Photo(solar)-Activated Hypochlorite Treatment: Radicals Analysis Using a Validated Model and Assessment of Efficiency in Organic Pollutants Degradation

Abstract

Reactive oxygen species (ROS) and reactive chlorine species (RCS) and their involvement in the degradation process are explored in this work by thorough kinetic modeling of the solar-activated hypochlorite degradation of Rhodamine B (RhB) dye. The kinetic modeling enabled the determination of rate constants for both radical and non-radical pathways of hypochlorite and the oxidation of RhB by free radicals. Using COPASI^® software, fed with a kinetics mechanism of 144 chemical reactions, the free radical kinetic model accurately fitted experimental data under various conditions, including temperatures ranging from 25 to 55 °C and initial hypochlorite concentrations from 300 to 1000 µM, at a controlled pH of 11. Results indicate that increasing hypochlorite dosages and temperatures enhance free radical concentrations and RhB degradation rates. ^•OH and ClO^• radicals were quantified as primary contributors to RhB degradation, while ozone played a minor role. The model provides profiles for ROS and RCS, details on radicals distribution in RhB degradation, and predictions of rate constants for the photolysis of ClO⁻: k_R1 = 2.67 × 10⁻⁴ s⁻¹ for the radical pathway (ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O^•− + Cl^•), and k_R2 = 1.88 × 10⁻⁵ s⁻¹ and k_R3 = 0 s⁻¹ for the non-radical pathway (i.e., ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O(³P) + Cl⁻ and ClO⁻ $\stackrel{h\nu }{\to }$ O(¹D) + Cl⁻, respectively). The rate constants for RhB reactions with O^•−, Cl^•, Cl₂^•− and ClO^• were predicted to be 4.8 × 10⁹ M⁻¹ s⁻¹, 1.45 × 10⁹ M⁻¹ s⁻¹, 2.5 × 10⁷ M⁻¹ s⁻¹ and 8.7 × 10⁴ M⁻¹ s⁻¹, respectively. Lower rate constants were predicted for RhB reactions with HOCl^•−, HO₂^•, O₂^•−, and O(³P), with values of 4.1 × 10⁴ M⁻¹ s⁻¹, 7.3 × 10⁵ M⁻¹ s⁻¹, 3.6 × 10⁴ M⁻¹ s⁻¹, and 0.40 M⁻¹ s⁻¹, respectively.

1. Introduction

Water utilities are increasingly challenged by the presence of organic pollutants in drinking water sources, including textile dyes, aromatics, pharmaceuticals, and pesticides. These substances find their way into the environment through various pathways such as industrial effluents, municipal sewage, and agricultural runoff [1,2,3,4]. Conventional water treatment systems often struggle to effectively remove these contaminants [5]. While oxidants differ in their selectivity, the generation of highly reactive radicals, such as hydroxyl radicals (^•OH), sulfate radicals (SO₄^•−), chlorine radicals (Cl^•), chlorite radicals (ClO^•), and di-chloride radicals (Cl₂^•−), enables the oxidation of a wide range of organic compounds, thanks to their high measured rate constants (~10⁶–10⁹ M⁻¹ s⁻¹) toward organic pollutants [6]. Advanced oxidation processes (AOPs) utilize these radicals to break down numerous organic pollutants [7,8,9,10]. UV-based techniques are regarded as robust AOPs for water purification [11,12,13]. Chlorine is particularly advantageous for sunlight-driven processes because it can absorb longer wavelengths, unlike H₂O₂ and persulfate-based AOPs, which are limited to UV-C light due to their lower molar absorptivity [14,15,16]. Chlorine photolysis stands out for its use of a cost-effective and widely available disinfectant, making it the more economical choice [17]. Additionally, chlorine photolysis eliminates the need for residual chlorine quenching, a necessary procedure in UV/persulfate and UV/H₂O₂ systems [5].

UV/chlorine systems commonly employ low-pressure mercury vapor (LP-UV) lamps as their primary UV light source. However, for the enhancement of the energy efficacy of the UV/chlorine process, various UV lamps have been used, including medium-pressure mercury vapor (MP-UV) lamps, excimer lamps, and UV light-emitting diodes (UV-LED). Furthermore, a limited number of studies have investigated the feasibility of utilizing solar radiation as a UV light source for UV/chlorine AOP [18,19,20]. The photocatalytic activation of chlorine is affected by several parameters (solution pH, irradiation wavelength, chlorine dose, water matrices, solution temperature, etc.). For example, in the experimental work of Yin et al. [21], UV-LEDs emitting at four wavelengths within the UV-C and near-UV-C spectrum (i.e., 257.7, 268, 282.3, and 301.2 nm) have been used to examine the influence of wavelength on chlorine photolysis and the subsequent generation of reactive radicals. The fluence-based photodecay rates of hypochlorous acid (HOCl) and hypochlorite (OCl⁻) exhibited a monotonic correlation with their respective molar absorption coefficients and quantum yields. Notably, chlorine photodecay rates were significantly more influenced by molar absorption coefficients (0.949) than by quantum yields (0.055). The modeling outcomes indicated that the maximum fluence-based rate constant (1.46 × 10⁻⁴ m² J⁻¹) was achieved at 289.7 nm and a pH of 9.95. The influence of wavelength on photodecay rates was more pronounced under alkaline conditions compared to acidic conditions. Additionally, the sensitivity to pH was greatest at the longest wavelength studied. The generation of ^•OH and reactive chlorine species (RCS) exhibited inverse wavelength dependencies at pH 6. However, at pH 7, the formation of ^•OH and RCS increased with increasing wavelength. Furthermore, a higher concentration of ^•OH was observed at pH 6 compared to pH 7, while the formation of RCS demonstrated the opposite pH dependence [21]. On the other side, the UV/chlorine process could be enhanced in the presence of catalysts. In the recent work of Cheng et al. [22], using the UV₃₆₅/TiO₂/chlorine process for the removal of carbamazepine, the obtained apparent first-order rate constant was found to be 34.2 and 3.9 times higher compared to those without TiO₂ and chlorine, respectively. In this process, chlorine plays a dual role. Firstly, it acts as a catalyst, increasing the production of hydroxyl radicals (^•OH) without undergoing consumption. Secondly, chlorine serves as a radical precursor, contributing to the formation of ^•OH and reactive chlorine species [22]. The k’_CBZ increased with rising TiO₂ dose within the range of 1.0 to 20.0 mg/L and with increasing light intensities from 0.1 to 0.33 mW/cm² and with decreasing chlorine dose from 5.0 to 1.0 mgCl₂/L. While an increase in pH led to a higher overall concentration of reactive species, the transformation of ^•OH and Cl^• into the less reactive ClO^• species resulted in a decrease in k’_CBZ from pH 6.0 to 9.0. In [23], the visible light/g-C₃N₄/chlorine process significantly enhanced the pseudo-first-order degradation rate constant of carbamazepine. Compared to processes without g-C₃N₄ or chlorine, this combined approach achieved degradation rates that were 16 and 7 times higher, respectively. Moreover, the system demonstrated sustained performance over repeated use cycles. A notable advantage of the vis-light/g-C3N4/chlorine process is its exceptional performance in the presence of natural organic matter (NOM). Unlike UV/TiO₂ or UV/chlorine AOPs, this system is less affected by NOM due to its reduced light absorption at visible wavelengths and its lower scavenging of surface-bound reactive species [23].

The photolysis of chlorine involves intricate chemistry influenced by the wavelength and the form of chlorine present, which depends on the solution’s pH [24]. Hypochlorous acid (HClO), with a pK_a of 7.5, is predominant in the pH range of 4.0–6.0, whereas hypochlorite (ClO⁻) dominates at pH levels above 10, such as in commercial sodium hypochlorite solutions with a pH around 12 [25]. HClO primarily absorbs at 236 nm (ε~102 M⁻¹ cm⁻¹), while the absorption band of hypochlorite extends up to 400 nm, peaking at 294 nm (ε~375 M⁻¹ cm⁻¹), making it suitable for solar-light activation. Photolysis of HClO produces ^•OH and Cl^• radicals (HOCl +hν → ^•OH + Cl^•) [26]. The quantum yield (Φ) for this reaction decreases with increasing wavelength: 0.278 at 253.4 nm, 0.127 at 313 nm, and 0.08 at 365 nm [26]. In contrast, the photolysis of hypochlorite (OCl⁻) is more complex, generating O^•− (ClO⁻ + hν → O^•− + Cl^•), and two excited oxygen states, O(¹D) and O(³P) (OCl⁻ +hν → O(¹D) + Cl⁻ and OCl⁻ + hν → O(³P) + Cl⁻) [26]. For the O(³P) pathway, the quantum yield of ClO⁻ photolysis increases with wavelength: 0.074 at 253.7 nm, 0.075 at 313 nm, and 0.28 at 365 nm [26]. Therefore, there is significant potential to utilize solar light, which includes UV light between 300 and 400 nm, to activate hypochlorite, particularly as it has an absorption band with a peak at 303 nm. However, the O(¹D) pathway is not favored with solar light, as the quantum yield of ClO⁻ at 365 nm is zero [26]. ClO^• from the reaction of Cl^• or ^•OH with chlorine (Reactions 30, and 44 in

Table 1. Model equations for RhB degradation in basic solar-photoactivated hypochlorite solution (pH 11). Abbreviations: TBA: tert-butyl alcohol, BA: benzoic acid, Prod_1 to Prod_11: unspecified species.

