temperature between the solution and the glass vessel.

Although prolonged thermal storage resulted in minimal degradation, more severe conditions, such as those encountered during electrochemical processes, may induce the formation of additional degradation products. Therefore, the electrochemical degradation of Pr 1:3 is investigated in the following sections.

### 2.2. Electrochemical Characterization

Prior to evaluating electrochemical degradation, the ESW of Pr 1:3 was characterized to define its operational limits. For this purpose, cyclic voltammetry (CV) of dry Pr 1:3 were recorded at 60 °C over various potential ranges using glassy carbon rod electrodes with a submerged area of 3.2 cm². The measurements were performed by sweeping first toward cathodic potentials at a scan rate of 20 mV s⁻¹. As shown in **Figure 2a** (inset, black plot), no redox signals were observed between –1.5 and 1 V versus Ag/AgCl, defining an ESW of 2.5 V for Pr 1:3. Extending the potential range beyond this window leads to the appearance of additional oxidation, A₁ and A₂, and reduction, C₁, peaks, indicating solvent degradation. These features evolve across cycles and are attributed to the formation and transformation of degradation products. A more detailed description is provided in the Figure S2, Supporting Information.

A comparison between dry (<0.1wt% H₂O) and standard (ca. 1.5 wt% H₂O) Pr 1:3 is shown in Figure 2b for the potential range inside the electrochemical window [–1.5, 1] V versus Ag/AgCl. Two reduction peaks, C₂ and C₃, are observed at –0.5 and –0.9 V versus Ag/AgCl, respectively. These peaks are absent in dry and degassed Pr 1:3 with no oxygen or water content, indicating that their appearance is related to the presence of oxygen and/or water. Additionally, a decrease in the anodic limit of the ESW is clearly evidenced in standard Pr 1:3 by the fast increase in current density from 0.03 to 0.11 mA cm⁻² when reaching 1 V versus Ag/AgCl (Figure 2b), compared to dry Pr 1:3. Finally, a new oxidation peak A₃ appears in standard Pr 1:3 above 0.7 V
<figure><img src="image_7.png" /></figure>

<figure><img src="image_8.png" /><figcaption>**Figure 2.** a) Cyclic voltammograms of dry Pr 1:3 within the ESW (black) and at extended potentials during the first (red) and second cycle (blue) and b) comparison between dry (black) and standard (green) Pr 1:3 containing ca. 1.5 wt% water. All scans correspond to the second cycle unless otherwise specified. Measurements were recorded at 20 mV s⁻¹, sweeping first toward cathodic potentials.</figcaption></figure>

versus Ag/AgCl. While this peak may reflect the oxidation of degradation products generated during cathodic processes as occurs in Figure 2a red plot), it is also possible that its appearance is due to surface redox reactions triggered by the presence of water and/or oxygen, without electrolyte degradation.

With the ESW defined, the following section explores the forced electrochemical degradation of Pr 1:3 in the absence of any redox-active species. Investigating the degradation products and decomposition pathways under such conditions is essential to understand how the electrolyte may behave when pushed beyond its stability limits, for example, due to process malfunctions or extreme electrochemical conditions.

## 2.3. Electrochemical Degradation

Working outside the ESW can happen during process malfunction or when imposing harsh redox conditions (high potential or high

current density) and force the DES degradation. Investigating the degradation products and decomposition pathways under such conditions is essential to understand how the electrolyte may behave when pushed beyond its stability limits, and provides a list of the eventual degradation products that may appear due to electrochemically induced DES aging.

The long-term electrochemical degradation of Pr 1:3 was initially performed in the absence of metal species over glassy carbon electrodes. Chronopotentiometry (CP) experiment was performed at 1 mA cm⁻² over 144 h (6 days) using the electrochemical cell 1. In the absence of redox-active species, both the cathode and anode potentials stabilized outside the ESW of Pr 1:3, inducing its decomposition. As shown in Figure 3, the potential measured at the anode gradually increased from 1.5 to 1.9 V versus Ag/AgCl over the first 44 h, after which it remained stable. At the cathode, the potential remained stable at 1.5 V versus Ag/AgCl for ca. 40 h, after which it decreased until reaching -2.4 V versus Ag/AgCl.

In order to identify the degradation products formed during the electrolysis of Pr 1:3, aliquots were taken at different times and analyzed by GC-MS. Over time, the concentration of degradation products increased, and after 144 h, more than 20 different by-products were detected (Figure S3, Supporting Information). In this section, a description of the compounds identified after 144 h (Figure 4) is discussed, focusing on the most concentrated or those that require special attention because of their potentially harmful properties.

As previously reported, the quaternary ammonium cation of ChCl may undergo a reversed Menshutkin's reaction upon electrochemical reduction[18] or a thermal decomposition,[16] forming either dimethylaminoethanol 1 and a methyl radical or trimethylamine and an ethanol radical, depending on whether an electron is added to one of the three methyl groups or to the hydroxyethyl group of choline, respectively. Consistent with this mechanism, dimethylaminoethanol 1 was detected as a broad peak, appearing in the chromatogram between 8.6 and 9.8 min (Figure 1). This peak continued to expand until the end of the chromatogram, leading to the presence of a base peak with

<figure><img src="image_9.png" /><figcaption>**Figure 3.** (Red) anode and (blue) cathode potentials during chronopotentiometry of Pr 1:3 at 1 mA cm⁻² over 144 h.</figcaption></figure>
<figure><img src="image_10.png" /></figure>

<figure><img src="image_11.png" /><figcaption>**Figure 4.** GC-MS chromatogram of Pr 1:3 showing identified compounds after 144 h of chronopotentiometry at 1 mA cm⁻² and 60 °C.</figcaption></figure>

$m/z = 58$. The same compound was also identified during the thermal degradation study (Figure 1).

