# Degradation of Reactive Blue 19 (RB19) by a Green Process Based on Peroxymonocarbonate Oxidation System

Authors: Thi Bich Viet Nguyen; Ngan Nguyen-Bich; Ngoc Duy Vu; Hien Ho Phuong; Hanh Nguyen Thi
Journal: Journal of Analytical Methods in Chemistry
 (2021)

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## Abstract

The effectiveness of peroxymonocarbonate (HCO4−) on the degradation of Reactive Blue 19 (RB19) textile dye was investigated in this study. The formation kinetics of HCO4− produced in situ in a H2O2−HCO3− system was studied to control the experimental conditions for the investigation of RB19 degradation at mild conditions. The effects of metallic ion catalysts, the pH, the input HCO3− and Co2+ concentrations, and UV irradiation were studied. The obtained result showed that Co2+ ion gave the highest efficiency on accelerating the rate of RB19 degradation by the H2O2–HCO3− system. In the pH range of 7–10, the higher pH values resulted in faster dye degradation. The reaction orders of the RB19 degradation with respect to Co2+ and HCO3– were determined to be 1.2 and 1.7, respectively. The UV irradiation remarkably enhanced the radical formation in the oxidation system, which led to high degradation efficiencies. The COD, TOC removal, and HPLC results clearly revealed complete mineralization of RB19 by the H2O2−HCO3−−Co2+ system.


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## Body

## 1. Introduction

Reactive dyes play an important role in textile industries. However, they may cause serious environmental problems as they are often nonbiodegradable and toxic to aquatic systems due to their complex aromatic molecular structures [1]. Once being discharged to the environment without destructive treatment, these water-soluble dyes can remain for a long time and alter the quality of water bodies by preventing light penetration and hindering photosynthesis, thereby affecting the ecosystems [2]. Among reactive dyes, anthraquinone dyes contain reactive groups with reinforced structures, which makes them difficult to degrade naturally and can be bioaccumulative in animals [3]. For example, an estimated half-life of C.I. Reactive Blue 19 at pH 7.0 and 25°C based on kinetic studies is ca. 46 years [4]. The concern over the ecotoxicity of these reactive dyes has led to the need to develop more efficient methods for their remediation from industrial effluents before discharging to water bodies. Textile wastewater treatment is not only decolorizing but also degrading and mineralizing dye molecules.Among technologies for dye treatment, advanced oxidation processes (AOPs) have attracted increasing attention due to their efficiency and ability to thoroughly remove pollutants from wastewater effluents. These AOPs base on the strong activity ofin situ formed free radicals, among which the most common radical is O•H with a redox potential of +2.8 V. Thanks to their strong oxidizing capacity, organic substances can be completely mineralized into CO2 and H2O. Some typical processes for active radical generation include Fenton, peroxone, H2O2/UV, O3/UV, and TiO2/UV.A photochemical-based AOP of textile wastewater containing RB19 dye has been studied using the UV/K2S2O8 system [5]. The results showed that only 50% of RB19 dye was removed after 5 hours in the dark system, while the UV-irradiated one gave a complete color removal in less than 30 min and 78.5% COD has been removed after 3 hours of irradiation time. The decolorization rate was fitted to the pseudo-first-order kinetic model regarding dye concentration. Another study on UV irradiation effect on RB19 decolorization conducted by Tehrani et al. who compared the RB19 decolorization efficiency by ozonation and UV-enhanced ozonation revealed that a UV irradiation by mercury lamp only increased COD removal efficiency, not the decolorization efficiency [6].In advanced oxidation processes, besidesO•H, O2−•, and O21 radicals, other oxidizing agents such as HCO4−, CO3−•, and HCO3• also play a crucial role in oxidation performance [7]. Peroxymonocarbonate (HCO4–) ion is known to be an active oxidant in the oxidation of organic compounds [8, 9]. This oxidizing agent was proved to have 100–500 times stronger reactivity in sulfide organic oxidation compared to hydrogen peroxide [10]. Although CO3−• and HCO3• radicals are reported to be less reactive than O•H, their occurrence at higher concentration and longer lifetime may provide sufficient oxidation potential. Another advantage of CO3−• and HCO3• radicals is a simple production technique.One of the major benefits of the oxidation process based on the peroxymonocarbonate system is its greenness toward the environment. As it can be prepared by mixingH2O2 with HCO3− or CO32−, the oxidation process does not require neither expensive nor toxic chemicals nor energy demand. Furthermore, it does not generate secondary sludge waste or require chemicals to adjust pH as compared to the Fenton process.The aim of this work was two‐fold: (i) to evaluate the effectiveness ofin situ generated HCO4− oxidizing agent in the HCO3−−H2O2 system in degrading the RB 19 textile dye under different experimental conditions and (ii) to verify the products of RB19 degradation by oxidation system based on HCO4−.