Name		Reaction	Rate Constant	Ref.
R1	ClO− photolysis Rxns	ClO− $\stackrel{h\nu }{\to }$ O●− + Cl●	(0.028–2.67) × 10−4 s−1	This study
R2		ClO− $\stackrel{h\nu }{\to }$ O(3P) + Cl−	(1.88–2.46) × 10−5 s−1	This study
R3		ClO− $\stackrel{h\nu }{\to }$ O(1D) + Cl−	0 s−1	This study
R4	O(1D)/O(3P) Rxns	O(1D) + H2O → 2●OH	1.2 × 1011 M−1s−1	[5]
R5		O(3P) + O2 → O3	4.0 × 109 M−1s−1	[5]
R6		O(3P) + ClO− → ClO2−	9.4 × 109 M−1s−1	[5]
R7		O(3P) + H2O2 → ●OH + HO2●	1.6 × 109 M−1s−1	[5]
R8		O(3P) + HO2− → ●OH + O2●−	5.3 × 109 M−1s−1	[5]
R9		O(3P) + OH− → HO2−	4.2 × 108 M−1s−1	[5]
R10		O3 → O2 + O(3P)	4.5 × 10−6 s−1	[5]
R11	Speciations Rxns	HOCl → H+ + ClO−	1.41 × 103 s−1	[5]
R12		H+ + ClO− → HOCl	5.0 × 1010 M−1s−1	[33]
R13		HOCl + Cl− → Cl2 + H2O	0.182 M−1s−1	[5]
R14		Cl2 + H2O → HOCl + Cl− + H+	0.27 M−1s	[14]
R15		HOCl + Cl− → Cl2OH−	1.5 × 104 M−1s−1	[5]
R16		Cl2 + OH− → HOCl + Cl−	1.0 × 109 M−1s−1	[5]
R17		HCl → H+ + Cl−	8.6 × 1016 s−1	[14]
R18		H+ + Cl− → HCl	5.0 × 1010 M−1s−1	[14]
R19		H2O2 → H+ + HO2−	0.13 s−1	[33]
R20		H+ + HO2− → H2O2	5.0 × 1010 M−1s−1	[33]
R21		H2O → H+ + OH−	0.001 s−1	[14]
R22		H+ + OH− → H2O	1.0 × 1011 M−1s−1	[14]
R23	●OH Rxns	●OH + HOCl → ClO● + H2O	2.0 × 109 M−1s−1	[34]
R24		2●OH → H2O2	5.5 × 109 M−1s−1	[34]
R25		●OH + ClO● → ClO2− + H+	1.0 × 109 M−1s−1	[5]
R26		●OH + H2O2 → HO2● + H2O	2.7 × 107 M−1s−1	[33]
R27		●OH + O2●− → O2 + OH−	7.0 × 109 M−1s−1	[14]
R28		●OH + Cl− → Cl● + OH−	1.1 × 109 M−1s−1	[5]
R29		●OH + Cl2●− → HOCl + Cl−	1.0 × 109 M−1s−1	[5,14]
R30		●OH + ClO− → ClO● + OH−	8.8 × 109 M−1s−1	[14]
R31		●OH + HO2● → O2 + H2O	6.6 × 109 M−1s−1	[5,14]
R32		●OH + Cl− → HOCl●−	4.3 × 109 M−1s−1	[14]
R33		●OH + ClO2− → OH− + ClO2●	4.5 × 109 M−1s−1	[5]
R34		●OH + ClO3− → ClO3● + OH−	1.0 × 106 M−1s−1	[5]
R35		●OH + HO2− → HO2● + OH−	7.5 × 109 M−1s−1	[33]
R36		●OH + HO2− → H2O + O2●−	7.05 × 109 M−1s−1	[33]
R37		●OH + OH− → O●− + H2O	1.25 × 1010 M−1s−1	[5]
R38		●OH + O●− → HO2−	2.0 × 1010 M−1s−1	[5]
R39		●OH + ClO2● → ClO3− + H+	4.0 × 109 M−1s−1	[5]
R40		●OH + O3●− → HO2● + O2●−	8.5 × 109 M−1s−1	[5]
R41		●OH + O3 → O2 + HO2●	1.05 × 108 M−1s−1	[5]
R42	Cl● Rxns	2Cl● → Cl2.	8.8.8 × 107 M−1s−1	[5,14]
R43		Cl● + HOCl → Cl− + ClO● + H+	3.0 × 109 M−1s−1	[35]
R44		Cl● + ClO− → Cl− + ClO●	8.3 × 109 M−1s−1	[35]
R45		Cl● + Cl2 → Cl3●	5.3 × 108 M−1s−1	[5]
R46		Cl● + H2O → HOCl●− + H+	2.05 × 105 M−1s−1	[5]
R47		Cl● + Cl− → Cl2●−	8.5 × 109 M−1s−1	[14]
R48		Cl● + OH− → HOCl●−	1.8 × 1010 M−1s−1	[14]
R49		Cl● + H2O2 → HO2● + Cl− + H+	2.0 × 109 M−1s−1	[14]
R50		Cl● + ClO3− → Prod_1	1.0 × 106 M−1s−1	[5]
R51		Cl● + ClO2● → Prod_2	1.0 × 109 M−1s−1	[5]
R52		Cl● + Cl2●− → Cl2 + Cl−	2.1 × 109 M−1s−1	[5]
R53		Cl● + ClO2− → ClO2● + Cl−	7.0 × 109 M−1s−1	[5]
R54	Cl2●− Rxns	2Cl2●− → Cl2 + 2Cl−	8.0 × 108 M−1s−1	[5]
R55		Cl2●− + H2O → HOCl●− + Cl− + H+	1.3 × 103 M−1s−1	[5]
R56		Cl2●− + OH− → HOCl●− + Cl−	4.5 × 107 M−1s−1	[5]
R57		Cl2●− + ClO− → ClO● + 2Cl−	2.9 × 108 M−1s−1	[5]
R58		Cl2●− + O2●− → O2 + 2Cl−	2.0 × 109 M−1s−1	[5,14]
R59		Cl2●− + HO2● → O2 + 2Cl− + H+	3.0 × 109 M−1s−1	[14]
R60		Cl2●− + H2O2 → HO2● + 2Cl− + H+	1.4 × 105 M−1s−1	[5]
R61		Cl2●− → Cl● + Cl−	6.0 × 104 s−1	[5]
R62		Cl2●− + ClO2● → Prod_3	1.0 × 109 M−1s−1	[5]
R63	ClO● Rxns	2ClO● → Cl2O2	2.5 × 109 M−1s−1	[36]
R64		Cl2O2 + H2O → HOCl + ClO2− + H+	4.5 × 107 M−1s−1	[36]
R65		Cl2O2 + OH− → ClO− + ClO2− + H+	2.5 × 109 M−1s−1	[36]
R66		ClO● + ClO2− → ClO2● + ClO−	9.4 × 108 M−1s−1	[5]
R67	Ozone Rxns	O3 + OH− → O2 + HO2−	48 M−1s−1	[6]
R68		O3 + HO2− → O3●− + HO2●	5.5 × 106 M−1s−1	[5]
R69		O3 + HO2● → H+ + O2 + O3●−	1.6 × 109 M−1s−1	[5]
R70		O3 + ClO− → 2O2 + Cl−	110 M−1s−1	[5]
R71		O3 + ClO− → O2 + ClO2−	30 M−1s−1	[5]
R72		O3 + Cl− → O2 + ClO−	0.0016 M−1s−1	[5]
R73		O3 + H2O2 → O2 + ●OH + HO2●	0.0272 M−1s−1	[5]
R74		O3 + HO2− → O2 + ●OH + O2●−	5.5 × 106 M−1s−1	[5]
R75		O3 + Cl2●− → Prod_4	9.0 × 107 M−1s−1	[5]
R76		O3 + O2●− → O2 + O3●−	1.55 × 109 M−1s−1	[5]
R77		O3 + ClO2● → O2 + ClO3●	1230 M−1s−1	[5]
R78		O3 + ClO3− → Prod_5	0.0001 M−1s−1	[5]
R79		O3 + H+ → Prod_6	0.0004 M−1s−1	[5]
R80		O3 + ClO2− → O3●− + ClO2●	2.01 × 106 M−1s−1	[5]
R81	ROS Rxns	H2O2 + Cl2 → 2HCl + O2	1.3 × 104 M−1s−1	[5]
R82		H2O2 + HOCl → HCl + H2O + O2	1.1 × 104 M−1s−1	[5]
R83		H2O2 + ClO− → Cl− + H2O + O2	1.7 × 105 M−1s−1	[5]
R84		H2O2 + HO2● → O2 + ●OH + H2O	3 M−1s−1	[5]
R85		H2O2 + O2●− → O2 + ●OH + OH−	0.13 M−1s−1	[5]
R86		H2O2 + O●− → O2●− + H2O	4.0 × 108 M−1s−1	[5]
R87		HO2− + O●− → HO2● + 2OH-	5.0 × 108 M−1s−1	[5]
R88		HO2− + O●− → ●OH + O2●−	4.0 × 108 M−1s−1	[5]
R89		HO2− + ClO2● → HO2● + ClO2-	9.57 × 104 M−1s−1	[5]
R90		HO2● + HOCl → Cl● + OH− + O2 + H+	7.5 × 106 M−1s−1	[5]
R91		HO2● + Cl2 → Cl2●− + H+ + O2	1.0 × 109 M−1s−1	[5]
R92		HO2● → H+ + O2●−	1.6 × 105 M−1s−1	[5]
R93		HO2● + O2●− → H2O2 + O2 + OH-	7.9 × 107 M−1s−1	[5]
R94		HO2● + O2●− → HO2− + O2	9.7 × 107 M−1s−1	[5]
R95		2 HO2● → H2O2 + O2	8.3 × 105 M−1s−1	[5]
R96		HO2● + ClO2● → Prod_7	1.0 × 106 M−1s−1	[5]
R97		O2●− + HOCl → Cl● + OH− + O2	7.5 × 106 M−1s−1	[5]
R98		O2●− + HOCl → Cl− + ●OH + O2	7.5 × 106 M−1s−1	[5]
R99		O2●− + ClO− → Cl● + 2 OH− + O2	2.0 × 108 M−1s−1	[5]
R100		O2●− + Cl2 → Cl2●− + O2	1.0 × 109 M−1s−1	[5]
R101		O2●− + O●− → 2OH− + O2	6.0 × 108 M−1s−1	[5]
R102		O2●− + Cl− → Prod_8	140 M−1s−1	[5]
R103		O2●− + ClO2● → O2 + ClO2-	3.15 × 109 M−1s−1	[5]
R104		O2●− + ClO2− → Prod_9	40 M−1s−1	[5]
R105		O2●− + H+ → HO2●	5.0 × 1010 M−1s−1	[33]
R106		O2●− + ClO3− → Prod_10	3.2 × 103 M−1s−1	[5]
R107		O3●− → O2 + O●−	3.2 × 103 s−1	[5,6]
R108		O3●− + ClO● → ClO− + O3	1.0 × 109 M−1s−1	[5]
R109		O3●− + H+ → O2 + ●OH	9.0 × 1010 M−1s−1	[5]
R110		O3●− + O●− → 2O2●−	7.0 × 108 M−1s−1	[5]
R111		O3●− + ClO2● → O2 + ClO3−	1.8 × 105 M−1s−1	[5]
R112		2O3●− → Prod_11	9.0 × 108 M−1s−1	[5]
R113		O3●− + ClO2● → ClO2− + O3	3.15 × 109 M−1s−1	[5]
R114		O●− + ClO− + H2O → ClO● + 2OH−	2.3 × 108 M−1s−1	[5]
R115		O●− + H2O → ●OH + OH−	1.8 × 106 M−1s−1	[14]
R116		O●− + O2 → O3●−	3.5 × 109 M−1s−1	[5]
R117		O●− + ClO2− → OH− + ClO2●	1.95 × 108 M−1s−1	[5]
R118		2O●− → O22 −	4.65 × 109 M−1s−1	[5]
R119		O●− + ClO2● → ClO3−	2.7 × 109 M−1s−1	[5]
R120	Other RCS Rxns	HOCl●− → Cl− + ●OH	6.1 × 109 M−1s−1	[37]
R121		HOCl●− + H+ → Cl● + H2O	2.1 × 1010 M−1s−1	[5]
R122		HOCl●− + Cl− → Cl2●− + OH−	1.0 × 104 M−1s−1	[5]
R123		ClO2● → Cl● + O2	6.7 × 109 M−1s−1	[5]
R124	Pollutants Rxns	P + ●OH → Prod_●OH	2.5 × 1010 M−1s−1	[33]
R125		P + O●− → Prod_O●−	4.8 × 109 M−1s−1	This study
R126		P + O3 → Prod_O3	2450 M−1s−1	[38]
R127		P + Cl● → Prod_Cl●	1.45 × 109 M−1s−1	This study
R128		P + ClO● → Prod_ClO●	8.70 × 104 M−1s−1	This study
R129		P + Cl2●− → Prod_Cl2●−	2.50 × 107 M−1s−1	This study
R130		P + HOCl●− → Prod_HOCl●−	4.10 × 104 M−1s−1	This study
R131		P + HO2● → Prod_HO2●	7.30 × 105 M−1s−1	This study
R132		P + O2●− → Prod_O2●−	3.60 × 104 M−1s−1	This study
R133		P + O(3P) → Prod_O(3P)	0.40 M−1s−1	This study
R134	TBA Rxns	TBA + Cl● → Prod_Cl●-TBA	3 × 108 M−1s−1	[39]
R135		TBA + ClO● → Prod_ClO●-TBA	1.30 × 107 M−1s−1	[39]
R136		TBA + Cl2●− → Prod_Cl2●−-TBA	700 M−1s−1	[39]
R137		TBA + ●OH → Prod_●OH-TBA	3.80 × 108 M−1s−1	[40]
R138		TBA + O3 → Prod_O3-TBA	0.003 M−1s−1	[41]
R139		TBA + O●− → Prod_O●−-TBA	5 × 108 M−1s−1
R140	BA Rxns	BA + Cl● → Prod_Cl●-BA	1.8 × 1010 M−1s−1	[39]
R141		BA + ClO● → Prod_ClO●-BA	<3 × 106 M−1s−1	[42]
R142		BA + Cl2●− → Prod_Cl2●−-BA	2 × 106 M−1s−1	[5]
R143		BA + ●OH → Prod_●OH-BA	5.27 × 109 M−1s−1	[5]
R144		BA + O●− → Prod_O●−-BA	4 × 107 M−1s−1	[5]

