Electrolyte-Driven Oxidant Generation on Ti/IrO₂–SnO₂–Sb₂O₅ Electrodes for the Efficient Removal of Alachlor and Isoproturon from Water

Abstract

In this study, anodic oxidation (AO) was evaluated using Ti/IrO₂–SnO₂–Sb₂O₅ electrodes in chloride, sulfate, and mixed electrolytes, along with electro-Fenton (EF) and photoelectro-Fenton (PEF) at pH 3.0, for the degradation of alachlor and isoproturon, each 50 mg L⁻¹. Active chlorine species were monitored using UV–Vis, while the removal of both herbicides was quantified using High Performance Liquid Chromatography (HPLC), along with the reduction in Total Organic Carbon (TOC), mineralization current efficiency (MCE), and specific energy per TOC removed (EC_TOC). The results show that electrolyte composition influences AO more than current density. In a chloride medium, isoproturon was eliminated within minutes, whereas alachlor required mixed electrolytes of Cl⁻/SO₄²⁻, allowing simultaneous combination of HClO/ClO⁻, ^●OH, and S₂O₈²⁻/SO₄^●−, or coupling with EF. An optimal current density of ~30 mA cm⁻² limited voltage rise and radical scavenging. EF introduced measurable mineralization (15% TOC), whereas PEF achieved rapid alachlor reduction and TOC reductions of up to 76% at low Fe²⁺. Overall, sequential AO followed by PEF maximized mineralization per unit of energy, and the mixed electrolytes provided a controllable pathway to scale up oxidant speciation generation.

1. Introduction

Herbicides such as isoproturon (ISP) and alachlor (ALA) are persistent micropollutants in industrial effluents [1]. They are frequently detected in surface and groundwater, where they exert adverse ecological effects and raise human health concerns, including endocrine-disrupting activity reported for ISP [2] and a carcinogenic effect reported in animal studies for ALA [3]. Their persistence, toxicity, and resistance to conventional treatment motivate the development of advanced oxidation technologies capable of achieving mineralization rather than transformation [4]. Beyond their occurrence in agricultural run-off, ISP and ALA are increasingly reported in mixed industrial and municipal effluents, often coexisting with chloride, carbonate alkalinity, and natural organic matter at µg to mg L⁻¹ levels. For ALA, surveys document 0.06–31.9 µg L⁻¹ in surface waters and 0.05 to 4.85 µg L⁻¹ in groundwaters, with drinking water wells in the USA occasionally reaching <0.1–16.6 µg L⁻¹. ISP is typically found at ng L⁻¹-µg L⁻¹ in rivers, with maxima up to 38 µg L⁻¹ reported; industrial effluents from manufactured pesticide can be orders of magnitude higher at 10–100 mg L⁻¹ [5,6,7]. These complex matrices can suppress oxidative pathways via radical scavenging and light attenuation, while promoting the formation of halogenated by-products under chlorinating conditions. Consequently, technologies that (i) tolerate variable ionic strength, (ii) steer oxidant speciation toward high-reactivity routes, and (iii) quantify mineralization rather than sole decay are required to mitigate cumulative ecological and human health pressure along river basins and groundwater bodies.

Electrochemical advanced oxidation processes (EAOPs), including anodic oxidation (AO), electro-Fenton (EF), and photoelectro-Fenton (PEF), are especially attractive because they generate oxidants in situ (Equation (1)) and can be tuned via electrode material, electrolyte support combination, and current density [8,9]. In chloride-bearing waters, electro-chlorination at the anode produces active chlorine species whose distribution is pH-dependent (Equations (2)–(4)) [10], while in persulfate media (S₂O₈²⁻), sulfate radicals (SO₄^−●) can be formed (Equations (5)–(7)) [11]. In EF and PEF, bulk hydroxyl radicals (^●OH) arise from the Fenton reaction (Equation (8)) and from photochemical cycling of Fe (III)-Fe (II) (Equation (10)). Together, these pathways offer complementary oxidizing power to abate recalcitrant organics and lower total organic carbon (TOC) [12,13,14]. From a process-intensification standpoint, EAOPs are attractive for scale-up because they do not require external chemical oxidant species. Reactor performance can be tuned by electrode selection (DSA and BDD), support electrolyte (Cl-/SO₄²⁻) mixtures, and hydrodynamic control. These knobs allow translation from bench-stirred cells to flow reactors (e.g., FM01-LC) while preserving current distributions and mass transfer that are essential to maintaining kinetic energy efficiency at a larger scale.

M (electrode surface) + H₂O → M(^●OH) + H⁺ + e⁻

(1)

2 Cl⁻ → Cl₂ + 2e⁻

(2)

Cl₂ + H₂O ⇄ HClO + Cl⁺ + H+

(3)

HClO ⇄ ClO⁻ + H⁺     pKa = 7.4

(4)

S₂O₈²⁻ + e⁻ → SO₄^●− + SO₄²⁻

(5)

SO₄^●− + e⁻ → SO₄²¯

(6)

2 SO₄²⁻ → S₂O₈²¯ + 2 e¯

(7)

Positioned within Sustainable Development Goal 6 (Clean water and Sanitation), the decisive metric for advanced treatment is not only the rate of compound removal but the energy-normalized mineralization achieved through by-product liability. Accordingly, MCE (mineralization current efficiency) and EC_TOC (specific energy per TOC removed) under controlled hydrodynamics provide the quantitative basis to compare operational windows and to judge cost and carbon removal efficiency. An electrolyte guide strategy that broadens the oxidant species while maintaining acidic operation is central to balancing performance, by-product control, and energy demand.