In general, for DES containing anions, the anion degradation occurs at the anodic limit. More specifically, the oxidation of chloride ions $Cl^-$ into trichloride ions $Cl_3^-$ was reported several times in chloroaluminate ionic liquids[19] and other chloride-containing ionic liquids.[20-22] The same behavior is expected to occur in a chloride-rich medium like Pr 1:3, as observed for Et 1:2, where UV-Vis measurements upon electrolysis revealed a band around 240 nm associated to $Cl_3^-$ ions.[18] This is in good agreement with other measurements for $Cl_3^-$ in chlorinated-rich mediums.[18,21] A reaction between trichloride anions $Cl_3^-$ and the ethanol radical, via addition at the antiMarkovnikov position, can lead to the formation of 2-chloroethanol 5, which was detected at Rt = 11.48 min. 2-chloroethanol is a highly toxic compound[23] and is already present in the fresh Pr 1:3 at 0 h (Figure S3, Supporting Information). Its early presence, both in the electrochemical degradation and the thermal stability study (Figure 1), suggests that its formation may also proceed through a direct reaction between chloride anions $Cl^-$ and the ethanol radical, without necessarily involving $Cl_3^-$, which would not be present at significant levels in the fresh electrolyte. Over time, its concentration appears to decrease (Figure S3, Supporting Information), likely due to its transformation into degradation by-products. The same ethanol radical can explain the formation of ethylene glycol 12 after hydrolysis of the radical with residual water in the DES. Ethylene glycol was detected in the freshly synthesized Pr 1:3 during the thermal stability study (Figure 1), but could not be confirmed in the electrochemical degradation samples. For this reason, label 12 is not indicated in Figure 4.

ChCl can also undergo a Hoffman elimination, resulting in trimethylamine and acetaldehyde.[16,18] Trimethylamine was not detected, and acetaldehyde 16 was identified in the gas phase (see below) and in the liquid phase as acetic acid 7 upon further oxidation at Rt = 12.09 min. Acetic acid 7 was also detected as a degradation product after 7 months of storage in the thermal stability study (Figure 1). In addition, degradation by-products derived from acetic acid were observed, notably methyl acetate 10 at Rt = 13.21 min, likely formed by esterification of 7 with methanol.[24]

Propylene glycol 11 was detected at Rt = 13.48 min as the most intense peak. At the anode, oxidation of the secondary alcohol of PG results in hydroxyacetone 2 at Rt = 10.59 min. Both compounds were confirmed by the injection of the pure substances. Hydroxyacetone 2 can further react with PG to form a cyclic ketal (2,4-dimethyl-1,3-dioxolan-2-yl)methanol (Figure 5), a well-known reaction in organic chemistry used to protect carbonyl groups from nucleophilic attacks.[25] This compound has two stereogenic centers at carbons 2 and 4, leading to a pair of diastereomers 9 and 9' distinguishable by GC-MS at Rt = 12.94 min and Rt = 13.11 min, respectively, with the most intense fragments at $m/z = 101$ and 43, appearing as the most concentrated degradation compound. To confirm the presence of these compounds, a reaction between hydroxyacetone and PG was conducted separately by heating the mixture at 80 °C for 1 h in the presence of 1% of HCl catalyst. The reaction mixture was then analyzed by GC-MS, showing two peaks at Rt = 12.94 min and Rt = 13.11 min and with the same $m/z$ pattern, confirming the origin of compounds 9 and 9' (Figure S4, Supporting Information). Additionally, both diastereomers were successfully identified by nuclear magnetic resonance spectroscopy (NMR).
<figure><img src="image_12.png" /></figure>

<figure><img src="image_13.png" /><figcaption>**Figure 5.** Proposed oxidation pathway of propylene glycol at the anode. Propylene glycol is first oxidized to hydroxyacetone, which can undergo tautomerization to form lactaldehyde. Both hydroxyacetone and lactaldehyde can further react with excess of propylene glycol to yield the corresponding dioxolane derivatives.</figcaption></figure>

confirming the formation of the dioxolane (Figure S5, Supporting Information). Interestingly, the presence of two diastereomers (9 and 9') should exhibit two different peaks on the chromatogram. Nevertheless, the condensation of hydroxyacetone and PG shows 3 different peaks. In fact, hydroxyacetone can also undergo tautomerisation (Figure 5), reaching an equilibrium between the enol and enediol forms through 1,2- and 2,3-enolization, respectively. The enediol tautomer can further equilibrate with 2-hydroxypropionaldehyde (lactaldehyde).[26] Lactaldehyde was not directly detected, but lactic acid 3 formed by its oxidation is present at Rt = 10.97 min. In a similar way than hydroxyacetone, lactaldehyde can further react with propylene glycol forming the acetal 1(4-methyl-1,3-dioxolan-2-yl)ethan-1-ol, which appears as a couple of diastereomers 8 and 8' at Rt = 12.74 min and Rt = 12.97 min, respectively, 8' overlapping with (2,4-dimethyl-1,3-dioxolan-2-yl)methanol 9. The same peaks appear during the reaction of hydroxyacetone and propylene glycol conducted in parallel, confirming the involvement of hydroxyacetone in the formation of these dioxolanes (Figure S4, Supporting Information, top). The most intense signal in the mass spectra corresponds to *m/z* = 87, likely representing the 4-methyl-1,3-dioxolane formed by ionization of 1-(4-methyl-1,3-dioxolan-2-yl)ethan-1-ol (Figure S4, Supporting Information, bottom).[27]

The trichloride ions $Cl_3^-$ can induce the chlorination of some degradation products, in particular the cyclic ketal (2,4-dimethyl-1,3-dioxolan-2-yl)methanol, resulting in 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane 4, identified at Rt = 11.27 min. To confirm

this, the previously synthesized (2,4-dimethyl-1,3-dioxolan-2-yl) methanol was isolated by liquid-liquid extraction using cyclohexane, and then chlorinated in the presence of $PCl_5$ by heating at 60 °C for 1 h. The resulting GC-MS shows the presence of a peak at the same retention time at 11.27 min and the same major fragments at *m/z* = 101 and 43 (Figure S6, Supporting Information), confirming the nature of compound 4. In the same way, the corresponding dioxolane formed by acetalization of lactaldehyde and propylene glycol is chlorinated forming 2-(1-chloroethyl)-4-methyl-1,3-dioxolane 6 seen at Rt = 11.65 min. Although the stereoisomerism of both dioxolanes remains after chlorination, both compounds are detected as a single peak in the chromatogram, the GC-MS program not being suitable for the separation of all the stereoisomers. No information was found on the toxicity of these compounds. However, chlorinated compounds are usually known for their toxicity and potential environmental persistence.[28] Therefore, special attention should be paid to them.