) and Cl₂^•− from the reaction of Cl^• with Cl⁻ (Reaction 47 in Table 1) are examples of further chlorine radicals that can be produced during chlorine photo-dissociation. Cl^• is more selective than the non-selective ^•OH and interacts preferentially with substrates that are rich in electrons [27]. Cl₂^•− preferentially interacts with a variety of organic compounds, although being typically less reactive than ^•OH and Cl^• [28,29]. These radicals are essential for the degradation of several water contaminants and have high redox potentials (2.43 V for Cl^•, 2.13 V for Cl₂^•−, and 1.5–1.8 V for ClO^•) [30]. With regard to organic molecules, the rates of reaction of these radicals are around~10²–10⁶ M⁻¹ s⁻¹ for Cl₂^•−, ~10⁷–10⁹ M⁻¹s⁻¹ for ClO^•, and 10⁸–10¹¹ M⁻¹s⁻¹ for Cl^• and ^•OH [15,24,31,32].

The UV/chlorine process demonstrated high degradation rates under various conditions. For example, in their analysis of NOM degradation through the UV/chlorine process, Wang et al. [32] found that this method is particularly efficient at removing chromophores (~ 80%) and fluorophores (76.4–80.8%), while the reduction of dissolved organic carbon (DOC) is less pronounced at 15.1–18.6%. The contribution of hydroxyl radicals was found to be 1.4 times more than that of chlorine radicals (Cl^•). Similarly, in [31] a systematic investigation was conducted to assess the degradation of three lipid regulators—gemfibrozil, bezafibrate, and clofibric acid—via UV/chlorine treatment. The chlorine oxide radical (ClO^•) was identified as a primary contributor to the degradation of gemfibrozil and bezafibrate, with corresponding second-order rate constants of 4.2 (±0.3) × 10⁸ M⁻¹ s⁻¹ and 3.6 (±0.1) × 10⁷ M⁻¹ s⁻¹, respectively. In contrast, the degradation of clofibric acid was predominantly attributed to UV photolysis and hydroxyl radicals. A linear correlation was observed between the first-order rate constants (k’) for the degradation of gemfibrozil and bezafibrate and increasing chlorine dosage [31]. Conversely, an inverse relationship was found between the k’ values for the degradation of gemfibrozil, bezafibrate, and clofibric acid and pH within the range of 5.0 to 8.4. Despite this, the contribution of reactive chlorine species (RCS) exhibited an upward trend across the investigated pH range. In [43], hydroxyl radicals (^•OH) and chlorine radicals (Cl^•) were found to be significant contributors to the degradation of benzoic acid, whereas the involvement of other reactive species, including the dichloride radical (Cl₂^•−) and the superoxide radical (O^•−), was deemed insignificant. The overall rate of benzoic acid (BA) degradation diminished as the pH ascended from 6 to 9 [43]. Notably, the relative contributions of HO^• and Cl^• to the degradation process shifted from 34.7% and 65.3%, respectively, at pH 6 to 37.9% and 62%, respectively, at pH 9 under the experimental conditions considered. Deng et al. [44] examined the degradation of Ciprofloxacin (CIP) using the UV/chlorine process. Over 30 min, UV photolysis and dark chlorination achieved only 41.2% and 30.5% CIP degradation, respectively. In contrast, the synergistic UV/chlorine process resulted in a significantly enhanced CIP removal of 98.5% within 9 min. Under neutral aqueous conditions, Ciprofloxacin exhibited the highest pseudo-first-order reaction rate constants for degradation. Among the reactive species involved, e_aq⁻ was the primary contributor, followed by Cl^•, HO^•, and UV photolysis.