Despite substantial progress in EAOPs, several gaps remain. Many studies emphasize rapid compound decay or decolorization but report limited mineralization to alachlor and isoproturon. Systematic mapping of oxidant speciation as a function of pH and electrolytes is often incomplete, and energy metrics, such as mineralization current efficiency (MCE) and the specific energy per TOC removed (EC_TOC), are not consistently quantified under matched hydrodynamic and electrical conditions. Moreover, the chloride-sulfate “co-electrolyte” strategy, potentially enabling simultaneous HClO/ClO⁻, ^●OH, and S₂O₈^●− generation, has not been evaluated alongside mineralization and energy consumption of ALA and ISP.

The present work addresses these shortcomings through an integrated experimental framework that first characterizes active chlorine electrogeneration and speciation by UV–Vis (λ = 292 nm for ClO⁻ and 235 nm for HClO). Next, it determines anodic oxidation on Ti/IrO₂-SnO₂-Sb₂O₅ (DSA) across sulfate, chloride, and mixed electrolytes at controlled current densities; and after, it quantifies the incremental contributions of electro-Fenton and photoelectro-Fenton at pH 3.0 to both degradation kinetics and mineralization via MCE (mineralization current efficiency) and EC_TOC (specific energy consumption per unit mass of TOC removed). This work considers the relevant radical production pathways, including anodic chlorine evolution (Equations (2)–(4)) [15], anodic persulfate formation in sulfate media (Equations (5)–(7)) [16], the Fenton reaction (Equation (8)), and the chlorine-assisted Fenton step (Equation (9)) [17]; while under irradiation, photoreduction/photolysis sustains Fe (III)-Fe (II) cycling (Equation (10)) [18]. By coupling rigorous analytics, HPLC for target compounds, ion chromatography for inorganic oxidants, and TOC for mineralization, with energy accounting and process testing, it delineates an operating window in which AO delivers rapid primary abatement, whereas EF and PEF provide deeper carbon removal, clarifying how the support electrolyte influences the formation of oxidizing species and, ultimately, overall treatment efficiency.

Finally, to bridge laboratory findings with practice, the operating window delineated here targets chloride- and sulfate-bearing effluents typical of industrial facilities and is compatible with pre-discharge treatment. The emphasis on maintaining acid conditions, mixed Cl⁻/SO₄²⁻ electrolytes, and moderate current density provides transferable information for treating real wastewaters in recirculating flow cells, where matrix effects can be mitigated while preserving high oxidant availability.

Fe²⁺ + H₂O₂ → Fe³⁺ + ^●OH + OH⁻

(8)

Fe²⁺ + HClO → Fe³⁺ + ^●OH + OH⁻

(9)

[Fe (OH)]²⁺ + hv → Fe²⁺ + ^●OH

(10)

Recent EAOP studies on herbicides show that anodic oxidation (AO) typically achieves rapid parent compound decay yet delivers limited mineralization unless operated for prolonged times or at high specific energy. In contrast, electro-Fenton (EF) and photoelectro-Fenton (PEF) reliably increase mineralization because homogeneous ^●OH is produced via an Fe²⁺ and H₂O₂ reaction and photoregeneration of Fe²⁺, respectively. Under an acidic pH and low Fe²⁺, PEF commonly outperforms WF and AO in total organic carbon (TOC) abatement at a comparable or lower energy per unit of mineralization. Within this landscape, an underexplored yet decisive lever is electrolyte engineering; mixed chloride and sulfate can tune oxidant speciation while preserving the acidic window required for EF and PEF, thereby improving kinetics and selectivity by merely increasing current density [19,20].

To ensure scale awareness, this work adopts energy-normalized performance metrics as core figures of merit. It reports current efficiency (MCE) as specific energy per TOC removal. Using these metrics alongside removal and TOC decay allows us to identify an energy-rational operating window and to quantify when electrolyte tailoring outperforms simply raising current density. Accordingly, the contribution of this manuscript is a process advancement, electrolyte engineering combined with AO, EF, and PEF sequencing under an EF-compatible acidic pH, rather than a material novelty.

2. Experimental Section

2.1. Chemicals

Analytical Analytical-grade sodium chloride, sodium sulfate, and ferrous sulfate heptahydrate were purchased from Sigma-Aldrich (St. Louis, MO, USA), while reagent-grade sodium hypochlorite was supplied by Merck (Darmstadt, Germany). Isoproturon (C₁₂H₁₈N₂O, 206.3 g mol⁻¹, solubility 70 mg L⁻¹ at 20 °C) and alachlor (C₁₄H₂₀ClNO₂, 269.76 g mol⁻¹ at 25 °C) were obtained from Aragonesas Agro S.A. (Madrid, Spain); their physicochemical specifications are summarized in Table 1.

Sodium hypochlorite solution (0.8 M) served as the hypochlorite standard, constructing the UV–Vis absorption calibration curve of active chlorine species (HClO/ClO⁻). All test solutions were formulated with deionized water whose conductivity was adjusted via the controlled addition of supporting electrolytes (NaCl and Na₂SO₄). Three representative matrices were prepared: (i) 30 mM NaCl, giving a conductivity of 3.94 mS cm⁻¹, (ii) 30 mM Na₂SO₄, yielding 3.54 mS cm⁻¹, and (iii) an equimolar mixture of 15 mM NaCl + 15 mM Na₂SO₄, supplying 3.75 mS cm⁻¹. These compositions were selected to emulate chloride-rich, sulfate-rich, and mixed ionic environments commonly encountered in agro-industrial effluents, while keeping ionic strength comparable across experiments to ensure reproducible mass transfer and ohmic conditions during electrochemical treatments.

2.2. Analytical Procedures

The concentration of HClO/CIO⁻ generated during NaCl electrolysis was quantified by UV–Vis spectrophotometry (Cintra 1010). The calibration curves for active chlorine species were established by recording the absorbance at 292 nm for ClO⁻ and 235 nm for HClO.