## 2. Materials and Experimental Procedure

### 2.1. Chemicals and Apparatus

Reactive Blue 19 dye was purchased from Sigma-Aldrich. Other chemicals and reagents used in this study were of analytical grade and used as received without any further purification.The RB19 degradations were monitored by measuring RB19 absorbance at 588 nm using a Biochrom Libra S60 UV-Vis Spectrophotometer. The degradation efficiency was calculated using the following equation: RB19 degradation (%) = (Co − Ct)/Co × 100%, where Co and Ct are the initial and remaining RB19 concentrations (ppm) at time t (min), respectively. These concentrations were determined using the standard curve Abs = (8.4 ± 0.1).103CRB19 (mg/L) (R2 = 0.999) with the LOD and LOQ of 0.7 and 2.3 mg/L, respectively. pH values were measured by a Lab850 pH meter (with BlueLine 14 pH electrode). COD values were determined by the oxidation standard method [11]. Total carbon and total inorganic carbon were measured on a multi-N/C 2100 TOC analyser (Analytik Jena AG). A Shimadzu RP-HPLC equipped with a PDA-M20A detector was used to investigate the products of the degradation. A mixture of acetonitrile and phosphate buffer at pH of 4.7 (50/50, v/v) was used as a mobile phase with a flow rate of 1 mL/min, an injection volume of 20 μL, and an oven temperature of 40°C.


### 2.2. Experimental Procedure

#### 2.2.1.In Situ Peroxymonocarbonate Preparation

Peroxymonocarbonate (HCO4−) solutions were prepared in situ by mixing sodium bicarbonate with hydrogen peroxide at a molar ratio of 1 : 2 at ambient temperature. Kinetics of HCO4− formation was evaluated by monitoring HCO4− concentration at different times in three hours. In this experiment, HCO4− was analyzed by a modified iodometric titration method [12]. In which, HCO4− was first allowed to react with iodide at low temperature (below −10°C) to prevent the effect of H2O2. The formed I2 was then titrated with thiosulfate using a starch indicator.


#### 2.2.2. Degradation of RB19 byIn Situ Generated HCO4− in the H2O2−HCO3− System

The effects of different parameters on the RB19 degradation were investigated as described in Table1 including the metallic ion catalyst (trials 1–9), the pH (trials 10–13), the HCO3− concentration (trials 14–18), the Co2+ catalyst concentration (trials 19–22), and UV irradiation (trials 24–28). The RB19 concentration was 100.0 mg/L in all trials. Each trial was replicated three times. In trials 1–5, 8–22, and 26, the mixture of HCO3− and H2O2 with a molar ratio of 1 : 2 was prepared in 50 minutes before adding to the RB19 solution to conduct the dye decolorization. For trials 1–23, the reactions were conducted in a 250 mL batch reactor controlled at 26 ± 1°C using a thermobath and thoroughly mixed by a magnetic stirrer as depicted in Figure 1(a). For trials 24–28, the reaction solutions were irradiated with the UVC light (254 nm, 12 W) in the UV chamber regulated by a thermobath as described in Figure 1(b). The flow speed through the UV chamber was maintained at 100 mL/min using a circulating pump. Dye decolorization was monitored by sampling at certain intervals and measuring light absorbance at the wavelength of 592 nm.Table 1
Conditions for RB19 decolorization by peroxymonocarbonate.
TrialpH[HCO3−] (mM)[H2O2] (mM)[M2+] catalyst (mg/L)UV irradiation181020—No2–5810200.1 (Ni2+, Mn2+, Zn2+, Co2+)No68—200—No78—2000.1 (Co2+)No88100200—No981002000.1 (Co2+)No10–137, 8, 9, 1010200.1 (Co2+)No14–1885, 10, 15, 25, 30400.1 (Co2+)No19–22820400.01, 0.02, 0.04, 0.06 (Co2+)No238———No248———Yes258—20—Yes268—200.1 (Co2+)Yes2781020—Yes28810200.1 (Co2+)YesFigure 1
The reaction system (a) and the continuous UV-irradiated system (b) for the RB19 degradation.
(a)(b)