The electrochemical reduction of PG can lead to the formation of propylene oxide 17,[29] detected in the gas phase (see below), along with the evolution of hydrogen or water (dehydration). Propylene oxide 17 is susceptible to nucleophilic attack due to the significant angle strain in its three-membered ring. In the presence of an excess of PG, a hydroxypropylation occurs, resulting in the formation of 1,1'-oxybis(propan-ol) 14, observed at Rt = 15.39 min. This compound was detected before the electrochemical aging of Pr 1:3 at Rt = 0 h (Figure S3, Supporting Information) and during the thermal stability study (Figure 1), indicating that it is formed during the synthesis of Pr 1:3. In this
<figure><img src="image_14.png" /></figure>

eventuality, propylene oxide would be likely formed by a thermal
dehydration of PG. Finally, in region 13 (Rt = 13.70–15.00 min) of
the chromatogram, numerous degradation products appear as
low intensity, overlapping peaks, making their identification
particularly complicated. The first analysis of mass spectra of
each peak tends to indicate the presence of acetates, dioxolanes,
dioxanes, and ether derivatives, suggesting that hydrolysis, trans-
esterification, acetalization, or ketalization reactions are occur-
ring. These processes involve PG, propylene oxide, and its
degradation products, leading to a complex mixture of oligo-
meric structures.

Only a few of the degradation compounds discussed above were detected during the first hours (≤5 h) of chronopotentiometry (Figure S3, Supporting Information), including 2-chloroethanol 5, ethylene glycol 12, 1,1'-oxybis(propan-ol) 14, hydroxyacetone 2, (2,4-dimethyl-1,3-dioxolan-2-yl)methanol 9/9', 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane 4 and 1-(4-methyl-1,3-dioxolan-2-yl)ethan-1-ol 8/8'. The absence of 2-(1-chloroethyl)-4-methyl-1,3-dioxolane 6 or the low intensity peak of 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane 4 suggest that chlorination becomes more favorable with longer degradation times. This is expected, as their formation is limited at the beginning due to the low concentration of their precursors.

Finally, a qNMR analysis was performed to estimate the order of magnitude of the concentration of degradation compounds formed. The concentration of PG decreased from an initial value of 8.0–7.4 M after 144 h of chronopotentiometry, representing ca. 7.5 wt% of the initial mass due to evaporation and degradation. Meanwhile, the concentration of the choline cation only slightly changed, remaining at ca. 2.8 M. As a result, the molar ratio of propeline decreased from Pr 1:2.9 to Pr 1:2.6. However, it should be noted that the quantification of both species, specially the choline cation, may be affected by partial overlap of their characteristic ¹H NMR signals with those of certain degradation products, introducing a degree of uncertainty in the reported concentration values. To better understand the losses due to degradation and vaporization, an analysis of the gas phase is presented in the following section.

## 2.4. VOC Analysis

During the chronopotentiometry experiment, the initially colorless Pr 1:3 gradually turned to light-yellow after 2 h. Throughout this period, a gas generation was observed at the cathode. In the study of electrochemical degradation of Et 1:2, Haerens et al.[18] observed the same phenomenon and attributed it to hydrogen evolution coming from residual water electrolysis. Later, Sinclair et al.[30] also attributed the gas as hydrogen but concluded that its primary source was the reduction of ethylene glycol at the cathode, rather than solely water reduction. Given the structural similarities between Et 1:2 and Pr 1:3, one can expect the same reaction to occur in the present system. As the experiment progressed, the color of Pr 1:3 continued to change. After 2 days, the color darkened from yellow to orange-brown, eventually turning brown-black, making further gas evolution difficult to observe.

To analyze the degradation compounds present in the gas phase, a Markes Air Toxic Analyzer adsorbent tube was connected to the N₂ gas outlet of electrochemical cell 1. Gas samples were collected over ≈40 min at 5 and 30 h of electrolysis. The trapped compounds were then released by thermal desorption and subsequently analyzed using GC-MS/FID.

Several degradation compounds identified in the gas phase were already detected in the liquid phase, including, acetic acid 7, hydroxyacetone 2, (2,4-dimethyl-1,3-dioxolan-2-yl) methanol 9/9', 1-(4-methyl-1,3-dioxolan-2-yl)ethan-1-ol 8/8' and 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane 4. The conditions and chromatographic column employed in this analysis enabled the separation of diastereomers of the latter compound into 4 and 4' (Figure 6). Among the newly detected compounds, acetaldehyde 16 at Rt = 3.39 min and propylene oxide 18 at Rt = 5.23 min provide further support for the mechanisms proposed in the previous section. The detection of chloromethane 15 at Rt = 3.00 min and chloroform 21 at Rt = 10.15 min can be explained by the recombination of the methyl radicals, formed during reduction of ChCl at the cathode, with chloride ions (Cl⁻).[18] Additionally, another

<figure><img src="image_15.png" /><figcaption>**Figure 6.** GC-MS chromatogram of the gas phase of Pr 1:3 after 30 h of chronopotentiometry at a current density of 1 mA cm⁻². Identified VOCs are labeled on the chromatogram.</figcaption></figure>
<figure><img src="image_16.png" /></figure>

chlorinated compound, chloroacetone 23 was identified at Rt = 14.07 min. Its formation is likely due to chlorination of PG or the enolic form of acetone by chlorine ions.

Several dioxolane derivatives were also detected. The formation of some of them can be reasonably explained, such as 4-methyl-1,3-dioxolane **22** at Rt = 11.67 min or 2,2,4trimethyl-1,3- dioxolane **24** at Rt = 14.20 min, which result from the combination of propylene glycol **11** with formaldehyde or acetone **20**, respectively. Acetone **20** was detected at Rt = 5.65 min, while formaldehyde was not detected under the current TD-GC-MS conditions, which are not optimized for very volatile compounds. Nevertheless, their formation could be expected through dehydration and CC bond cleavage in hydroxyacetone **2**,[31] respectively. The identification of other dioxolane derivatives is not conclusive, and their formation pathways are unclear. Therefore, they were excluded from this analysis. A list with all the degradation compounds identified in gas an liquid phase can be found in Table S1, Supporting Information.