For herbicide kinetics, 20 μL aliquots, pre-filtered through 0.45 μm nylon membranes, were injected into an Agilent 1100 HPLC/DAD equipped with a Luna C18 column (150 × 3 mm, 5 μm, 100 Å) operated isocratically at 0.5 mL min⁻¹ with 60% acetonitrile and 40% formic acid (25 mM). The detector wavelengths were 240 nm for isoproturon (time retention = 2.8 min) and 220 nm for alachlor (time retention = 6.2 min), affording calibration curves with R² ≥ 0.999. Total organic carbon (TOC) was determined on a Shimadzu TOC-VCSN analyzer to evaluate mineralization efficiency. All measurements were duplicated, with the instrument’s precision better than $\pm$ 3%. The resulting TOC data obtained were used to calculate overall mineralization current efficiency (MCE), according to Equation (11) [21].

$$\% MCE={\displaystyle \frac{n F V \Delta (TOC{)}_{exp}}{4.32 \times {10}^{7}m I t}}\times 100$$

(11)

where n = 58 e⁻ is the electrons for isoproturon mineralization and n = 68 e⁻ is the electrons for alachlor, F is the Faraday constant (96,457 C mol⁻¹), V is the solution volume (L), $\Delta$(TOC)_exp is the experimental TOC decay (mg C L⁻¹), 4.32 × 10⁷ is a conversion factor to homogenize units (3600 s h⁻¹ × 12,000 mg C mol⁻¹), m = 12 is the number of carbon atoms of isoproturon and m = 14 is the number of carbon atoms of alachlor, I is the applied current (A), and t is the electrolysis time (h).

Assuming isoproturon undergoes complete conversion to its final products, the stoichiometrically balanced equation (Equation (12)) is:

C₁₂H₁₈N_2O + 23H₂O → 12CO₂ + 2NH₄⁺+56H⁺ + 58e⁻

(12)

For alachlor, presuming it is entirely oxidized to its final products, the reaction can be expressed by Equation (13).

C₁₄H₂₀ClNO₂ + 26H₂O → 14CO₂ + NH₄⁺ + Cl⁻ + 68H⁺ + 68e⁻

(13)

The specific energy consumption per unit mass of TOC removed (EC_TOC) at a given time (h) for each experiment was determined with Equation (1) [22].

$${EC}_{TOC}({\left(kWh·gTOC\right)}^{-1})={\displaystyle \frac{E_{cell }I t}{V \Delta (TOC{)}_{exp}}}$$

(14)

where E_cell is the average cell voltage (V).

2.3. Reactor Design

Degradation tests were performed in a batch cell, magnetically stirred (160 rpm) to ensure homogeneous mass transfer. A 500 mL working volume was operated at room temperature. The electrochemical cell employed a titanium substrate coated with IrO₂-SnO₂-Sb₂O₅ (Ti/IrO₂-SnO₂-Sb₂O₅) as the anode. It exhibited a geometric area of 3 cm² and was positioned 1 cm opposite a BDD cathode of identical area. The Ti/IrO₂-SnO₂-Sb₂O₅ was manufactured by an adapted Pechini method and was reported in detail by our group previously [23]. Constant current densities of 15, 30, or 60 mA cm⁻² were supplied by a DC source (RS PRO IPS 3610D). For the electro-Fenton process, Fe²⁺ (0.05–1 mM) was dosed at the start and compressed air bubbled to generate H₂O₂ in situ, while photoelectro-Fenton trials were run using a solar simulator (Suntest, XLS + Heraeus Noblelight GmbH, Hanau, Germany) equipped with a 2.2 kW xenon arc lamp and a special glass filter (daylight) that cuts off UV irradiation at 290 nm. The radiation from the lamp was measured using a PMA2100 pyranometer (Solar Light, Glenside, PA, USA).

All experimental runs of AO, EF, and PEF were performed in duplicate to ensure repeatability across conditions, electrolytes, current density, and Fe²⁺ concentration. The operating window identified as optimal was repeated in triplicate. Unless stated otherwise, results are reported as mean ± standard deviation (SD). Error bars in time course plots denote ± SD.

Quantification of ISP and ALA was based on external calibration curves by linear least squares regression, yielding R² ≥ 0.999. Each condition included duplicate runs and for energy figures (EC_TOC), uncertainty was calculated from the averaged cell voltage, applied current, electrolysis time, treated volume, and TOC change; tabulated values are reported as mean ± SD. The model fit for the pseudo first order kinetics was evaluated by ordinary least squares; parameter estimates are reported with their standard errors (SEs) and coefficient of determination (R²).

The electrolysis target solutions consisted of ISO and ALA, each prepared at an initial concentration of 50 mg L⁻¹. Actual starting concentrations for each essay are reported at t = 0 in the corresponding figures.

2.4. Operating Conditions

Before each trial, both the anode and cathode were electrochemically activated by polarization in 0.1 M H₂SO₄ at 30 mA cm⁻² for 60 min to remove surface impurities and stabilize the potential response. Electrolyte solutions, 30 mM NaCl, 30 mM Na₂SO₄, or an equimolar 15 mM NaCl + 15 mM Na₂SO₄, were prepared with deionized water, adjusted to pH 3.0 with H₂SO₄, and transferred to the batch reactor. The galvanostatic program was initiated, and time-zero samples were taken. Aliquots (2–4 mL) were withdrawn periodically up to 240 min, filtered through 0.45 μm nylon membranes, and analyzed. HPLC for isoproturon and alachlor, pH, conductivity, and cell voltage was logged continuously. For electro-Fenton and photoelectro-Fenton assays, Fe SO₄.7H₂O was dosed to the desired Fe²⁺ level (0.05–1 mM) just before electrolysis and compressed air was sparged at 100 mL min⁻¹. In the photo-assisted experiments, the trials were conducted using the previously described solar simulator.