## 2.1. Chemicals and Apparatus

Reactive Blue 19 dye was purchased from Sigma-Aldrich. Other chemicals and reagents used in this study were of analytical grade and used as received without any further purification.The RB19 degradations were monitored by measuring RB19 absorbance at 588 nm using a Biochrom Libra S60 UV-Vis Spectrophotometer. The degradation efficiency was calculated using the following equation: RB19 degradation (%) = (Co − Ct)/Co × 100%, where Co and Ct are the initial and remaining RB19 concentrations (ppm) at time t (min), respectively. These concentrations were determined using the standard curve Abs = (8.4 ± 0.1).103CRB19 (mg/L) (R2 = 0.999) with the LOD and LOQ of 0.7 and 2.3 mg/L, respectively. pH values were measured by a Lab850 pH meter (with BlueLine 14 pH electrode). COD values were determined by the oxidation standard method [11]. Total carbon and total inorganic carbon were measured on a multi-N/C 2100 TOC analyser (Analytik Jena AG). A Shimadzu RP-HPLC equipped with a PDA-M20A detector was used to investigate the products of the degradation. A mixture of acetonitrile and phosphate buffer at pH of 4.7 (50/50, v/v) was used as a mobile phase with a flow rate of 1 mL/min, an injection volume of 20 μL, and an oven temperature of 40°C.


## 2.2. Experimental Procedure

### 2.2.1.In Situ Peroxymonocarbonate Preparation

Peroxymonocarbonate (HCO4−) solutions were prepared in situ by mixing sodium bicarbonate with hydrogen peroxide at a molar ratio of 1 : 2 at ambient temperature. Kinetics of HCO4− formation was evaluated by monitoring HCO4− concentration at different times in three hours. In this experiment, HCO4− was analyzed by a modified iodometric titration method [12]. In which, HCO4− was first allowed to react with iodide at low temperature (below −10°C) to prevent the effect of H2O2. The formed I2 was then titrated with thiosulfate using a starch indicator.


### 2.2.2. Degradation of RB19 byIn Situ Generated HCO4− in the H2O2−HCO3− System

The effects of different parameters on the RB19 degradation were investigated as described in Table1 including the metallic ion catalyst (trials 1–9), the pH (trials 10–13), the HCO3− concentration (trials 14–18), the Co2+ catalyst concentration (trials 19–22), and UV irradiation (trials 24–28). The RB19 concentration was 100.0 mg/L in all trials. Each trial was replicated three times. In trials 1–5, 8–22, and 26, the mixture of HCO3− and H2O2 with a molar ratio of 1 : 2 was prepared in 50 minutes before adding to the RB19 solution to conduct the dye decolorization. For trials 1–23, the reactions were conducted in a 250 mL batch reactor controlled at 26 ± 1°C using a thermobath and thoroughly mixed by a magnetic stirrer as depicted in Figure 1(a). For trials 24–28, the reaction solutions were irradiated with the UVC light (254 nm, 12 W) in the UV chamber regulated by a thermobath as described in Figure 1(b). The flow speed through the UV chamber was maintained at 100 mL/min using a circulating pump. Dye decolorization was monitored by sampling at certain intervals and measuring light absorbance at the wavelength of 592 nm.Table 1
Conditions for RB19 decolorization by peroxymonocarbonate.
TrialpH[HCO3−] (mM)[H2O2] (mM)[M2+] catalyst (mg/L)UV irradiation181020—No2–5810200.1 (Ni2+, Mn2+, Zn2+, Co2+)No68—200—No78—2000.1 (Co2+)No88100200—No981002000.1 (Co2+)No10–137, 8, 9, 1010200.1 (Co2+)No14–1885, 10, 15, 25, 30400.1 (Co2+)No19–22820400.01, 0.02, 0.04, 0.06 (Co2+)No238———No248———Yes258—20—Yes268—200.1 (Co2+)Yes2781020—Yes28810200.1 (Co2+)YesFigure 1
The reaction system (a) and the continuous UV-irradiated system (b) for the RB19 degradation.
(a)(b)


## 2.2.1.In Situ Peroxymonocarbonate Preparation

Peroxymonocarbonate (HCO4−) solutions were prepared in situ by mixing sodium bicarbonate with hydrogen peroxide at a molar ratio of 1 : 2 at ambient temperature. Kinetics of HCO4− formation was evaluated by monitoring HCO4− concentration at different times in three hours. In this experiment, HCO4− was analyzed by a modified iodometric titration method [12]. In which, HCO4− was first allowed to react with iodide at low temperature (below −10°C) to prevent the effect of H2O2. The formed I2 was then titrated with thiosulfate using a starch indicator.