A semi-quantitative analysis was then performed from FID using toluene as a reference to construct a calibration curve based on peak areas and concentrations. The measured areas of the identified compounds were used to estimate their mass relative to toluene. Since the FID is sensitive to carbon atoms, compounds with less carbon atoms than toluene, for example, PG, lead to an underestimation of their actual concentrations. While this analysis is not fully quantitative, it provides a reasonable estimation of the order of magnitude of concentrations. Using this method, 159 µg L⁻¹ of PG was detected in the gas phase, representing 61% of the total VOC, while degradation compounds accounted for 39% with a 100 µg L⁻¹ concentration. These values, obtained after 44 min of sampling, were used to extrapolate the evaporation losses over 144 h of experimentation. The evaporated PG represents ca. 0.24 wt% of the initial PG mass in the DES, while the evaporated degradation products correspond to ca. 0.09% of the initial DES mass. As a reminder, an inert atmosphere of N₂ gas at flow rate of 40 sccm was applied to the electrochemical cell, causing a faster evaporation of PG and degradation products compared to conditions without gas. These results, together with the qNMR data presented in the previous section, indicate that PG was primarily degraded into various compounds (7.26 wt%), while only a small fraction was lost by evaporation (0.24 wt%). The most concentrated compound detected in the gas phase, after PG, was 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane **4**, representing 16% of the total gas phase, with a concentration of 42 µg L⁻¹ (or 6.68 ppmv). No European reference for occupational exposure limit has been established for 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane. Therefore, no definitive conclusions regarding its safety can be done at this point and a standard exposure control should be implemented as a precaution, including proper ventilation and effective gas waste treatment to prevent the accumulation of degradation compounds. The formation of those additional products is nonetheless undesirable and roots from parasitic redox reactions that decreases the reaction faradaic efficiency. Avoiding forming those products is therefore preferable, which requires the reaction to stay in the ESW region.

## 2.5. Stability Study during Ag Electroleaching and Electrodeposition

We identified the degradation products that may be formed outside the ESW window on glassy carbon electrode. Though, the use of metallic electrodes may catalyze the solvent degradation even in reasonable conditions. To detect the eventual formation of those species, a series of simultaneous EL-ED experiments using silver foil electrodes were performed under increasing current densities. The electrochemical cell 2 was used for this purpose, with solutions of Pr 1:3 containing 70 mM AgCl to initiate Ag deposition from the start of the ED process and to prevent electrolyte degradation at the cathode due to the initial absence of Ag⁺ ions. The immersed surface area of both the anode and cathode was ca. 6 cm² (15 × 20 mm per face). In this configuration, silver dissolution occurred at the anode, while silver deposition took place at the cathode. Fresh Pr 1:3 samples were used for each experiment, and the electrodes were carefully polished using RS Pro silicon carbide emery paper (grit P2000, ≈10 µm particle size) and cleaned with ethanol before use.

A first CV was recorded to evaluate Ag oxidation and reduction potentials over silver electrodes in Pr 1:3. The results are depicted in **Figure 7** for the working electrode at a potential range between -1.5 and 2 V versus Ag/AgCl starting from open circuit potential and sweeping first toward cathodic potentials with a scan rate of 20 mV s⁻¹. During the first cycle, a small reduction peak C1 is observed at -0.2 V versus Ag/AgCl, corresponding to the reduction of Ag⁺ ions present in solution. This is followed by a broad oxidation peak A1 centered at 1.2 V versus Ag/AgCl, attributed to the oxidation of silver. These features indicate that both the EL and ED of silver are taking place. The decrease in anodic current observed after 0.4 V may reflect diffusion limitations of the Ag(I) species formed toward the bulk solution, leading to the formation of AgCl on the metal surface and a limitation of the anodic current. Scanning electron microscopy (SEM) and

<figure><img src="image_17.png" /><figcaption>**Figure 7.** First cycle of the cyclic voltammogram of Pr 1:3 with 70 mM AgCl, recorded at 60 °C and a scan rate of 20 mV s⁻¹. The potential was initially swept toward cathodic values.</figcaption></figure>
<figure><img src="image_18.png" /></figure>

EDX confirm the presence of AgCl in some regions of the working electrode surface (Figure S7, Supporting Information). The formation of AgCl during silver oxidation in rich chloride aqueous solutions is well-known and has been extensively studied[32–34] and reported in similar DES media.[35,36] If the faradaic yield is quantitative, the higher the current density, the faster the EL-ED process. To determine the maximum current density for Ag EL-ED in our system without solvent degradation, chronopotentiometry experiments were performed for two hours using different current densities (Table 1 and 2). The recording of anodic and cathodic potentials over time, as well as the faradaic yield computation, depending on the applied current will provide information regarding eventual side reactions that may happen. After experiments, the experimental mass of electrodeposited silver has been measured using an analytical balance Ohaus EX324AD after the cleanning and drying procedure described above, and allowed the computation of the cathodic and anodic reactions yield following the Faraday’s law (see Equation (1), Δm<sub>exp</sub> being the electrode experimental mass variation, Q the total charge that went through the electrode, F the Faraday constant). The respective yield values are given in Table 1.

The first experiment was performed at 1 mA cm⁻². The potentials at the anode and cathode remained stable at 0.075 V versus Ag/AgCl and –0.1 V versus Ag/AgCl, respectively (Table 2). These potentials correspond to the dissolved Ag(I) reduction and silver metal oxidation, according to the CV presented on Figure 7, which remain within the ESW of the DES, and are therefore representative of suitable electrochemical conditions for performing EL-ED in the system. After 2 h, the anode surface exhibited a shiny silver color. SEM images show a homogeneous silver surface, with localized regions containing AgCl, observed as distinct features of various sizes ranging between 6 µm and 370 nm, as confirmed by EDX analysis (Figure S8, Supporting Information). AgCl presence was also observed on top of the silver deposits at the cathode. These AgCl features may result from precipitation during the cleaning step, which involved repeatedly rinsing the electrodes in 96% ethanol (Carlo ERBA Reagents), followed by rinsing with deionized water, and drying at 60 °C overnight.

According to Faraday’s law, at a given current density, it is possible to estimate the theoretical mass of silver that will be dissolved at the anode or deposited at the cathode. The faradaic oxidation

<figure><img src="image_19.png" /></figure>

and reduction yields can then be calculated by comparing the experimental mass change (Δm<sub>exp</sub>) with the theoretical value derived from the total charge passed (Q). The yield is given by