3. Results and Discussion

3.1. UV–Vis Analysis of a NaClO Solution

A total of 0.8 M of sodium hypochlorite was serially diluted to generate free-chlorine standards, enabling identification of the characteristic UV absorption bands of hypochlorous ion (CIO⁻). In aqueous media, NaCl hydrolyzes to produce HClO according to Equation (15) [24]:

NaCl + H₂O → Na⁺ + HClO + OH⁻

(15)

However, the concentration of the weak acid HClO and its conjugate base ClO⁻ are subjected to the pH value of the solution by Equation (16) with pKa = 7.4 [25]

HClO ⇄ ClO⁻ + H⁺

(16)

Figure 1 depicts the profile of active chlorine in a 1.0 mM NaClO solution at 25 °C. The diagram indicates that hypochlorous acid (HClO) is the principal form between pH 2.0 and 8.0, and hypochlorite ion (ClO⁻) prevails above pH 8.0.

Figure 2 shows the UV–Vis spectra of the 1.0 mM NaCl solution recorded at various pH values. At pH 3.0 a band was well-defined at 235 nm, attributable to HClO. In alkaline medium (pH 10), the spectrum showed a single intense peak at around 292 nm, confirming that ClO⁻ is the predominant chlorine species under these conditions.

3.2. Electrogeneration of Hypochlorite Ion on Ti/IrO₂-SnO₂-Sb₂O₅

NaCl solutions (30 mM) pre-adjusted to initial pH values of 3.0, 5.0, 7.0, and 10 were electrolyzed in a batch cell equipped with a Ti/IrO₂-SnO₂-Sb₂O₅ anode for 60 min. The corresponding UV–Vis spectra (Figure 3) show the hypochlorite band at 292 nm; the signal was maximal at pH 10 and minimal at pH 3.0. This trend is consistent with the pH-dependent HClO/ClO⁻ acid–base equilibrium (Figure 2), whereby alkaline conditions favor ClO⁻. The hypochlorite concentration exhibited a clear dependence on the initial pH; between pH 3.0 and 7.0, [ClO⁻] rose approximately linearly. At pH 8.0, a pronounced jump was observed, followed by a maximum value at pH 10. The lack of the hypochlorous acid band at 235 nm for solutions initiated at pH < 7.0 is consistent with an electrolysis-induced pH drift to >8.0, which shifts the active chlorine speciation toward ClO⁻.

3.3. Effect of NaCl Concentration

Figure 4 shows the UV–Vis spectra of NaCl solutions (30 mM, 15 mM, and 10 mM) electrolyzed for 60 min using the Ti/IrO₂-SnO₂-Sb₂O₅ anode. An absorption band was determined at 235 nm, attributable to electrogenerated hypochlorous acid. The band intensity increased with NaCl concentration, indicating that under identical operating conditions, higher chloride levels promote greater hypochlorous acid production.

3.4. Anodic Oxidation of Isoproturon and Alachlor

3.4.1. Effect of Current Density

Anodic oxidation experiments were conducted at current densities of 15, 30, and 60 mA cm⁻² in different supporting electrolytes: NaCl (30 mM), Na₂SO₄ (30 mM), and their equimolar mixture (15 mM NaCl + 15 mM Na₂SO₄). These conditions are widely used in EAOP studies with DSA-type anodes [26].

Galvanostatic trials with the Ti/IrO₂-SnO₂-Sb₂O₅ anode revealed a clear dependence of herbicide abatement on the applied current density (j). Figure 5a shows that in sulfate medium (30 mM Na₂SO₄, pH 3.0), raising j from 15 to 30 mA cm⁻² increased isoproturon removal from 24% to 35% and alachlor (ALA) removal from 36% to 58% after 240 min (Figure 5b). However, a further step to 60 mA cm⁻² afforded only marginal gains (ISP 37%, ALA 64%) while trebling the cell voltage (4.4 to 11.2 V). These data point to 30 mA cm⁻² as the kinetic-energetic optimum for sulfate-based AO, beyond which mass transfer limits and parasitic O₂ evolution reduce efficiency, which diverts charge away from organic oxidation and elevates cell voltage; this behavior is well documented for AO systems [20,27].

3.4.2. Effect of Electrolyte Support

Substituting chloride for sulfate profoundly modified the oxidative environment during anodic oxidation (Figure 6). At a constant current density of 30 mA cm⁻², the Ti/IrO₂–SnO₂–Sb₂O₅ anode exhibited a markedly different behavior depending on the supporting electrolyte. When NaCl was used, the system generated active chlorine species in situ, substantially accelerating pollutant removal relative to sulfate medium. In this matrix, isoproturon (ISP) degradation was nearly complete within 15 min, while alachlor (ALA), the more recalcitrant chloroacetanilide, achieved about 65% conversion after 240 min. This strong performance evidences the dominant role of indirect oxidation via Cl₂/HClO/ClO⁻ and highlights the intrinsic catalytic capability of the DSA surface toward chlorine evolution reactions [28,29].

However, the improved kinetics in chloride medium were accompanied by a significant pH drift (from 3.0 up to nearly 7.0) and a parallel increase in cell voltage compared with sulfate electrolyte, implying greater energy demand and the possible accumulation of chlorinated intermediates under quasi-neutral conditions [30]. This pH rise originates from the cathodic hydrogen evolution reaction in the undivided cell and from the anodic oxidation of chloride (Equation (17)), which continuously generates hydroxide ions that diffuse into the bulk solution. Although the hydrolysis of molecular chlorine (Equation (18)) releases protons, they are insufficient to neutralize the cathodic OH⁻ flux, producing a net alkalinization that shifts the active chlorine equilibrium toward ClO⁻, the less reactive but more stable form [31]. Consequently, the process becomes progressively governed by slower chlorine-mediated pathways rather than direct anodic ^●OH attack.