Comprehending the generation of reactive species is essential for the efficient application of chlorine photolysis in water treatment. The distribution of reactive oxidants in a UV/chlorine system is mostly determined by the pollutant’s reactivity with the produced species as well as processing variables including chlorine dose, solution pH, and irradiation wavelength. To assess the contributions of ^•OH, RCS, and other oxidants to pollutant degradation, an innovative approach has been developed. This method combines specific radical quenchers with a free radical kinetics model, enabling the determination of the profiles of different species after model validation through experimental quenching tests. Utilizing this technique, Djaballah et al. [45] analyzed the distribution of key free radicals involved in the degradation of reactive green 12 at 254 nm, with varying solution pH and chlorine dosages. Their findings indicated that ^•OH was the predominant oxidant, as determined using the quenchers technique. Bulman et al. [5] employed the same combination approach to study the generation of reactive oxidants during chlorine photolysis across different irradiation wavelengths (254, 311, and 365 nm) and pH levels (6.0–10.0). Nitrobenzene, cinnamic acid, and benzoate were selected as probe chemicals for hydroxyl radicals, RCS, and ozone. According to the study, steady-state concentrations of ^•OH and Cl^• were higher in acidic conditions under 254 or 311 nm radiation, while ozone levels peaked at 254 nm radiation in alkaline conditions. Because of the greater molar absorptivity of hypochlorite under high-wavelength irradiation, the chlorine decay rate constants increased with pH. The scientists came to the conclusion that kinetic modeling is an effective method for investigating the processes of radicals photochemistry in the UV/chlorine process. However, both of the aforementioned studies utilized UV light sources, and no similar approach has been applied to quantify free radical generation and utilization in solar-activated chlorine, particularly focusing on hypochlorite. The investigation of this approach—utilizing solar light under highly basic conditions—is of paramount importance compared to other light sources. This is due to its potential to reduce future operating costs of this technique (UV/chlorine) by (i) harnessing sunlight (simulated here by a solar light simulator), a renewable energy source, and (ii) its adaptability to the basic pH levels commonly found in industrial effluents.

Therefore, this study aims to investigate the solar-activated hypochlorite degradation of Rhodamine B (RhB) through detailed kinetic modeling, providing insights into the roles of reactive oxygen species (ROS) and reactive chlorine species (RCS) in the degradation process. The kinetic modeling facilitated the determination of key rate constants for both radical and non-radical pathways of hypochlorite and the oxidation of RhB by free radicals. Using COPASI^® software, which was fed with a kinetic mechanism involving 144 chemical reactions, the free radical kinetic model was tested under various conditions, including temperatures ranging from 25 to 55 °C and initial hypochlorite concentrations from 300 to 1000 µM, at a controlled pH of 11. RhB was chosen as the pollutant model for several reasons: (i) it is a common dye pollutant whose degradation has been widely studied in the literature [46,47,48], (ii) the reaction rate constants of RhB with ^•OH and O₃ are available, reducing the number of unknown parameters in the reaction kinetics model, and (iii) its reaction with chlorine is negligible in basic medium where ClO⁻ predominates, allowing for a purely free radical reaction mechanism.

2. Experimental and Modeling Package

2.1. Reagents and Setup Specifications

Rhodamine B (RhB; C₂₈H₃₁N₂O₃Cl; MW: 479.01 g/mol) was sourced from Sigma-Aldrich. NaOCl solution with approximately 16% accessible chlorine was supplied by Henkel-Algeria. Tert-butanol, benzoic acid, NaOH, and H₂SO₄ were purchased from Sigma-Aldrich and were of the highest purity grade available. All reagent solutions were prepared in distilled water.

Reference [49] contains comprehensive details on the experimental setup, including specifications and descriptions. A 250 mL Pyrex glass reactor with a water jacket was part of the arrangement, and it was housed within a Suntest CPS+ simulator from Atlas. This simulator ran at 500 W of radiation intensity using a Xe arc lamp whose emission range was limited to 280–800 nm. About 0.5% of the photons released by the lamp were in the wavelength range of less than 300 nm, and about 7% were in the range of 300 to 400 nm. The emission spectrum from 400 to 800 nm was similar to that of sunlight. The solution’s surface and the light source were maintained at a constant distance. A thermocouple submerged in the reaction mixture was used to measure the temperature during the experiments, which were carried out between 25 and 55 °C. Water from a controlling bath (RC6 Lauda) was circulated through the reactor jacket to regulate the temperature. Using a Jenway 3505 pH meter, the pH of the solution was monitored during the process. A magnetic stirrer set at 300 rpm was used to constantly mix the 200 mL of irradiation RhB solution. Samples were taken periodically to measure the dye absorbance at λ_max = 551 nm using a UV-Visible spectrophotometer (Jasco V-730). Each experiment was conducted in triplicate, and the mean values, along with error bars, were reported in the figures.

2.2. Kinetic Model, Simulator, and Computational Approach

Table 1 summarizes the 144 chemical reactions used to study the degradation kinetics of RhB. This scheme includes the initial constituents (HOCl/ClO⁻, RhB, H₂O, O₂, H⁺ and OH⁻), various reactive radicals (Cl^•, ClO^•, ^•OH, ClOH^•−, Cl₂^•−, HO₂^•, O₂^•−, O₃^•−), several non-radical intermediates/products (O₃, ClO₂⁻, ClO₃⁻, H₂O₂, HO₂⁻, and Cl₂O₂), and the reactions of free radicals and oxidants with RhB and scavengers such as tert-butyl alcohol (TBA) or benzoic acid (BA). The chemical reactions and their rate constants were sourced from various specialized references (as noted in Table 1), except for the rate constants for Reactions R1-R3, R125 and R127-R133, which were determined in this study.

The kinetic modeling of RhB degradation, based on the reactions outlined in Table 1, was carried out using the open-source COPASI^® kinetic simulation software. This robust tool enables the determination of concentration-time profiles for all species involved in the reaction scheme. It simultaneously optimizes multiple unknown reaction rate constants to achieve the best fit for the degradation curve of the micropollutant [45,50]. In this study, the Nelder-Mead Simplex deterministic method, provided by the software, was selected for optimizing the unknown reaction rate constants (R1–R3, R125 and R127–R133).

The software’s input parameters comprise the starting concentrations of the following species: ClO⁻, RhB, O₂, H₂O, OH⁻, H⁺, and scavengers (TBA or BA), as well as the reaction scheme from Table 1 and their corresponding rate constants. Following that, the simulator generates concentration profiles for every species over the course of the oxidation period. The experimental degradation profile of RhB over time was fed into the software in order to derive the unknown rate constants, which included those for ClO⁻ photolysis (Reactions R1–R3) and the reactions of RhB with other reactive species (Reactions R125 and R127–R133). This profile was adjusted by optimizing the unknown rate constants as parameters.

2.3. Reactive Species Contribution

The contribution of each reactive species (RS: ^•OH, O^•−, O(³P), Cl^•, ClO^•, Cl₂^•−, O₃, etc.) to the overall degradation rate of the micropollutant was quantified by calculating their selectivity. This is given by:

$$S_{RC/RhB}={\displaystyle \frac{\left(r_{RC\text{-}prod}\right)}{\left({-r}_{RhB}\right)}}$$

(1)

Here, (−r_RhB) is the rate at which RhB disappears, and (r_RC-prod) is the rate at which the organic product forms due to the direct reaction of each specific reactive species (RC) with RhB. Both rates were quantified (from the optimized profiles) during the initial progress of the reaction. The selectivity approach enables an assessment of the impact of each reactive species on the degradation process, providing a clear understanding of their relative contributions during the early stages of the reaction.

3. Results and Discussion

3.1. Validated RhB Degradation Kinetics Using Scavenger Tests

Figure 1a–c compare the experimental and simulated (using the model from Table 1) photodegradation profiles of RhB (C₀ = 10 µM) at pH 11 in the presence of ClO⁻ (1000 µM) and either tert-butyl alcohol (TBA: 100 mM) or benzoic acid (BA: 10 mM), which act as specific radical scavengers. As shown in Figure 1a, the degradation of RhB rapidly diminished (exponentially) over time in the UV/ClO⁻ system. In contrast, no change in dye concentration was observed when subjected to either chlorination (1000 µM) or solar lighting separately. This demonstrates that the combined action of chlorine and solar illumination in the Sun Test model results in a fully synergistic treatment. The free radical-oxidation mechanism is the primary pathway responsible for the observed synergism in RhB degradation during the UV/ClO⁻ photo-assisted treatment at pH 11. This conclusion was supported by the addition of TBA and BA, which significantly slowed down RhB decomposition, as illustrated in Figure 1b,c.

Moreover, it is evident from Figure 1a–c that the kinetic model accurately matches the experimental degradation profiles, both in the absence and presence of TBA or BA. Notably, the reaction rate constant for BA with the ClO^• at a basic medium (k₁₄₁ in Table 1) is reported to be less than 3 × 10⁶ M⁻¹s⁻¹ [42,51] (this reaction rate constant is much lower (7.26 × 10³ [52] and 3.13 × 10³ M⁻¹ s⁻¹ [53]) at an acidic pH). Therefore, the fitting process in Figure 1c was conducted using three different values for k₁₄₁: 1 × 10⁵, 2 × 10⁵, and 1 × 10⁶ M⁻¹s⁻¹. Although all three values provided a good fit between the experimental and predicted profiles, the best among them was observed with k₁₄₁ set at 2 × 10⁵ M⁻¹s⁻¹ (Figure 1c).

The observed RhB degradation rate constant from the experimental profiles in the control run is k_obs = 3 × 10⁻³ s⁻¹ (R² = 0.9785). However, it decreased to k_BA = 8 × 10⁻⁴ s⁻¹ (R² = 0.9816) in the presence of BA and k_TBA =1 × 10⁻⁴ s⁻¹ (R² = 0.9305) in the presence of TBA. The predicted values from the modeling profiles closely matched the experimental ones, with 3 × 10⁻³ s⁻¹ for the control run (R² = 1), 9 × 10⁻⁴ s⁻¹ (R² = 0.9997) with BA, and 1 × 10⁻⁴ s⁻¹ (R² = 0.9933) with TBA. The predicted profiles in Figure 1a–c were obtained using optimized specific rate constants for several reactions listed in Table 1, a topic that will be discussed in the subsequent section. However, before delving into that, it is important to estimate the overall contribution of some reactive species in the dye degradation based on the experimental tendency of BA and TBA.