<table><caption>**Table 1.** Quantitative results of chronopotentiometry experiments summarize in Table 2.</caption><thead><tr><th>j [mA cm<sup>-2</sup>]<sup>a)</sup></th><th>Time [h]</th><th>Charge [C]</th><th>Theoretical Ag<br>leached/deposited [mg]</th><th>Δm anode [mg]</th><th>Faradaic oxidation<br>Yield [%]</th><th>Δm cathode [mg]</th><th>Faradaic reduction<br>yield [%]</th></tr></thead><tbody><tr><td>1</td><td>2</td><td>43.2</td><td>48.3</td><td>-48.3</td><td>≈100</td><td>49.2</td><td>102</td></tr><tr><td>2</td><td>2</td><td>86.4</td><td>96.2</td><td>-95.7</td><td>99</td><td>99.4</td><td>103</td></tr><tr><td>2.6</td><td>2</td><td>112.3</td><td>125.6</td><td>-95.2</td><td>76</td><td>138.5</td><td>110</td></tr><tr><td>2<sup>b)</sup></td><td>476</td><td>20,563.2</td><td>22,989.7</td><td>9.6</td><td>≈100</td><td>-30.8</td><td>≈100</td></tr><tr><td>2<sup>c)</sup></td><td>2</td><td>43.2</td><td>48.3</td><td>-3</td><td>6</td><td>0.4</td><td>Not applicable</td></tr></tbody></table>

<footnote><sup>a)</sup>The table includes the applied current density, *j* (mA cm⁻²), total experiment duration (h), total charge passed (C), theoretical silver oxidation based on Faraday’s law (mg), measured mass change of the anode and cathode (mg), as well as the calculated faradaic oxidation and reduction yield (%). Yields higher than 100% can be explained by the presence of chloride at the surface of silver metal. These observations are confirmed by EDS spectra (see Figure S8, S9 and S11, Supporting Information); <sup>b)</sup>Current density applied for 476 h reversing the polarity each 2 h; <sup>c)</sup>Current density applied to degraded Pr 1:3 at 1 mA cm⁻² for 6 days.</footnote>
<figure><img src="image_20.png" /></figure>

$$ \eta = \frac{\Delta m_{\text{exp}}}{M_{\text{Ag}} * \frac{Q}{F}} \quad (1) $$

where MAg is the molar mass of silver and F is the Faraday constant. The calculated oxidation and reduction yields are presented in Table 1. The experimental mass of electrodeposited silver has been measured using an analytical balance Ohaus EX324AD after the cleanning, and drying procedure described above and compared to theoretical values (Table 1). From 1 to 2.6 mA cm⁻² experimental values for Ag electrodeposition are consistent with a quantitative faradaic yield (Table 1, entries 1–3). In some cases, the experimental faradaic yields slightly exceed 100%, reaching up to 110% at 2.6 mA cm⁻². While part of this deviation (0.09%–0.2%) can be attributed to the weighing uncertainty of the balance (±0.1 mg), other factors like an incomplete removal of DES can explain the observed deviation. In all cases, on the cathode, silver deposits were clearly visible to the naked eye at the end of the experiment (Figure 8a). SEM images revealed rounded silver particles along with some dendritic structures (Figure 8b,c), which became more abundant near the electrode edges.

When the experiment was carried out under a current density of 2 mA cm⁻², the measured potentials at the anode and cathode were slightly higher than for the lower current density at 1 mA cm⁻², but remained stable at values of ca. 0.15 V versus Ag/AgCl and -0.25 versus Ag/AgCl, respectively (Table 2, entry 2). These potentials still correspond to the Ag oxidation and Ag(I) reduction potentials. Interestingly, the surface of the anode darkened during the experiment, and SEM images revealed an irregular and rough silver morphology. When the current density was increased to 2.6 mA cm⁻², a continuous rise of the anode potential was observed, doubling from ca. 0.3 V versus Ag/AgCl to 0.6 V vs. Ag/AgCl over the course of 2 h (Table 2, entry 3). In contrast, the cathode potential remained stable at around -0.1 V versus Ag/AgCl. As for the previous experiment at 2 mA cm⁻², the anode surface turned black, and SEM images revealed an irregular and rough morphology in the central region of the electrode (Figure S9a, Supporting Information). A significant formation of AgCl was also observed near the electrode edges (Figure S9b, Supporting Information), which could explain the increase in the anode potential, as previously reported for AgCl formation

<figure><img src="image_21.png" /><figcaption>Figure 8. Visual and morphological characterization of the CE after chronopotentiometry at 1 mA cm⁻². Similar results were obtained for the other current densities. a) Photograph of the silver electrode showing deposition patterns b) SEM image showing a compact morphology with rounded silver particles and c) SEM image showing the formation of dendritic or crystalline structures.</figcaption></figure>

in silver substrates.[32,34] The difference between the theoretical mass of silver leaching and the experimental mass loss of the anode can also be attributed to the formation of AgCl, which has a higher molecular weight than metallic silver, resulting in a faradaic oxidation yield of 76% (Table 1, entry 3). It is worth noting that in this case, a voluminous silver deposit was formed on the cathode, significantly increasing the thickness of the electrode and complicating the removal of the DES after the EL-ED process, which may explain the 110% faradaic reduction yield.

To further evaluate the operational limit of our system, higher current densities of 4 and 6 mA cm⁻² were applied for several hours (Figure S10, Supporting Information). In both cases, the anode potential increased linearly with time, showing no sign of stabilization and rapidly exceeding the anodic stability limit of the DES. The continuous potential rise suggest that, under these conditions, silver dissolution leads to local saturation of Ag(I) near the electrode surface, eventually exceeding the solubility limit and causing AgCl precipitation. As the electrochemically active surface area decreases, the potential continues to rise, an effect that becomes more pronounced at higher current densities. Meanwhile, the cathode potential remains relatively stable during the experiments, with values around -1.5 V versus Ag/AgCl for 4 mA cm⁻² and slightly more negative for 6 mA cm⁻². Minor fluctuations are observed but no significant drift is detected over time.

These results show that in our system, the speed of silver electrodeaching is limited by the diffusion of Ag(I), leading to the formation of an AgCl layer on the anode when applying a high current density, rather than by the degradation of Pr 1:3. Based on these observations, a current density of 2 mA cm⁻² can be considered optimal for the experimental conditions applied in this study. With further optimization of cell design, agitation, and other operational parameters, it may be possible to increase this current density without pushing the system outside the ESW.