2 Cl⁻ → Cl₂ + 2 e⁻

(17)

Cl₂ + H₂O ⇄ HClO + Cl⁻ + H⁺

(18)

Although the formation of oxychlorinated by-products was not quantified in the present study, previous research using similar DSA-type anodes and chloride-containing electrolytes has shown that their generation is strongly suppressed under acidic conditions and moderate current densities. In our system, these conditions were maintained throughout the experiments, which minimizes the accumulation of ClO₃⁻ and ClO₄⁻ and favors selective oxidation via HClO/ClO⁻ and ^•OH rather than over-oxidation of chloride. This operational window aligns with reports demonstrating negligible formation of hazardous oxychlorine species under comparable EAOP conditions, thus supporting the environmental safety of the proposed process.

When a mixed electrolyte containing 15 mM Cl⁻ + 15 mM SO₄²⁻ was employed, the reaction environment became more balanced. Under these conditions, ISP removal remained rapid, approaching complete disappearance within two hours, while ALA degradation rose to about 80% after 240 min, outperforming both single-anion systems. Importantly, the pH remained close to 3.0 throughout electrolysis, preserving an acidic window compatible with subsequent electro-Fenton (EF) operation and minimizing the build-up of free hypochlorite. The sustained acidity suggests that sulfate acts as a pH-buffering component, moderating cathodic OH⁻ accumulation while still enabling moderate Cl⁻ mediated oxidation. This equilibrium allows a complementary interplay between oxidants: (i) surface-bound ^●OH generated at the DSA interface, (ii) electrogenerated HClO/ClO⁻ with redox potential E° = 1.49 V vs. SHE, and (iii) anodic persulfate S₂O₈²⁻ (E° = 2.01 V). Similar enhancements with Cl⁻ and SO⁻²₄ have been reported for electro-oxidation of herbicides and other micropollutants, where electrolyte tailoring improves kinetics relative to single salts (mixed pathways via active chlorine and sulfate oxidants) [29,32]. Their concurrent presence broadens the oxidative spectrum, enhancing degradation without resorting to excessive current densities or generating large potential drops.

Mechanistically, this synergy can be rationalized as a dual radical regime. In the chloride pathway, fast reactions of free or combined chloride and chlorine radicals with aromatic moieties drive rapid parent compound decay, comparable to or only modestly below ^●OH for selected substrates, explaining the accelerated removal seen in these results [33]. In parallel, sulfate chemistry near the anode enables the transient formation of SO₄^●− and S₂O₈²⁻ (Equations (5)–(7)), which are effective toward ring-opened and aliphatic intermediates and support partial mineralization; recent reviews and kinetics studies document electro-oxidation activation of sulfate to SO₄^•−/S₂O₈²⁻ on DSA and its contribution to deeper oxidant steps [34,35]. Such findings support the concept of electrolyte engineering as a cost-effective intensification strategy, in which the oxidative selectivity can be steered through controlled anion composition instead of higher current input.

Overall, the results demonstrate that electrolyte composition exerts a more decisive influence on anodic oxidation efficiency than the applied current density itself. Mixed Cl⁻/SO₄²⁻ electrolytes at 30 mA cm⁻² deliver both rapid and energy-rational ISP abatement and significant ALA degradation while maintaining acidic operation beneficial for subsequent EF and PEF stages. In contrast, increasing current density beyond 30 mA cm⁻² would only raise the extent of gas evolution (Cl₂ and O₂), causing bubble accumulation on the electrode, higher ohmic overpotential, and lower instantaneous current efficiency [36]. Thus, the mixed electrolyte configuration defines an optimal trade-off between kinetic performance, energy economy, and compatibility with downstream electro-Fenton mineralization.

3.5. Electro-Fenton and Photoelectro-Fenton Performance

In light of the electro-oxidation results, electro-Fenton and photoelectro-Fenton were investigated using a Ti/IrO₂-SnO₂-Sb₂O₅ anode and BDD cathode. Experiments were conducted at current densities of 15, 30, and 60 mA cm⁻² in different supporting electrolytes: NaCl (30 mM), Na₂SO₄ (30 mM), and their equimolar mixture (15 mM NaCl + 15 mM Na₂SO₄). The aim was to evaluate the performance of distinct EAOPs in the degradation of alachlor and isoproturon.

3.5.1. Effect of Fe²⁺ in EF Process

The incorporation of Fe²⁺ into the system profoundly altered the oxidative regime of anodic oxidation, marking the transition to the electro-Fenton (EF) domain (Figure 7). In acidic chloride electrolyte (30 mM NaCl, pH 3.0), introducing 0.15 mM Fe²⁺ created a complex interplay between direct anodic and homogeneous radical pathways. As shown in Figure 7a, isoproturon degradation remained rapid, complete within roughly 20 min, but slightly slower than in pure anodic oxidation (AO), where full abatement occurred in under 15 min. This modest deceleration does not reflect a loss of efficiency but a redistribution of oxidant pathways. In AO, oxidation proceeds mainly through electrogenerated chlorine species (Cl₂/HClO/ClO⁻) that react directly with target molecules near the anode. Once Fe²⁺ is present, a fraction of HClO is consumed in the Fenton-like step (Equation (19)), producing hydroxyl radicals in the bulk solution and thereby shifting the process from a surface-controlled to a solution-mediated oxidation regime rather than reducing overall efficiency; supporting results and kinetics in chlorine systems are reported in [37,38].

Fe²⁺ + HClO → Fe³⁺ + ^●OH + Cl⁻

(19)

This mechanistic redistribution expands the reactive zone and diversifies the oxidant spectrum. Chloride ions act simultaneously as promoters and scavengers: ^●OH may react with Cl⁻ to generate chlorine radicals (Cl^● and Cl₂^●−), less oxidative but longer-lived species that propagate secondary oxidation in the bulk [39,40]. Consequently, EF exhibits a steadier, more sustained degradation pattern compared with the rapid, surface-dominated AO process.