BA (i.e., benzoate form at basic pH) commonly quenches ^•OH/O^•− and Cl^• in a UV/chlorine photoinduced system, with rate constants k_BA-•OH = 5.27 × 10⁹ M⁻¹s⁻¹, k_BA-O•– = 4 × 10⁷ M⁻¹s⁻¹, and k_BA-Cl• = 1.8 × 10¹⁰ M⁻¹s⁻¹ [5]. However, its reactions with ClO^• and Cl₂^•− are relatively insignificant, with rate constants k_BA-ClO• < 3 × 10⁶ M⁻¹ s⁻¹ [42] and k_BA-Cl2•– = 2 × 10⁶ M⁻¹ s⁻¹, respectively [5]. Thus, the contribution of (^•OH/O^•− + Cl^•) can be estimated based on the retarding impact of BA by using the ratio (k_obs-k_BA)/k_obs, which is calculated as (0.003–0.0008)/0.003 = 0.73. Therefore, the contribution of (^•OH/O^•− + Cl^•) is approximately 73%.

Similarly, TBA quenches OH/O^•−, Cl^•, and ClO^•, with rate constants k_TBA-•OH = 3.80 × 10⁸ M⁻¹s⁻¹, k_TBA-O•- = 5 × 10⁸ M⁻¹s⁻¹, k_TBA-Cl• = 3 × 10⁸ M⁻¹s⁻¹, and k_{TBA- ClO•} = 1.30 × 10⁷ M⁻¹s⁻¹ [39,54]. However, TBA reacts insignificantly with Cl₂^•− and ozone, with rate constants k k_{TBA- Cl2•−} = 700 M⁻¹s⁻¹ [39,54] and k_{TBA- O3} = 0.03 M⁻¹ s⁻¹ [41]. The contribution of (^•OH/O^•− + Cl^• + ClO^•) can be estimated using the ratio (k_obs-k_TBA)/k_obs, which is calculated as (0.003–0.0001)/0.003 = 0.966 (96.6%). Thus, the contribution of ClO^• can be estimated as 96.6%–73% = 23.6%. The specific contributions of other radicals and oxidants are difficult to estimate experimentally due to the lack of specific quenchers, but these contributions will be predicted based on the best fitting provided by the model.

3.2. Determination of Unknown Rate Constants

The created free-radicals model correctly predicted RhB degradation profile under a variety of circumstances, as was covered in the previous section (Figure 1). The software’s Nelder-Mead optimization technique was used to identify the unknown rate constants by finding the best match to the experimental data (Figure 1a). At pH 11 and around 25 °C, the starting hypochlorite concentration of 1000 µM yields the following rate constants for the photolysis of ClO⁻ through various pathways: k_R1 = 2.67 × 10⁻⁴ s⁻¹ for R1: ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O^•− + Cl^•, k_R2 = 1.88 × 10⁻⁵ s⁻¹ for R2: ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O(³P) + Cl⁻ and k_R3 = 0 s⁻¹ for R3: ClO⁻ $\stackrel{h\nu }{\to }$ O(¹D) + Cl⁻. These results indicate that the radical pathway (R1) of ClO⁻ photolysis under solar light irradiation is more favorable than the non-radical pathways (R2 and R3), though this is specific to the current simulation conditions (1000 µM chlorine, pH 11 and 25 °C). This is because the activation pathway of chlorine can be influenced by various operating conditions, such as medium temperature, pH level, chlorine dose, and the presence of a catalyst. As a result, the following sections investigate these activation possibilities under different chlorine concentrations and solution temperatures to evaluate their significance within the overall mechanism, specifically the abatement of hypochlorite and the contributions of ROS (Reactive Oxygen Species) and RCS. It should be indicated that these rate constants (k_R1, k_R2, and k_R3) align well with the quantum yields determined by Buxton and Subhani [55] at 313 nm: Φ (R1) = 0.127, Φ (R2) = 0.075 and Φ (R3) = 0.02. At 356 nm, the quantum yield for R3 is effectively zero [55]. It is important to note that the photolysis of H₂O₂, HO₂⁻, and O₃ are not included in the reaction scheme because they are produced in lower concentrations and their molar absorptivity is negligible between 300 and 400 nm [5]. Additionally, at pH 11, ClO⁻ is the sole chlorine species in the solution, so photolytic reactions of hypochlorous acid are not significant. At 254 nm irradiation light (pH 9), Djaballah et al. [45] reported a rate constant of ~9 × 10⁻³ s⁻¹ for R2. The discrepancy between our value and Djaballah et al.’s was attributed to the higher energy (E = hC/λ) of the 254 nm irradiation light compared to our solar system, which includes only about 0.5% UV light < 300 nm and 7% between 300 and 400 nm (with lower intensity). Additionally, the Φ₂₅₄ (R2) = 0.278 [15] is much higher than the Φ₃₆₅ (R₂) = 0.08 [15]. The single rate constants for the photolysis of S₂O₈²⁻ and H₂O₂ at 254 nm irradiation was found to be 1.5 × 10⁻⁵ s⁻¹ [56] and 1.16 × 10⁻⁵ s⁻¹ [57]. As a result, our fitted values (both R1 and R2) for hypochlorite photolysis are higher than those for H₂O₂ and S₂O₈²⁻. This makes sense, because the molar absorption coefficient of ClO⁻ is higher than that of H₂O₂ and S₂O₈²⁻ (at 254 nm, ε_ClO⁻ = 66.0 M⁻¹ cm⁻¹ [58] compared to 19.0 M⁻¹ cm⁻¹ for H₂O₂ and 47.50 M⁻¹ cm⁻¹ for S₂O₈²⁻ [43]).

The rate constants for the reactions of RhB with ^•OH and O₃ have been experimentally determined to be k_R124 = 2.5 × 10¹⁰ M⁻¹ s⁻¹ [33] and k_R126 = 2450 M⁻¹ s⁻¹ [38], respectively. In contrast, the rate constants for the reactions of RhB with O^•−, Cl^•, Cl₂^•− and ClO^• were predicted to be k_R125 = 4.8 × 10⁹ M⁻¹ s⁻¹, k_R127 = 1.45 × 10⁹ M⁻¹ s⁻¹, k₁₂₉ = 2.5 × 10⁷ M⁻¹ s⁻¹ and k₁₂₈ = 8.7 × 10⁴ M⁻¹ s⁻¹ (Table 1). Additionally, lower rate constants were predicted for the reactions of RhB with HOCl^•−, HO₂^•, O₂^•−, and O(³P), with k_R130 = 4.1 × 10⁴ M⁻¹ s⁻¹, k_R131 = 7.3 × 10⁵ M⁻¹ s⁻¹, k_R132 = 3.6 × 10⁴ M⁻¹ s⁻¹, and k_R133 = 0.40 M⁻¹ s⁻¹, respectively (Table 1). These predicted values align with those reported for various organic pollutants, including dyes. The reactivity of O(³P) with chlorazol black was found to be negligible [45], while the reactions of HOCl^•−, HO₂^•, and O₂^•− with dyes are often negligible as well [30,45,59]. Measurements show that Cl^• is highly reactive towards organic aromatic solutes, such as toluene (k =1.80 × 10¹⁰ M⁻¹ s⁻¹), chlorobenzene (k = 1.80 × 10¹⁰ M⁻¹ s⁻¹) and benzoic acid (k = 1.80 × 10¹⁰ M⁻¹ s⁻¹). Furthermore, Cl₂^•− reacts with acid orange 7 azo at k = 3.65 × 10⁷ M⁻¹ s⁻¹ [56]; however, other values such as k = 2.50 × 10⁸, 2.80 × 10⁸ and 3.0 × 10⁶ M⁻¹ s⁻¹ were reported for phenol, p-hydroxybenzoic acid, and p-chlorobenzoic acid, respectively [15].

It is important to note that in the subsequent sections when the effects of solution temperature and initial concentrations of chlorine and RhB are examined, only the rate constants k_R1 and k_R2 (corresponding to the radical and non-radical photolysis pathways of ClO⁻) will be optimized based on the experimental profiles. The rate constants for RhB’s reactions with radicals (Reactions R125 and R127–R133 in Table 1) will remain unchanged.

3.3. Concentration Profiles of ClO⁻, Reactive Species and Degradation Products

3.3.1. Control Conditions (without Scavengers)

Figure 2 presents the simulated concentration profiles of the various reactants, reactive species, and reaction products in the RhB-hypochlorite photoactivated (solar) system, using the same fitting conditions as in Figure 1a. As depicted in Figure 2a–d, the depletion of ClO⁻ and RhB leads to the substantial formation of Cl₂OH⁻ (Figure 2a), a transient increase in reactive oxygen and chlorine species (ROS/RCS) (Figure 2b,c), and a high level of RhB degradation by-products (Figure 2d). The concentration of Cl₂OH⁻ increases over time, becoming the predominant hypochlorite photolysis by-product, while HOCl is initially produced at a lower yield, reaching a steady concentration of 1.89 × 10⁻⁵ M before being quickly consumed. After 100 s, Cl₂OH⁻ ions attain a concentration of 5.97 × 10⁻⁵ M, compared to 1.58 × 10⁻⁵ M for HOCl (Figure 2a). In contrast, other species such as ClO₂⁻, ClO₃⁻, H₂O₂, HCl, Cl⁻, and Cl₂O₂ are produced in negligible concentrations (Figure 2a).