In order to evaluate the long-term stability of Pr 1:3 in a silver EL-ED process at 2 mA cm⁻², 119 electrochemical cycles were performed over a period of 476 h (ca. 20 days). Each cycle consisted of 2 h of anodic polarization at +2 mA cm⁻², followed by 2 h of cathodic polarization at -2 mA cm⁻², enabling alternate oxidation and reduction of silver between the two electrodes. During anodic polarization, the working electrode potential remained relatively stable around +0.2 V vs. Ag/AgCl. In contrast, during cathodic polarization the potential exhibited a slight shift from -0.15 V versus Ag/AgCl at the beginning of the process to -0.2 V versus Ag/AgCl in the final cycles (Table 2, entry 4). Both voltage values ensure that the silver electrowinning occurs quantitatively. At the end of the experiment, both silver electrodes were consumed on the edges and visible silver deposits were formed in the central regions. SEM images of this region show silver deposits made of rounded particles, while dendritic formations were also observed near the edges. As silver is oxidized from one electrode and reduced onto the other during each half-cycle, the overall mass of both electrodes is supposed to remain unchanged at the end of each cycle and therefore after the 476 h. Experimentally, mass differences of +9.6 and -30.8 mg were measured for the WE and CE, respectively (Table 1).
<figure><img src="image_22.png" /></figure>

Considering a theoretical amount of ca. 22,990 mg of silver was
cycled between the electrodes over the course of the experiment,
these values correspond to a faradaic yield of ca. 100%.
Throughout the 20 days, Pr 1:3 remained colorless, and a com-
parison of GC-MS analysis before and after the experiment show
no significant degradation products (Figure 9), apart from those
formed during the synthesis of Pr 1:3. These results confirm that,
under the experimental conditions and set-up used, Pr 1:3 exhib-
its long-term electrochemical stability. Although no degradation
products were detected, qNMR analysis revealed a decrease in PG
concentration, with the molar ratio shifting from 1:3 to ≈1:2 after
20 days. This loss, associated to PG evaporation, was likely
enhanced by the continuous nitrogen flow used to maintain a
dry and inert atmosphere in the electrochemical cell. This evap-
oration highlights the need to replenish PG during long-term
operation, as changes in the molar ratio between the HBA and
HBD can alter the physicochemical properties of Pr 1:3 and con-
sequently impact the performance of the EL-ED process. A final
experiment was conducted using electrochemically degraded Pr
1:3 to compare the silver recovery capacity with that of fresh Pr
1:3. Aged Pr 1:3 was obtained by applying a current density of
1 mA cm⁻² to a fresh batch of Pr 1:3 for 6 days using electrochem-
ical cell 1, which are the same conditions employed in the elec-
trochemical degradation study earlier in this paper (see above).
The Ag EL-ED experiment was then carried out using 3 cm² silver
electrodes and a current density of 2 mA cm⁻² applied for 2 h. The
anode potential rapidly grew to 1.5 V vs. Ag/AgCl (Table 2), indi-
cating the formation of AgCl, which was confirmed by SEM
(Figure S11, Supporting Information). The theoretical faradaic
yield for silver oxidation at the anode was 6% demonstrating
poor performance, likely to a decrease in Cl⁻ ion concentration
caused by the degradation of the DES, although this still needs

to be confirmed. Unlike previous experiments, no additional
70 mM AgCl was added to the electrolyte beforehand, so the far-
adaic reduction yield should not be considered in this case.

Following the evaluation of electrochemical degradation Pr 1:3, we investigate the evolution of key physico-chemical properties throughout the degradation processes in the presence and absences of redox-active species.

## 2.6. Physico-Chemical Properties

Physico-chemical properties such as density, viscosity and elec-
trical conductivity are important parameters when considering
the suitability of DESs for industrial-scale applications.[37] A com-
parative study of these properties for Pr 1:3 and Et 1:2 was pre-
viously reported by Bertoloni et al.[14] at different molar ratios
and water content. In this section, we investigate the evolution
of density, viscosity and electrical conductivity of Pr 1:3 in the
temperature range from 30 to 100 °C during the electrochemical
degradation and the long-term silver EL-ED at 2 mA cm⁻².
To facilitate direct comparison with the extended EL-ED process,
the electrochemical degradation of Pr 1:3 has been extended to
20 days at 1 mA cm⁻².

The measured density of Pr 1:3 is consistent with the previ-
ously reported values,[14,38] showing a decrease from 1.07 to
1.02 g cm⁻³ for 30 and 100 °C, respectively (Figure 10a). At
60 °C, the measured density of Pr 1:3 was 1.05 g cm⁻³. During
electrochemical degradation, this value progressively increases,
reaching 1.08 g cm⁻³ after 20 days. This suggests a structural reor-
ganization, likely due to changes in the hydrogen bonding net-
work or the incorporation of degradation products that pack
more efficiently. Additionally, upon degradation, the ChCl:PG
molar ratio decreased from 1:3 to ca. 1:2 after 20 days (deter-
mined by qNMR). This change in molar ratio reflects the same
trend reported by Bertoloni et al.[14] who reported higher densi-
ties for Pr 1:2 (1.06 g cm⁻³) than for Pr 1:3 (1.05 g cm⁻³) at 60 °C.

In the same way, the presence of 70 mM of AgCl in Pr 1:3
slightly increases the density compared to the salt-free DES from
1.05 to 1.06 g cm⁻³ (Figure 10a, blue and green plot). After
20 days of Ag EL-ED process, the density further increased from
1.06 to 1.07 g cm⁻³. Since no degradation products were
detected after the experiment (see above), the observed change
is likely attributed to a decrease of ChCl:PG molar ratio due to
evaporation of PG.

Viscosity is another important parameter to consider, as it
directly affects mass and heat transfer, therefore impacting reac-
tion kinetics and conductivity, and can pose challenges for prac-
tical operations such as pumping, stirring, and solidliquid
separation.[37] The measured viscosities of Pr 1:3 are in good
agreement with the values reported by Bertoloni et al.,[14] show-
ing a decreasing tendency with increasing temperature from 54.9
to 6.5 mPa.s at 30 and 100 °C, respectively (Figure 10b). For a
given temperature, the viscosity increases as the electrochemical
degradation progresses, passing from 15.3 to 33.0 mPa s after
20 days at 60 °C (Figure 10b, blue and yellow plot). In the same
way, the addition of 70 mM of AgCl to Pr 1:3 slightly increased the

<figure><img src="image_23.png" /><figcaption>**Figure 9.** Comparison of GC-MS chromatograms of freshly synthesized Pr 1:3 (blue) and after 20 days of chronopotentiometry at 2 mA cm⁻² cycling Ag between electrodes (red). Degradation compounds formed are also identified.</figcaption></figure>
<figure><img src="image_24.png" /></figure>

<figure><img src="image_25.png" /></figure>

<figure><img src="image_26.png" /></figure>

<figure><img src="image_27.png" /><figcaption>**Figure 10.** Density a) viscosity b) and electrical conductivity c) measurements of Pr 1:3 at *t* = 0 days (blue), *t* = 7 days (orange), *t* = 14 days (black) and *t* = 20 days (yellow) of electrochemical degradation at 1 mA cm⁻² using electrochemical cell 1. Measurements of Pr 1:3 containing 70 mM of AgCl before (green) and after (red) 20 days of Ag EL-ED using electrochemical cell 2 are also shown.</figcaption></figure>

viscosity from 15.3 to 15.7 mPa s at 60 °C. After 20 days of EL-ED process, the viscosity increased from 15.7 to 20.4 mPa s at the same temperature. Similarly to the density evolution, the changes in viscosity can be explained by a reorganization of the hydrogen

bonding network due to the presence of degradation products, evaporation of PG or the increase of Ag⁺ concentration.