## 2.2.2. Degradation of RB19 byIn Situ Generated HCO4− in the H2O2−HCO3− System

The effects of different parameters on the RB19 degradation were investigated as described in Table1 including the metallic ion catalyst (trials 1–9), the pH (trials 10–13), the HCO3− concentration (trials 14–18), the Co2+ catalyst concentration (trials 19–22), and UV irradiation (trials 24–28). The RB19 concentration was 100.0 mg/L in all trials. Each trial was replicated three times. In trials 1–5, 8–22, and 26, the mixture of HCO3− and H2O2 with a molar ratio of 1 : 2 was prepared in 50 minutes before adding to the RB19 solution to conduct the dye decolorization. For trials 1–23, the reactions were conducted in a 250 mL batch reactor controlled at 26 ± 1°C using a thermobath and thoroughly mixed by a magnetic stirrer as depicted in Figure 1(a). For trials 24–28, the reaction solutions were irradiated with the UVC light (254 nm, 12 W) in the UV chamber regulated by a thermobath as described in Figure 1(b). The flow speed through the UV chamber was maintained at 100 mL/min using a circulating pump. Dye decolorization was monitored by sampling at certain intervals and measuring light absorbance at the wavelength of 592 nm.Table 1
Conditions for RB19 decolorization by peroxymonocarbonate.
TrialpH[HCO3−] (mM)[H2O2] (mM)[M2+] catalyst (mg/L)UV irradiation181020—No2–5810200.1 (Ni2+, Mn2+, Zn2+, Co2+)No68—200—No78—2000.1 (Co2+)No88100200—No981002000.1 (Co2+)No10–137, 8, 9, 1010200.1 (Co2+)No14–1885, 10, 15, 25, 30400.1 (Co2+)No19–22820400.01, 0.02, 0.04, 0.06 (Co2+)No238———No248———Yes258—20—Yes268—200.1 (Co2+)Yes2781020—Yes28810200.1 (Co2+)YesFigure 1
The reaction system (a) and the continuous UV-irradiated system (b) for the RB19 degradation.
(a)(b)


## 3. Results and Discussion

### 3.1. Formation of Peroxymonocarbonate

The peroxymonocarbonate ion was producedin situ by the reaction of HCO3− and H2O2 according to the following reversible reaction [13]: (1)HCO3−+H2O2⇌HCO4−+H2OTo determine the maximum formation ofHCO4− formed in situ in the system, the variation of HCO4– concentrations with time at a HCO3−: H2O2 molar ratio of 1 : 2 ([HCO3−] = 312.5 mM) has been studied and the result is given in Figure 2. It is revealed that HCO4− was formed at the beginning, reached a maximum concentration at around 50 minutes, and then slowly decreased. This phenomenon can be explained by the decomposition of HCO4− after formation.Figure 2
The variation of peroxymonocarbonate concentrations with time.Therefore, the dye decolorization experiments by theHCO4− agent in this study were carried out by mixing H2O2 with HCO3− 50 minutes before adding to the dye solution.


### 3.2. RB19 Degradation byH2O2−HCO3− System

#### 3.2.1. Effect of Metallic Ion Catalysts

The effect of metallic ion catalysts (i.e., Ni2+, Mn2+, Zn2+, and Co2+) was investigated by performing trials 1–5 (Table 1). The result shown in Figure 3 reveals the most significant effect of Co2+ on the degradation of RB19 by the H2O2−HCO3− system. RB19 was degraded ca. 90% after 200 min in the presence of Co2+, while the similar systems with other metallic ions such as Mn2+, Zn2+, and Ni2+ and the bare system (without metallic ions) gave the degradation efficiency of less than 10%.Figure 3
Degradation efficiency of RB19vs. time without and with different metallic ion catalysts (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Ni2+/Mn2+/Zn2+/Co2+; pH 8).To clarify the role of Co2+ in the RB19 degradation by the HCO4− agent, a comparison of RB19 degradation efficiencies by different systems (H2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+, trials 6–9) is presented in Figure 4. It clearly showed the remarkable effectiveness of Co2+ on the RB19 degradation of the H2O2-HCO3– system. It can be seen from Figure 4 that the RB19 degradation efficiency is ca. 80% at a reaction time of 40 min with the H2O2−HCO3−−Co2+ system, while three others (H2O2, H2O2−Co2+, and H2O2−HCO3−) gave the degradation efficiencies of less than 10%. This phenomenon can be explained by the great capability of Co2+ in catalyzing oxidation reactions. In peroxyacid solutions, Co2+ enhances the decomposition of peroxy compounds to generate radicals [14, 15]. Applying the same mechanism, radicals are proposed to be formed by the following reactions:(2)HCO4−+CO2+⟶CO32−+Co3++OH∗(3)HCO4−+Co2+⟶CO3•−+Co3++OH−Figure 4
Comparison of the degradation of RB19 byH2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+ systems.Co3+ is then reduced to regenerate Co2+:(4)HCO4−+Co3+⟶CO4•+Co2++H+These radicals then react with the organic compounds through many steps and eventually form CO2 and H2O. The oxidation process of organic compounds occurs continuously with a crucial role of the Co2+/Co3+ redox couple.