Figure 2c,d illustrate the presence of reactive oxygen species (ROS) and reactive chlorine species (RCS) in the system. The ROS include O₃, ^●OH, O^●−, and O(³P), while the RCS are ClO^●, Cl^●, and Cl₂^●−. Other ROS and RCS, such as O₂^●−, HO₂^●, Cl₂, ClO₂^●, and HOCl^●−, were generated in negligible amounts. As shown in Figure 2c,d, the steady-state concentrations of ^●OH, O^●−, and Cl^● reached 7.54 × 10⁻¹⁴, 1.22 × 10⁻¹⁴, and 1.38 × 10⁻¹⁴ M, respectively, with ratios [^●OH]_ss/[O^●−]_ss = 6.18 and [^●OH]_ss/[Cl^●]_ss = 5.45. Cl₂^●− and O(³P) achieved lower steady-state concentrations of 3.69 × 10⁻¹⁶ and 1.73 × 10⁻¹⁵ M, respectively. In contrast, ClO^● and O₃ were produced in significantly higher concentrations, with [ClO^●]_ss = 9.95 × 10⁻⁹ M and [O₃]_ss = 4.96 × 10⁻⁹ M, resulting in [ClO^●]_ss/[Cl^●]_ss and [O₃]_ss/[^●OH]_ss ratios of 7.19 × 10⁵ and 6.59 × 10⁴, respectively. The high concentration of ClO^● can be attributed to the rapid reaction of initially formed radicals (^●OH and Cl^●) with hypochlorite, which is in excess at the early stages of the reaction, as described by Reactions R30 and R44 in Table 1. The significant formation of O₃ is due to the high yield of O(³P) from reaction R2 (ClO⁻ → O(³P) + Cl⁻, k_R2 = 1.88 × 10⁻⁵ s⁻¹) and the subsequent fast reaction of O(³P) with dissolved O₂ to produce O₃ (R5, k_R5 = 4 × 10⁹ M⁻¹ s⁻¹). Notably, O₃ concentration continues to rise over time, reaching 1.31 × 10⁻⁸ M at t = 800 s, owing to the ongoing release of O(³P) via reaction R2, as hypochlorite levels remain high even at t = 800 s (Figure 2a). It is essential to mention that the high levels of O₃ and ClO^● do not necessarily indicate their dominant role in RhB degradation, as the degradation rate is determined by the concentrations of reactive species (RS) and RhB, as well as their reaction rate constants, r_i = k_RS-RhB[RS][RhB], which will be discussed later.

From Figure 2b, it can be observed that the concentration profiles of ^●OH-prod and ClO^●-prod increased over time as primary RhB by-products, with a slight presence of O^●−-prod also noted. No other RC-RhB by-products were formed. The initial stage of these profiles is of particular interest, as by-products become competitors for radical consumption after prolonged irradiation (the reaction model based on Table 1 did not account for this scenario). At the early stages of the reaction (before attaining t_1/2 [50]-), the reactions between the target dye (RhB) and radicals predominate (by-product levels remain low), thus enabling the assessment of each radical’s contribution to RhB degradation using Equation (1). This equation yielded contributions of 67%, 2.1%, and 29.6% for ^●OH, O^●−, and ClO^●, respectively, in the overall degradation rate of RhB in the chlorine-photoassisted process. The contributions of all other reactive species (including O₃, Cl^●, Cl₂^●−, HO₂^●, O₂^●−, O(³P), and ClOH^●−) under the given simulation conditions are negligible, although their contributions may change with operational conditions, such as the initial dosage of hypochlorite, as will be discussed in the subsequent sections.

In Section 3.1, the final two paragraphs discuss probing tests with TBA and BA, which indicate an estimated contribution of 23.6% for ClO^● and 73% for the combined effects of (^●OH/O^●− + Cl^●). Although these values align with the predictions made in this section (67% for ^●OH, 2.1% for O^●−, and 29.6% for ClO^●), the numerical simulations provide a more precise determination of the specific contributions of each radical species, highlighting no contribution from Cl^● and a dominant role for ^●OH over O^●−. Despite the photolysis being conducted at pH 11, where the pK_a of ^●OH is 11.8, the hydroxyl radical speciation still favors the ^●OH form, with an 86.3% dominance over the 13.6% presence of O^●−.

Regardless, as previously mentioned, the RhB degradation rate is determined by the concentrations of reactive species (RS) and RhB, as well as their reaction rate constants, r_i = k_RS-RhB[RS][RhB]. Therefore, even though O₃ and ClO^● have higher concentrations in the solution, their contributions are lower than that of ^●OH, primarily due to the significantly higher rate constant between RhB and ^●OH (2.5 × 10¹⁰ M⁻¹ s⁻¹) compared to those of RhB with O₃ (2450 M⁻¹ s⁻¹) and ClO^● (8.70 × 10⁴ M⁻¹ s⁻¹). A similar interpretation applies to the null contributions of Cl^●, Cl₂^●−, HOCl^●−, HO₂^●, O₂^●−, and O(³P), considering both the lower concentrations of these radicals (Figure 2b,c) and their much lower rate constants with RhB (R127, R129-R132 of Table 1) compared to those related to ^●OH.

3.3.2. UV_solar/ClO⁻/TBA and UV_solar/ClO⁻/BA Systems

Figure 3 and Figure 4 illustrate the simulated concentration profiles of various reactants, reactive species, and reaction products in the RhB-hypochlorite photoactivated system when TBA and BA are present, respectively (under the same conditions as in Figure 1b,c). The addition of TBA and BA resulted in a significant decrease in the steady-state concentration of ^•OH to 5.65 × 10⁻¹⁵ M and 2.89 × 10⁻¹⁴ M, representing reductions of 92.5% and 61.6%, respectively. The corresponding quenching effects on the key species in the reaction system are depicted in Figure 5. The quenching data provided in Figure 5 reveal substantial insights into the effects of TBA and BA on various reactive species within the RhB-hypochlorite photolysis system. The significant quenching percentages illustrate the direct or indirect interactions of these scavengers on different species. For the hydroxyl radical (^•OH), TBA exhibited a quenching effect of 92.50%, highlighting its high reactivity with hydroxyl radicals and effectiveness as a scavenger. In comparison, BA also significantly quenched ^•OH, but to a lesser extent at 61.59%, indicating that while effective (k_R143 = 5.27 × 10⁹ M⁻¹s⁻¹), it is not as potent as TBA (k_R137 = 3.80 × 10⁸ M⁻¹s⁻¹). This is due to the fact that we used a higher initial concentration of TBA (100 mM) compared to BA (1 mM) for solubility reasons.

The superoxide radical anion (O^•−) showed a similar trend, with TBA quenching 81.75% of O^•−, thus demonstrating TBA’s substantial impact. BA, on the other hand, reduced O^•− concentration by 48.41%, which is still significant but comparatively lower than the effect of TBA. For ozone (O₃), TBA caused a 12.28% reduction, suggesting that TBA slightly affects ozone formation (k_R141 = 0.003 M⁻¹s⁻¹). Conversely, BA exhibited null quenching due to the negligible reaction between O₃ and BA. Interestingly, both TBA and BA had no effect on ground-state atomic oxygen (O(³P)), with 0% quenching for both. This implies that neither TBA nor BA interacts significantly with O(³P) and that this species did not play a major role in the reactive system under study.

Chlorine radical (Cl^•) quenching was substantial with both TBA and BA. TBA reduced Cl^• concentration by 72.09%, while BA achieved a slightly lower quenching effect of 63.57%, indicating strong but differential interactions with Cl^•. For the chlorine oxide radical (ClO^•), TBA was extremely effective, achieving 100% quenching, thus completely scavenging this species due to the effective quenching of Cl^• and ^•OH, which are precursors of ClO^• (R33 and R44 of Table 1). BA also had a strong effect, with an 86.76% quenching rate, though it was slightly less effective than TBA. The quenching of the chlorine dioxide radical anion (Cl₂^•−) was also high, with TBA achieving 91.46% and BA 85.99%, indicating that both agents effectively reduce the concentration of this radical. This is attributed to the efficient scavenging effect of both TBA and BA on the precursor of Cl₂^•−, which is Cl^• (R134 and R140 in Table 1).

The hypochlorite ion (ClO¯) experienced a moderate quenching effect from TBA at 42.11%, while BA’s effect was lower at 23.89%, suggesting a less pronounced but still significant interaction. This is related to the direct impact of TBA and BA on quenching reacting radicals with ClO¯, as exemplified in reactions R134 and R140 in Table 1. Finally, for the dichlorine monoxide anion (Cl₂OH¯), TBA caused a 68.53% reduction, while BA resulted in a 60.21% reduction, highlighting their substantial but differential impacts on this species.