Figure 10c shows the conductivity values measured during the electrochemical degradation. The reported values[14] are in good agreement with those of Pr 1:3 at *t* = 0 days. As expected, the electrical conductivity increases with temperature, rising from 4.1 sM cm⁻¹ at 30 °C to 31.1 sM cm⁻¹ at 100 °C. This trend is attributed to the decrease in viscosity at higher temperatures, which allows faster ion diffusion. Upon degradation, the conductivity at 60 °C decreased from 11.9 to 8.3 sM cm⁻¹. In contrast, the addition of 70 mM of AgCl led to a slight increase in conductivity to 12.7 sM cm⁻¹ at 60 °C due to the introduction of additional charge carriers compared to the salt-free DES. After 20 h of EL-ED, the conductivity at 60 °C remained nearly unchanged, with a value of 12.4 sM cm⁻¹.

### 3. Conclusion

This study demonstrates the potential of Pr 1:3 as a viable and sustainable deep eutectic solvent for metal electrodeaching and electrodeposition, using silver electrodeaching electrodeposition as a model electrochemical reaction. Neither the 7 month-long term storage at 50 °C, nor the 476 h-long reasonable electrochemical operation (60°C and 2 mA cm⁻²), demonstrated the formation of toxic degradation compounds. The only aging effect observed was due to evaporation, which was enhanced due to the nitrogen gas flow used to maintain a dry atmosphere. The Propylene Glycol: Choline Chloride constituting the Propeline 1:3 shifted from 1:3 (molar ratio) to ≈1:2, leading to a slight improvement and a slight diminishing of the DES viscosity and conductivity respectively. This suggests that the DES properties could be restored through simple PG replenishment.

However, moving toward harsher conditions is more problematic. The forced electrochemical degradation of Pr 1:3 carried out by direct solvent electrolysis at 1 mA cm⁻² over 168 h led to the formation of more than 20 different compounds in the gas and liquid phase. Notably, several chlorinated compounds were identified, including 2-(chloromethyl)-2,4-dimethyl-1,3-dioxolane, which was the most concentrated compound in the gas phase after PG, as well as chloromethane, chloroethanol and chloroform. These chlorinated compounds pose toxicity and environmental concerns and should be taken into account for future safety measurements. Extending the degradation up to 20 days led to changes in density, viscosity, and electrical conductivity, which could negatively impact the applicability of Pr 1:3 for industrial applications, particularly in processes that rely on efficient fluid handling and ion transport. Worse, the depletion of Cl⁻ ions in the system led to the total inhibition of the solvent metal electrodeaching ability.

In a nutshell, long-term operation of propeline 1:3 does not degrade the solvent and provides a positive insight on this solvent durability as it could be used without additional chemical treatment for several electrolysis reaction batches. Yet, operation malfunction which would electrolyze the solvent does produce toxic degradation compounds and decrease the chlorure
<figure><img src="image_28.png" /></figure>

concentration, and—at term—nullify the solvent interest in ion-
ometallurgy. Further work would however be required to under-
stand more clearly how the produced degradation compounds
could influence the electrochemical reaction efficiency, and how
the solvent could be treated to remove them. Working at low
current density (2 mA cm⁻² in this work), is up to date essential
to prevent such degradation, but not particularly problematic
due to the limited volume of precious-metals to be electro-
wonned through DES EL-ED processes. However, further process
engineering input will be considered in future work to improve
hydrodynamic in the electrochemical cells involving DESs, and
therefore increase this current density limitation without
degradation.

**4. Experimental Section**

Reagents and DES Preparation

Choline chloride (≥98%), propylene glycol (≥99.5%) and silver chlo-
ride (99%) were all purchased from Sigma-Aldrich.

Silver foils (Premion 50 × 50, 0.25 mm thickness, 99.95%) were bought from Thermo Scientific. Standard Pr 1:3 was synthesized by mixing choline chloride, previously dried at 80 °C under vacuum for 2 h, and propylene glycol in a 1:3 molar ratio. The reaction mixture was heated at 80 °C and stirred with a magnetic bar for 2 h until a homogeneous solution was obtained. All solutions were stored at 50 °C in closed glass bottles.

A dry sample of Pr 1:3 was synthesized following the same procedure but inside a glove box under argon atmosphere. Choline chloride was first dried overnight at 80 °C. For deoxygenation, argon gas was bubbled through the DES for 30 min.

For the EL-ED process of silver, 70 mM AgCl solution of Pr 1:3 were prepared for each experiment by stirring the appropriate amount of silver chloride at 80 °C for 1 h.

Characterization Techniques: Gas Chromatography—Mass Spectrometry (GC-MS)

Liquid samples were prepared by dissolving 25 µL of Pr 1:3 in 1 mL of ethanol. For Propeline solutions containing AgCl, the samples were filtered through 0.2 µm syringe PTFE filters prior to analysis in order to remove any precipitate. Degradation compounds in the liquid phase were identified using an Agilent 5977E Series GC-MSD with electron ionization (EI). The oven temperature program was as follows: the temperature was maintained at 50 °C for 5 min, then increased at a rate of 15 °C min⁻¹ until 260 °C, where it was held for 1 min, resulting in a total run time of 20 min. The analysis was conducted with a split flow of 2.33 L min⁻¹, a carrier gas flow rate of 0.7 mL min⁻¹ and a solvent delay of 5 min. The column used was an Agilent 19091N-233HP-INNOWax (30 m × 250 µm × 0.5 µm). Masses between 16 and 800 m/z were scanned, and compound identification was performed using Agilent MassHunter Qualitative Analysis B.06.00 software and the NIST/EPA/NIH mass spectral library (2011 version).