For alachlor (Figure 7b), EF and AO yielded comparable conversions (~60–65% after 240 min). As a chloroacetanilide, ALA resists primary oxidation and requires either more energetic radicals or prolonged bulk oxidation. The similar efficiencies observed here confirm that the overall rate is governed more by the oxidant distribution than by the Fe²⁺ concentration itself. In chloride-rich matrices, ^●OH radicals generated via Fe²⁺/HClO interactions are partially quenched by Cl⁻, while anodic active chlorine species contribute concurrently. Thus, iron addition alone does not accelerate primary degradation but introduces pathways that enhance downstream mineralization. Literature comparisons [41,42,43] emphasize that EF often matches but rarely exceeds AO in short-term degradation kinetics in chloride systems; its advantage lies in the gradual oxidation of refractory intermediates.

The TOC profiles shown in Figure 7 correspond to the same degradation experiments presented for isoproturon (ISP) and alachlor (ALA) under identical operational conditions; thus, mineralization was evaluated for both target compounds using the same samples collected during the kinetic assays; Figure 7c illustrates this difference between degradation and mineralization. While AO achieved complete pollutant disappearance, it produced no measurable decrease in total organic carbon (TOC), confirming that rapid oxidation does not equate to carbon mineralization. In contrast, EF achieved ~15% TOC reduction at both 0.15 mM and 0.50 mM Fe²⁺, indicating the generation of homogeneous ^●OH capable of oxidizing secondary species. Increasing Fe²⁺ above the optimum failed to improve mineralization; rather, it inhibited the process because excess Fe²⁺ acts as a radical sink, consuming ^●OH according to Equation (19), which diminishes overall oxidizing capacity [44,45]. Additional Fe²⁺ can also promote Fe(OH)₃ precipitation or restrict mass transfer, limiting radical access to organics.

Fe²⁺ + ^●OH → Fe³⁺ + OH⁻

(20)

The temporal evolution of dissolved iron (Figure 7d) corroborates the expected Fe²⁺/Fe³⁺ redox cycling intrinsic to EF. During operation, Fe²⁺ is oxidized to Fe³⁺ through the Fenton reaction and electrochemically regenerated at the cathode, maintaining an approximately constant total iron concentration. The quasi-steady state observed confirms the stability of the Fe²⁺/Fe³⁺ loop that sustains radical generation under galvanostatic control. Maintaining pH 3.0 is crucial for this balance: under acidic conditions, iron remains soluble and active, minimizing precipitation and preserving catalytic turnover [46,47]. Deviations toward neutrality disrupt solubility and collapse radical generation.

The mineralization efficiency and energy implications of EF are summarized in Figure 8, which plots the mineralization current efficiency (MCE) and the specific energy consumption per TOC removed (EC_TOC) for 30 mM NaCl, 0.15 mM Fe²⁺, and j = 30 mA cm⁻². MCE decreases progressively with time, whereas EC_TOC increases. This inverse relationship reflects the gradual depletion of easily oxidizable species and the accumulation of more persistent intermediates. As electrolysis proceeds, each incremental TOC removal demands more charge, thereby raising energy consumption per gram of carbon mineralized. Both parameters were calculated using Equation (15).

The observed trends are characteristic of EF systems employing DSA-type anodes (Ti/IrO₂–SnO₂–Sb₂O₅) in chloride media. At pH 3, HClO dominates the chlorine speciation and participates in the Fenton-like mechanism (Equation (19)), supplementing anodic ^●OH and enhancing mineralization at early stages [30]. However, as organics are depleted and Fe³⁺ complexes accumulate, parasitic reactions and electrode side processes consume a larger fraction of the applied current, lowering MCE and raising EC_TOC [48,49]. This behavior underscores a kinetic-energetic trade-off: the initial period of EF operation is the most productive, combining high MCE with low energy cost, while extended electrolysis yields diminishing returns; this time dependence of MCE and energy metrics is consistent with prior EF, PEF, and AO reports, where early time operation offers the best energy rationality and later stages show declining MCE with rising energy cost [50].

Consequently, Figure 8 defines a practical energy-rational window for EF in acidic chloride media. Operating at 30 mA cm⁻² and 0.15 mM Fe²⁺ efficiently exploits the initial regime of high current efficiency and fast mineralization, avoiding unnecessarily long treatments where the energy cost per unit TOC removed escalates sharply. From an operational viewpoint, this condition represents the optimal compromise between kinetic effectiveness, energy economy, and system stability, providing a solid baseline for subsequent photoelectro-Fenton coupling, where photoreduction of Fe³⁺ can further enhance ^●OH regeneration and improve overall mineralization yield [51].

3.5.2. Effect of Fe²⁺ in Photoelectro-Fenton

The introduction of UV irradiation to the electro-Fenton system profoundly modifies the oxidation environment, transforming it into a photoelectro-Fenton (PEF) process (Figure 9). PEF is expected to outperform conventional EF because UV photons photoreduce Fe³⁺ back to Fe²⁺ and photolyze ferric complexes, sustaining the catalytic Fe²⁺/Fe³⁺ loop and enhancing hydroxyl radical formation. This dual photo-activation Fe³⁺ reduction (Equation (10)) and photolysis of carboxylated iron complexes (Equation (21)) maintains radical availability even when Fe²⁺ becomes temporarily depleted by parasitic reactions.

[Fe(OOCR)]²⁺ + hv → Fe²⁺ + CO₂ + R^●

(21)

Consistent with this rationale, Figure 9a shows that isoproturon is rapidly and completely removed across all tested conditions. AO achieves total disappearance in about 10 min, whereas EF and PEF at 0.15 mM Fe²⁺ reach full degradation within ~20 min. Increasing Fe²⁺ to 0.50 mM in PEF does not inhibit ISP oxidation, indicating that photoreduction counteracts the scavenging effect observed under dark EF. The slightly faster AO response is attributed to the strong action of anodically generated active chlorine (Cl₂/HClO/ClO⁻), which dominates near-electrode oxidation, whereas EF and PEF favor bulk ^•OH-driven oxidation. The similarity between EF and PEF for ISP suggests that, for readily oxidizable compounds, light mainly stabilizes Fe²⁺ cycling rather than altering the overall kinetics; this behavior is consistent with the well-documented ability of light to sustain Fe³⁺-Fe²⁺ cycling and offset Fe²⁺ scavenging pathways under acidic conditions [52].