3.4. Analysis of Initial Hypochlorite Dosage

3.4.1. RhB and Hypochlorite Decay

In Figure 6a, the simulated RhB concentration profiles are compared with the experimental profiles at three initial hypochlorite concentrations (300, 500, and 1000 µM) under the conditions of pH 11 and a temperature of 25 °C. The comparison reveals that both the experimental and simulated results agree in showing that the degradation rate of RhB increases with higher initial chlorine concentrations. There is a strong concordance between the experimental data and the modeling profiles. This trend of increased degradation with higher chlorine concentrations has also been observed in various experimental studies for different contaminants [60,61,62]. The RhB degradation rate constants (pseudo-first-order) increased to 2 × 10⁻⁴, 10⁻³ and 3 × 10⁻³ s⁻¹ for 300, 500 and 1000 µM [ClO⁻]₀, respectively. Correspondingly, the initial rates of degradation of RhB were increased to 4.07 × 10⁻⁹, 8.85 × 10⁻⁹, and 2.07 × 10⁻⁸ M s⁻¹.

In Figure 6b, the kinetics of hypochlorite depletion and Cl₂OH⁻ formation (the main product) are illustrated under the optimal fitting conditions from Figure 6a. Both rates increase with higher hypochlorite dosages, elucidating the behavior of RhB degradation observed in Figure 6a. A greater hypochlorite depletion leads to a higher yield of radical generation and an increased rate of Cl₂OH⁻ release, as demonstrated. According to Figure 6b, hypochlorite depletes at increasing initial rates of 8.84 × 10⁻⁸, 3.22 × 10⁻⁷, and 1.15 × 10⁻⁶ M s⁻¹, while Cl₂OH⁻ forms at rates of 1.92 × 10⁻⁸, 1.19 × 10⁻⁷, and 4.97 × 10⁻⁶ M s⁻¹ as the initial hypochlorite dosage rises from 300 µM to 500 µM and 1000 µM, respectively. Consequently, the yield of Cl₂OH⁻ formation from hypochlorite depletion increases to 21.7%, 37%, and 43.2% for [ClO⁻]₀ values of 300, 500, and 1000 µM, respectively.

In Figure 6c, the overall hypochlorite depletion rate constant (derived from the profiles in Figure 6b) and the specific ClO⁻ photolysis rate constants (k_R1 and k_R2, predicted based on the best fit in Figure 6a) are plotted against initial hypochlorite concentration [ClO⁻]₀. The k_R1 increases linearly, from 9.05 × 10⁻⁶ s⁻¹ at 300 µM hypochlorite to 2.67 × 10⁻⁴ s⁻¹ at 1000 µM hypochlorite. In contrast, k_R2 shows less variation, being 3 × 10⁻⁵ s⁻¹ at 300 µM, 5.54 × 10⁻⁵ s⁻¹ at 500 µM, and 1.88 × 10⁻⁵ s⁻¹ at 1000 µM. Additionally, Figure 6c demonstrates that the overall first-order rate constant (k_overall) for chlorine depletion is greater than the sum of the specific rate constants, k_R1 + k_R2. The ratio (k_R1 + k_R2)/k_overall is 0.195, 0.277, and 0.286, respectively, at the concentrations indicated, indicating that the combined contribution of Reactions R1 and R2 in hypochlorite depletion does not exceed 30%. This finding suggests that hypochlorite consumption also occurs via other pathways besides Reactions R1 and R2. Indeed, radicals such as O^•−, Cl^•, O₃, and ^•OH, generated from Reactions R1, R2, and R115, can break down hypochlorite at higher rate constants according to Reactions R30, R44, R71, and R114 of Table 1.

3.4.2. ROS and RCS Evolutions

Figure 7 illustrates the simulated profiles of the main ROS and RCS under the conditions set in Figure 6a. As the hypochlorite concentration increases, the concentration of RCS (Cl^•, ClO^•, and Cl₂^•−) and ROS (^•OH/O^•−) also increases. The steady-state concentration of ^•OH rises from 4.3 × 10⁻⁵ M at 300 µM to 2.76 × 10⁻¹⁴ M at 500 µM, and further to 7.54 × 10⁻¹⁴ M at 1000 µM. Similarly, O^•− shows an increase, albeit at a lower yield, with concentrations of 5.68 × 10⁻⁶ M at 300 µM, 3.9 × 10⁻¹⁵ M at 500 µM, and 1.22 × 10⁻¹⁴ M at 1000 µM. Conversely, the concentration of O₃ peaks at 500 µM with [O₃]_ss = 1.56 × 10⁻⁸ M, a trend corresponding to the optimal behavior of its precursor O(³P) and the decline in the photolytic rate constant of reaction R₂ (k_R2) beyond 500 µM, as shown in Figure 6c. For the same range of hypochlorite concentrations, Cl^• radicals reach steady-state values of 3.23 × 10⁻¹⁶ M, 2.9 × 10⁻¹⁵ M, and 1.38 × 10⁻¹⁴ M. ClO^• radicals, however, achieve higher steady-state concentrations compared to ^•OH and Cl^•, with values of 8.67 × 10⁻¹⁰ M, 3.8 × 10⁻⁹ M, and 9.95 × 10⁻⁹ M for 300, 500, and 1000 µM of [ClO⁻]₀, respectively. Similarly, Cl₂^•− radicals follow the same trend, although they are the least formed RCS.

Overall, except for O₃, the yield of reactive species increased with rising hypochlorite concentration. This explains the intensified effect of higher hypochlorite concentrations on the degradation rate of RhB, as observed in Figure 6a. These findings suggest that optimizing hypochlorite dosage is essential to maximize the degradation efficiency of RhB, considering both the generation of reactive species and their interaction dynamics.

3.4.3. Products and Reactive Species Contribution

Figure 8 presents the simulated profiles of the main considered degradation products of RhB, namely Prod_^•OH, Prod_O^•−, Prod_O₃, Prod-Cl^•, Prod-ClO^•, and Prod-Cl₂^•−, under the conditions described in Figure 6a. Prod_^•OH and Prod-ClO^• are the primary products, with their concentrations increasing over time and with higher hypochlorite dosages. Prod_O^•− and Prod_O₃ are formed to a lesser extent, while Prod-Cl^• and Prod-Cl₂^•− are not detected. These findings indicate that ^•OH and ClO^• radicals predominantly attack RhB. The contribution of each radical to the overall RhB degradation rate is calculated using the selectivity equation (Equation (1)). The terms of Equation (1) are calculated, and the results for three initial hypochlorite concentrations (300, 500, and 1000 µM) are shown in

Table 2. Computed RhB degradation rates, product formation rates, and the subsequent contribution of each reactive species to the overall degradation rate of RhB for various initial hypochlorite concentrations. The contribution is calculated using the selectivity equation (Equation (1)).

Species	Rate (M/s)	Contribution (%)
	[ClO−]₀ = 1000 µM
RhB	2.89 × 10−6
Prod_Cl•	1.93 × 10−8	0.67
Prod_•OH	1.94 × 10−6	67.01
Prod_O•−	6.05 × 10−8	2.09
Prod_ClO•	8.58 × 10−7	29.63
Prod_O₃	1.75 × 10−8	0.60
	[ClO−]₀ = 500 µM
RhB	1.25 × 10−6
Prod_Cl•	4.61 × 10−9	0.37
Prod_•OH	7.84 × 10−7	62.78
Prod_O•−	2.12 × 10−8	1.70
Prod_ClO•	3.69 × 10−7	29.54
Prod_O₃	7.00 × 10−8	5.61
	[ClO−]₀ = 300 µM
RhB	2.70 × 10−7
Prod_Cl•	5.60 × 10−10	0.21
Prod_•OH	1.31 × 10−7	48.43
Prod_O•−	3.32 × 10−9	1.23
Prod_ClO•	9.16 × 10−8	33.93
Prod_O₃	4.38 × 10−8	16.21

. According to the results, ^•OH is the dominant contributor to RhB degradation, accounting for 67% at 1000 µM, 62.78% at 500 µM, and 48.43% at 300 µM. ClO^• is the second most significant contributor, with around 29.5% at both 1000 and 500 µM and 33.93% at 300 µM. The contribution of ^•OH decreases with lower hypochlorite dosages, while ClO^• shows the opposite trend. The third contributing species, although less significant than the previous two, is O₃, contributing 16.21% at 300 µM and 5.61% at 500 µM, but its contribution is negligible (0.6%) at 1000 µM. Therefore, the ozonation pathway increases with decreasing hypochlorite dosage, whereas the free radical attack (involving ^•OH and ClO^•) becomes more significant at higher ClO⁻ concentrations. It is important to note that the evolution of product concentrations depends on two kinetic parameters: the rate constants and the reactive species concentration, as previously discussed. The interaction (multiplication) between these parameters determines the overall contribution of each reactive species in the dye degradation process.

3.5. Analysis of Solution Temperature Impact

In Figure 9a, the simulated RhB concentration profiles are compared with experimental profiles at four different temperatures (25, 35, 45, and 55 °C) under conditions of pH 11 and using 1000 µM of hypochlorite. The comparison shows that both experimental and simulated results agree, indicating that the RhB degradation rate increases with temperature. There is a strong concordance between the experimental data and the modeling profiles. The pseudo-first-order RhB degradation rate constants increased to 2 × 10⁻³ and 3 × 10⁻³ s⁻¹ for 45 and 55 °C, compared to 1 × 10⁻³ s⁻¹ at 25 °C. Correspondingly, the initial degradation rates increased to 4.58 × 10⁻⁶ and 5.87 × 10⁻⁶ M s⁻¹ at 45 and 55 °C, compared to 22.89 × 10⁻⁶ M s⁻¹ at 25 °C.