Volatile organic compounds (VOCs) release in gas phase were inves-
tigated by samplings on adsorbent tubes and subsequent analysis by
thermal desorption—gas chromatography—mass spectrometry/
flame ionization Detector. Sampling was carried out by sweeping
the electrolysis cell with a N₂ flow (40 sccm) and placing the

adsorbent tube (Markes Air Toxics Analyzer tubes) at the cell outlet,
maintained at ambient temperature. The analysis equipment used
consisted of a Thermal Desorber Ultra-xr series 2 from Markes and
a gas chromatograph 5975C VL MSD from Agilent Technologies.
Both MS and FID were used in parallel, the MS allowing the identifi-
cation of collected VOC and FID the quantification analysis. The FID
calibration was performed using toluene and the quantification of
individual VOC and total VOC (obtained from the sum of all chro-
matographic peaks) is given in toluene equivalent amounts.
Before use, adsorbent tubes were previously thermally conditioned
under inert gas and their cleanliness was controlled by TD-GC-MS/
FID analysis. Thermal desorption was performed using the following
conditions: tube desorption at 320 °C for 15 min at 50 mL min⁻¹; cold
trap at 30 °C without split; trap heating at 40 °C s⁻¹ up to 300 °C
(5 min) with a split ratio of 1:25 and a transfer line maintained at
210 °C. The analysis conditions for the GC were as follows: With
He as carrier gas at 1 mL min⁻¹ and an Agilent column DB-624-Ul
(30 m × 0.25 mm × 1.4 µm), a temperature program of 40 °C
(5 min), 5 °C min⁻¹ up to 250 °C (30 min) was applied.

Characterization Techniques: NMR

All NMR measurements were recorded on a Bruker Neo 400 MHz spectrometer equipped with a BBFO smartprobe. Quantitative ¹H NMR (qNMR) was carried out using anhydrous sodium acetate (NaAcetate, ≥99.0%, Sigma-Aldrich) as an internal standard. Before quantification, T₁ relaxation times of all relevant proton signals of Pr 1:3 and sodium acetate were measured using an inversion recovery experiment. The longest T₁ value was used to set the relaxation delay (D₁ = 5 x T₁) to ensure quantitative accuracy.

For the calibration curve, 1 mL of a 0.1 M sodium acetate stock solu-
tion in D₂O (99.9%, Sigma-Aldrich) was mixed with increasing vol-
umes of Pr 1:3 in five separate vials. The ¹H NMR spectra of these
calibration samples were acquired, and the areas of characteristic
peaks for choline chloride (3.22 ppm, singlet, N⁺-CH₃) and propylene
glycol (1.17 ppm, doublet, CH-CH₃) were measured relative to the
internal standard (1.93 ppm, singlet, CH₃-COO⁻). Calibration curves
were obtained by plotting the ratio of the peak areas (A<sub>Analyte</sub>/
A<sub>NaAcetate</sub>) against the concentration ratios. For sample analysis,
1000 µL of the 0.1 M sodium acetate solution in D₂O was mixed with
40 µL of degraded Pr 1:3 and transferred to an NMR tube for analysis.
Spectra were acquired using the standard Bruker pulse program zg
with 32 scans. Processing was performed in Bruker TopSpin 4.3.0,
with baseline correction and manual phase adjustment.

Characterization Techniques: SEM

SEM imaging of silver electrodes was performed using a Nova
NanoSEM 200 (FEI Company) in low vacuum mode with a Helix detec-
tor for secondary electrons. Energy dispersive X-ray (EDX) analysis
was conducted using EDAX Genesis v5.10 software (Ametek) to iden-
tify elemental composition.

Characterization Techniques: Viscosimeter/Densimeter

Dynamic viscosity measurements were performed using a LOVIS
2000 M/ME falling ball viscometer (Anton Paar) equipped with
a 1.8 mm capillary (Anton Paar) and gold beads of 1.5 mm diameter
and 7.94 g cm⁻³ density. Measurements were carried out between
30 and 100 °C. Density measurements were conducted using
a DMA 4100 M densimeter (Anton Paar) over the same tempera-
ture range.
<figure><img src="image_29.png" /></figure>

## Characterization Techniques: Karl-Fischer Titration

The water content of the DES was determined by KarlFischer coulometric titration using a 899 Coulometer (Metrohm) and Hydranal Coulomat AG (Honeywell) titration solution. Approximately 0.1 mL of each sample was analyzed.

## Electrochemical Experiments

Two different three-electrode electrochemical cells were used for this study (Figure S1, Supporting Information). The electrochemical cell 1 consists in a reservoir capable of holding ca. 35 mL of electrolyte with two glassy carbon disk electrodes of 15.9 cm². The electrodes were arranged in a parallel configuration separated by a distance of 2 cm. A double-walled glass chamber allows for temperature control by circulating silicon oil, maintaining the system at 60 °C. Additionally, two lateral connections enable the flow of nitrogen gas at ca. 40 sccm, establishing an inert and dry atmosphere. No stirring was applied in this configuration. The reference electrode consisted of a silver wire insulated in a glass tube filled with a 10 mM AgCl solution of Pr 1:3, separated from the main solution by a ceramic frit. All potentials were reported versus Ag/AgCl.

The electrochemical cell 2 consists of a reservoir capable of holding ca. 45 mL of electrolyte. Both the anode and cathode were silver plates of 15 × 20 mm overall surface and 0.25 mm thickness separated about 1–2 cm apart. The cell was double-walled to allow temperature control at 60 °C using silicon oil. An inert and dry atmosphere was established by perforating two septa to insert connections for nitrogen gas flow (40 sccm). A cross-shaped magnetic stir bar rotating at 600 rpm was used for agitation. The reference electrode was identical to the one used in the electrochemical cell 1. For the experiments described in the Electrochemical characterization section, glassy carbon rod electrodes (5 mm diameter × 150 mm length; Neyco), spaced 2 cm apart, were used in electrochemical cell 2.

All CV and CP measurements were performed using a Biologic VSP potentiostat controlled with EC-Lab software v11.61. Experimental parameters such as scan rate, number of cycles, current density and electrolysis duration are provided in the main text.

## Acknowledgements

This work was supported by a grant from the French government managed by the National Research Agency (ANR) under the France 2030 program, reference number ANR-22PERE-0003.

## Conflict of Interest

The authors declare no conflict of interest.

## Data Availability Statement

The data that support the findings of this study are available on request from the corresponding author. The data are not publicly available due to privacy or ethical restrictions.

**Keywords:** degradation compounds · deep eutectic solvents · durability · electrodeposition · electroleaching · silver

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