For alachlor (ALA) (Figure 9b), the benefit of photoactivation becomes striking. AO and EF (0.15 mM Fe²⁺) exhibit partial degradation, reaching only about 60% removal after 240 min, reflecting ALA’s higher recalcitrance. In contrast, PEF achieves complete ALA degradation within ~60 min at 0.15 mM Fe²⁺ and ~75 min at 0.50 mM Fe²⁺. The slight delay at the higher iron dose can be ascribed to excess Fe²⁺ scavenging of •OH and increased light attenuation due to Fe³⁺ complexes in the solution [53]. These observations confirm that light-assisted Fe³⁺ reduction efficiently sustains the Fenton cycle while maintaining acidic conditions that prevent iron precipitation and favor homogeneous radical propagation.

Mineralization data (Figure 9c) corroborate these kinetic trends. AO, despite rapid parent decay, produces no measurable TOC decrease, confirming that indirect chlorine oxidation primarily fragments the molecule without achieving mineralization. EF enhances mineralization to about 15% TOC reduction at both Fe²⁺ concentrations, demonstrating partial oxidation of intermediates via homogeneous ^•OH formation, but it is still limited by radical scavenging from chloride and chlorinated by-products [44,54]. Upon UV irradiation (PEF), TOC removal increases dramatically, reaching up to ~76% for 0.15 mM Fe²⁺. The higher efficiency at a lower iron dosage aligns with the principle that excessive Fe²⁺ enhances light screening and promotes recombination, while moderate Fe²⁺ maintains optimal photon penetration and radical yield [49,55].

Table 1 compares the most efficient advanced oxidation processes (AOPs) reported for the removal of alachlor, paraquat, diquat, and isoproturon, offering a useful framework to interpret our own findings. Although several literature systems such as UV/Fe–citrate, gamma radiolysis with H₂O₂, or coupled photo-Fenton–biological treatments achieve TOC removals close to or above 80%, it is important to note that most of these studies were performed at very small volumes (typically 15–100 mL). In contrast, our PEF–BDD system operated at a substantially larger reaction volume (500 mL) and still achieved 79% TOC removal, placing our results within the upper range of reported efficiencies despite the more demanding operational scale. Compared with photocatalytic and catalytic ozonation processes, which commonly reach ≤60–80% mineralization, our system demonstrates competitive performance and highlights the robustness of electrochemical technologies for treating recalcitrant herbicides at higher working volumes. This contextual comparison underscores both the strengths of our approach and the opportunities for further optimization.

Table 1. Summary of the best advanced oxidation processes reported for the degradation and mineralization of alachlor, paraquat, diquat, and isoproturon.

Process	Contaminant	Experimental Conditions	%TOC Removal	ECTOC	Ref.
PEF (BDD/ADE, 300 mA)	Alachlor (0.60 mM)	Volume 100 mL; pH 3; 0.05 M Na2SO4; Fe2+ = 0.5 mM; UVA 6 W; 300 mA; T = 25 °C	98%	≈5 Ah·L−1	[56]
BDD electro-oxidation (1.55 mA/cm2)	Paraquat (70 mg/L)	Volume 0.65 L; 0.05 M Na2SO4; J = 1.55 mA/cm2; flow 500 mL/min	91%	0.10 kWh/gTOC	[57]
BDD electro-oxidation (1.0 mA/cm2)	Diquat (70 mg/L)	Volume 0.65 L; 0.05 M Na2SO4; J = 1.0 mA/cm2; flow 500 mL/min	92%	0.24 kWh/gTOC	[57]
UV/Fe(II)–citrate/H2O2	Alachlor (10 mg/L)	Volume 20 mL; pH 5; Fe2+ 2 × 10−4 M; H2O2 4 × 10−3 M; citrate 5 × 10−4 M; Xe lamp 990 W	≈100% (26 h)	Not reported	[43]
Gamma irradiation + H2O2 (1.0 mM)	Alachlor (40 µM)	Volume 15 mL; Co-60 source; 20 kGy; H2O2 = 1 mM; aerobic; 20 °C	83.8%	Not reported	[58]
Cu/Al2O3-catalyzed ozonation	Alachlor (100 mg/L)	Volume 75 mL; pH 6.39; Cu/Al2O3 honeycomb; 0.488 mg O3/min; 20 °C	≈60% (180 min)	Not reported	[59]
Solar TiO2/H-MOR (15 wt%)	Isoproturon (1.14 × 10−4 M)	Volume 50 mL; catalyst load 1.5 g/L; solar light 11–15 h; open dish reactor	≈80% (5 h)	Not reported	[60]
Photoelectro-Fenton (DSA/BDD)	Isoproturon (50 mM) + Alachlor (50 mM)	Volume 500 mL; pH 3; J = 30 mA/cm−2; 30 mM NaCl; 240 min, 0.15 mM Fe2+; UV source 290 nm, 32 W m−2	76% (240 min)	1.4 kWh/gTOC	This work

Table 1. Summary of the best advanced oxidation processes reported for the degradation and mineralization of alachlor, paraquat, diquat, and isoproturon.