In Figure 9b, the effect of liquid temperature on the kinetics of hypochlorite depletion and Cl₂OH⁻ formation (the main product) is illustrated under the best-fitting conditions from Figure 6a. Both rates increase with higher temperature, elucidating the behavior of RhB degradation observed in Figure 9a. Greater hypochlorite depletion in heated solutions leads to a higher yield of radical generation and an increased rate of Cl₂OH⁻ release, as demonstrated. According to Figure 9b, hypochlorite depletes at increasing initial rates of 1.36 × 10⁻⁶, 2.0 × 10⁻⁶, and 2.77 × 10⁻⁶, while Cl₂OH⁻ forms at rates of 6.02 × 10⁻⁷, 9.29 × 10⁻⁷, and 1.31 × 10⁻⁶ M s⁻¹ as the temperature increases to 35, 45, and 55 °C, respectively. Consequently, the yield of Cl₂OH⁻ formation from hypochlorite depletion increases from 43.2% at 25 °C to 44.3%, 46.0%, and 47.1% at 35, 45, and 55 °C, respectively.

In Figure 9c, the impact of liquid temperature on the overall hypochlorite depletion rate constant (derived from the profiles in Figure 9b) and the specific ClO⁻ photolysis rate constants (k_R1 and k_R2, estimated based on the best fit in Figure 9a) is shown. Both k_R1 and k_R2 increase with liquid temperature, with k_R1 exhibiting a more substantial rise compared to k_R2 (more than 10-fold). Additionally, the overall first-order rate constant (k_overall) for hypochlorite depletion is greater than the sum of the specific rate constants, k_R1 + k_R2. The ratio (k_R1 + k_R2)/k_overall remains around 0.28–0.30 (Figure 9c) across the investigated temperature range, indicating that the combined contribution of Reactions R1 and R2 to hypochlorite depletion does not exceed 30%, similar to the findings for varying initial hypochlorite dosages. Consequently, hypochlorite dissociation also occurs through other pathways besides Reactions R1 and R2, as previously discussed. Indeed, radicals such as O^•−, Cl^•, O₃, and ^•OH, generated from Reactions R1, R5, and R115, can degrade hypochlorite at higher rate constants according to Reactions R302, R44, R71, and R14 in Table 1. The efficiency of these reactions becomes significantly more pronounced at elevated temperatures, particularly at 55 °C. This is evident as the k_overall for hypochlorite depletion peaks against the sum of k_R1 and k_R2 at this temperature (Figure 9c).

Figure 10 depicts the simulated profiles of the main ROS and RCS under the conditions outlined in Figure 9a. An increase in liquid temperature had a favorable effect on both ROS and RCS, with their concentrations rising with temperature. For instance, the steady-state concentration of ^•OH increased from 7.54 × 10⁻¹⁴ M at 25 °C to 9.28 × 10⁻¹⁴ M at 35 °C, 1.47 × 10⁻¹³ M at 45 °C, and 2.22 × 10⁻¹³ M at 55 °C. Similarly, ClO^• rose from 99.95 × 10⁻⁹ M at 25 °C to 1.1 × 10⁻⁸ M at 35 °C, 1.38 × 10⁻⁸ M at 45 °C, and 1.7 × 10⁻⁸ M at 55 °C. Cl^• increased from 1.38 × 10⁻¹⁴ M at 25 °C to 1.7 × 10⁻¹⁴ M at 35 °C, 2.7 × 10⁻¹⁴ M at 45 °C, and 4.06 × 10⁻¹⁴ M at 55 °C. Overall, the yield of reactive species increased with rising liquid temperature. This accounts for the intensified effect of solution heating on the degradation rate of RhB observed in Figure 9a. These findings indicate that controlling liquid temperature is crucial for maximizing the degradation efficiency of RhB, as it influences the generation of reactive species.

Figure 11 displays the simulated profiles of the primary RhB degradation products, specifically Prod_^•OH, Prod_O^•−, Prod_O₃, Prod-Cl^•, Prod-ClO^•, and Prod-Cl₂^•−, under the conditions outlined in Figure 9a. Prod_^•OH and Prod-ClO^• emerge as the main products, with their concentrations increasing over time and with temperature. All other products are formed at negligible concentrations. The contribution of each radical to the overall RhB degradation rate is determined using the selectivity equation (Equation (1)). The terms of Equation (1) are calculated, and the results for various liquid temperatures are presented in

Table 3. Computed RhB degradation rates, product formation rates, and the subsequent contribution of each reactive species to the overall degradation rate of RhB for various liquid temperatures. The contribution is calculated using the selectivity equation (Equation (1)).

Species	Rate (M/s)	Contribution (%)
	Temp. 25 °C
RhB	2.89 × 10−6
Prod_Cl•	1.93 × 10−8	0.67
Prod_•OH	1.94 × 10−6	67.01
Prod_O•−	6.05 × 10−8	2.09
Prod_ClO•	8.58 × 10−7	29.63
Prod_O₃	1.75 × 10−8	0.60
	Temp. 35 °C
RhB	3.35 × 10−6
Prod_Cl•	2.27 × 10−8	0.68
Prod_•OH	2.32 × 10−6	69.25
Prod_O•−	7.25 × 10−8	2.16
Prod_ClO•	9.18 × 10−7	27.39
Prod_O₃	1.72 × 10−8	0.51
	Temp. 45 °C
RhB	4.58 × 10−6
Prod_Cl•	3.18 × 10−8	0.69
Prod_•OH	3.38 × 10−6	73.77
Prod_O•−	1.06 × 10−7	2.31
Prod_ClO•	1.03 × 10−6	22.50
Prod_O₃	3.33 × 10−8	0.73
	Temp. 55 °C
RhB	5.87 × 10−6
Prod_Cl•	4.07 × 10−8	0.69
Prod_•OH	4.55 × 10−6	77.54
Prod_O•−	1.43 × 10−7	2.44
Prod_ClO•	1.03 × 10−6	17.53
Prod_O₃	3.50 × 10−8	0.60

. According to these results, ^•OH is the dominant contributor to RhB degradation, accounting for 67% at 25 °C, 69% at 35 °C, 73.7% at 45 °C, and 77.5% at 55 °C. ClO^• is the second most significant contributor, with its contribution decreasing as the temperature rises: 29.63% at 25 °C, 27.4% at 35 °C, 22.5% at 45 °C, and 17.5% at 55 °C.

4. Conclusions

Important new information about the functions of various reactive species (RCS and ROS) in the degradation process has been revealed by modeling the oxidation kinetics of RhB in the UV_Solar/hypochlorite treatment. Using this method made it possible to determine important rate constants for the oxidation of RhB by free radicals and both radical and non-radical hypochlorite routes. Free radical concentrations and degradation kinetics are influenced by several operational parameters, including temperature and ClO⁻ initial dose, which may be quantitatively understood by kinetic modeling. It was discovered that RhB could be effectively degraded by solar-activated hypochlorite. The kinetics of this degradation were precisely modeled using a free radical kinetic model with COPASI^® software, under a range of conditions that included temperatures between 25 and 55 °C and initial hypochlorite concentrations between 300 and 1000 µM. To ensure ClO⁻ was the only form of chlorine present in the solution, all experiments were conducted at pH 11.

The validated kinetic model enabled the establishment of ROS and RCS profiles, highlighting their dependence on temperature and hypochlorite concentration. The simulated product profiles further identified the contributions of different reactive species to RhB degradation, with ^•OH emerging as the primary ROS and ClO^• as the main RCS. ^•OH displayed a dominant contribution over ClO^•, while ozone’s contribution was negligible but could reach up to 6% at lower hypochlorite dosages. Higher radical concentrations were achieved by increasing the solution temperature and initial hypochlorite dosage, resulting in more efficient hypochlorite decomposition and RhB degradation.

The photolysis rate constants for ClO⁻ via various pathways were predicted as follows: k_R1 = 2.67 × 10⁻⁴ s⁻¹ for R1: ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O^•− + Cl^•, k_R2 = 1.88 × 10⁻⁵ s⁻¹ for R2: ClO⁻ $\stackrel{\mathrm{h}\mathsf{\nu }}{\to }$ O(³P) + Cl⁻ and k_R3 = 0 s⁻¹ for R3: ClO⁻ $\stackrel{h\nu }{\to }$ O(¹D) + Cl⁻. The rate constants for RhB reactions with O^•−, Cl^•, Cl₂^•− and ClO^• were predicted to be k_R125 = 4.8 × 10⁹ M⁻¹ s⁻¹, k_R127 = 1.45 × 10⁹ M⁻¹ s⁻¹, k₁₂₉= 2.5 × 10⁷ M⁻¹ s⁻¹ and k₁₂₈ = 8.7 × 10⁴ M⁻¹ s⁻¹, respectively. Lower rate constants were predicted for RhB reactions with HOCl^•−, HO₂^•, O₂^•−, and O(³P), with k_R130 = 4.1 × 10⁴ M⁻¹ s⁻¹, k_R131 = 7.3 × 10⁵ M⁻¹ s⁻¹, k_R132 = 3.6 × 10⁴ M⁻¹ s⁻¹, and k_R133 = 0.40 M⁻¹ s⁻¹.

In summary, even though the model was only tested using the RhB decay profile, it shows impressive accuracy in a variety of demanding scenarios. First of all, every anticipated rate constant, including those for RhB interactions with free radicals and ClO⁻ photolysis, agrees with data from previously published studies on related pollutants. Second, the RhB degradation is well matched by these expected rate constants, under various hypochlorite dosages and temperatures. Finally, it correctly depicts the suppression of RhB degradation in the presence of TBA and BA scavengers. This extensive validation confirms the resilience of the model.