Process	Contaminant	Experimental Conditions	%TOC Removal	ECTOC	Ref.
PEF (BDD/ADE, 300 mA)	Alachlor (0.60 mM)	Volume 100 mL; pH 3; 0.05 M Na2SO4; Fe2+ = 0.5 mM; UVA 6 W; 300 mA; T = 25 °C	98%	≈5 Ah·L−1	[56]
BDD electro-oxidation (1.55 mA/cm2)	Paraquat (70 mg/L)	Volume 0.65 L; 0.05 M Na2SO4; J = 1.55 mA/cm2; flow 500 mL/min	91%	0.10 kWh/gTOC	[57]
BDD electro-oxidation (1.0 mA/cm2)	Diquat (70 mg/L)	Volume 0.65 L; 0.05 M Na2SO4; J = 1.0 mA/cm2; flow 500 mL/min	92%	0.24 kWh/gTOC	[57]
UV/Fe(II)–citrate/H2O2	Alachlor (10 mg/L)	Volume 20 mL; pH 5; Fe2+ 2 × 10−4 M; H2O2 4 × 10−3 M; citrate 5 × 10−4 M; Xe lamp 990 W	≈100% (26 h)	Not reported	[43]
Gamma irradiation + H2O2 (1.0 mM)	Alachlor (40 µM)	Volume 15 mL; Co-60 source; 20 kGy; H2O2 = 1 mM; aerobic; 20 °C	83.8%	Not reported	[58]
Cu/Al2O3-catalyzed ozonation	Alachlor (100 mg/L)	Volume 75 mL; pH 6.39; Cu/Al2O3 honeycomb; 0.488 mg O3/min; 20 °C	≈60% (180 min)	Not reported	[59]
Solar TiO2/H-MOR (15 wt%)	Isoproturon (1.14 × 10−4 M)	Volume 50 mL; catalyst load 1.5 g/L; solar light 11–15 h; open dish reactor	≈80% (5 h)	Not reported	[60]
Photoelectro-Fenton (DSA/BDD)	Isoproturon (50 mM) + Alachlor (50 mM)	Volume 500 mL; pH 3; J = 30 mA/cm−2; 30 mM NaCl; 240 min, 0.15 mM Fe2+; UV source 290 nm, 32 W m−2	76% (240 min)	1.4 kWh/gTOC	This work

This pronounced improvement stems from two complementary photo-mechanisms: (i) Fe³⁺ photoreduction, which continuously regenerates the active Fe²⁺ catalyst, and (ii) photolysis of ferric carboxylate complexes, which releases CO₂ and organic radicals that initiate secondary oxidation cascades (Equation (21)). Both mechanisms sustain radical generation even as the organic load diminishes, yielding extended mineralization efficiency beyond the electrochemical steady state.

The temporal profile of dissolved iron (Figure 9d) confirms this regenerative behavior. Total dissolved iron remains nearly constant during PEF, while Fe²⁺ and Fe³⁺ fractions oscillate, evidencing the dynamic Fe²⁺/Fe³⁺ cycling characteristic of homogeneous photochemical catalysis. Maintaining pH 3.0 is critical to this stability: at acidic pH, Fe³⁺ remains soluble as photoactive hydroxy complexes, Fe(OH)²⁺ and [Fe(OH)₂]⁺, which undergo photoreduction and directly generate ^•OH upon irradiation [61,62]. Deviations toward neutrality would precipitate Fe(OH)₃, terminating the Fenton cycle and reducing photon utilization efficiency.

Overall, the results establish a clear hierarchy among the processes. Anodic oxidation (AO) provides the fastest parent removal via indirect oxidants but negligible mineralization; electro-Fenton (EF) introduces true carbon oxidation through homogeneous ^•OH production; and photoelectro-Fenton (PEF) achieves a step-change in both kinetics and mineralization efficiency by coupling light-driven Fe³⁺ regeneration with electrochemical radical formation. Operated at pH 3.0 and moderate Fe²⁺ concentrations, PEF maximizes radical flux while minimizing scavenging and UV screening, delivering an energy-rational and environmentally sustainable pathway for deep oxidation of persistent herbicides [63].

4. Conclusions

This work demonstrates that active chlorine electrosynthesis on Ti/IrO₂-SnO₂-Sb₂O₅ is robust and tunable (pH and [Cl⁻]) and that electrolyte support outweighs current density for anodic oxidation. A mixed Cl⁻/SO₄⁻² matrix at moderate j = 30 mA cm⁻² delivers rapid isoproturon abatement, markedly improves alachlor conversion via a synergistic ^●OH/HClO-ClO⁻/S₂O₈⁻² pathway, and preserves the acidic window needed for the Fenton process. While AO affords very fast removal, only electro-Fenton introduces measurable mineralization, and photoelectro-Fenton produces a step-change in both kinetics (notable for alachlor) and TOC reduction by sustaining Fe (II) and Fe (III) cycling and complementary photochemical routes. Process metrics indicate that mineralization current efficiency declines and the specific energy per TOC removed increases with time, defining an energy-rational operating window in which AO is used for primary abatement and EF-PEF are applied to achieve deeper degradation and mineralization under acidic conditions (pH 3.0).

In chloride-rich AO, pH drift and elevated cell voltage can arise, increasing energy demand and the likelihood of oxychlorinated by-products (chlorate and perchlorate) during prolonged operation. The mineralization gains observed for EF and PEF are constrained by radical scavenging by chloride and reaction intermediates and attenuation at higher Fe²⁺ levels, while effective iron management remains strongly dependent on maintaining pH to avoid precipitation. Performance in real effluents, containing natural organic matter, carbonates, and variable salinity, may be attenuated relative to that in synthetic effluents. Future work could prioritize comprehensive by-product speciation, including chlorite (ClO⁻₂), chlorate (ClO⁻₃), and perchlorate (ClO⁻₄), together with a toxicological assessment under the identified operating window. It will also be important to enhance EF and PEF via cathode engineering for improved H₂O₂ electrogeneration, enhance low Fe²⁺ regimes under UV/solar to limit radical scavenging, and conduct electrode stability studies (lifetime and passivation) supported by surface analytics. Techno-economic and life cycle assessments are needed to achieve regulatory targets at minimum cost and carbon footprint.