#### 3.2.2. Effect of pH

The effect of pH was investigated by varying pH values from 7 to 10 and keeping constant other parameters (trials 10–13). The result shown in Figure5 revealed that the increase in pH value causes faster dye degradation. However, the H2O2−HCO3− system acts as a buffer solution of pH ca. 8. Therefore, the pH value used during this study was adjusted to 8 in order to be the closest to the intrinsic pH of the solution.Figure 5
Degradation efficiency of RB19vs. time at different pH values (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Co2+; pH 7, 8, 9, and 10).


#### 3.2.3. Effect ofHCO3− and Catalyst Concentrations: Kinetic Study of the RB19 Degradation Catalyzed by Co2+

To study the effect ofHCO3− and Co2+ catalyst concentrations on the RB19 degradation, as well as to study reaction kinetics, the RB19 degradation was carried out either at different HCO3– concentrations (trials 14–18) and at different Co2+ concentrations (trials 19–22). The RB19 concentration decreased as a function of time.It can be seen from Figure6(a) that RB19 degradation efficiency and rate remarkably increase with an increase in HCO3− concentration from 5 mM to 30 mM while keeping constant other parameters (H2O2 and Co2+ concentrations, pH). It is because the increase in HCO3− concentration leads to the increase in HCO4− concentration according to the equilibrium reaction:(5)HCO3−+H2O2⇌HCO4−+H2OFigure 6
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different HCO3− concentrations and 40 mM H2O2, 0.1 mg/L Co2+, and pH 8.
(a)(b)Similarly, Figure7(a) shows a proportional relationship between the Co2+ concentration and the RB19 degradation efficiency and reaction rate. The reason for this may be the increase in Co2+ complexes as mentioned above.Figure 7
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different Co2+ concentrations and 10 mM HCO3−, 20 mM H2O2, and pH 8.
(a)(b)The plots oflnCo/Ctvs. time shown in Figures 6(b) and 7(b) reveal that the kinetic degradation of RB19 is well fitted with the first-order kinetic model given by the following equation: (6)lnCoCt=k1t.The reaction rate of RB19 degradation was assumed to be thepseudo-first-order kinetics with respect to RB19 as the following expression: (7)v=−dCtdt=k0RB19HCO3−n1H2O2n2Co2+n3H+n4=k1RB19,where(8)k1=k0HCO3−n1H2O2n2Co2+n3H+n4is constant.Integrating (7) gives (6): lnCo/Ct=k1t.Thepseudo-first-order rate constants, k1 (min−1), were calculated from the slope of the plots of ln (Co/Ct) vs. time t and subsequently used to calculate the experimental order of HCO3− and Co2+.(1) Order ofHCO3−. To determine the order of HCO3−, experiments were conducted at different HCO3− concentrations when [H2O2], [Co2+], and [H+] were kept constant.From equation (8), (9)k1=k2HCO3−n1,where k2=k0H2O2n2Co2+n3H+n4.Taking the logarithm of both sides of (9) gives(10)lnk1=n1lnHCO3−+lnk2.(2) Order of Co2+. Similarly, the order of Co2+ was determined when changing [Co2+] and keeping concentrations of other species constant.(11)whenH2O2,HCO3−,andH+are constant,k1=k3Co2+n3,where k3=k0HCO3−n1H2O2n2H+n4.Taking the logarithm of both sides of (11) gives(12)lnk1=n3lnCo2++lnk3.The experimental orders ofHCO3−, n1, and Co2+, n3, were determined from the slope of the plot of ln(k1) vs. lnHCO3− and ln(k1) vs. ln[Co2+], respectively. The calculated data are shown in Table 2, and the result gives the experimental orders of HCO3− and Co2+ of 1.7 and 1.2.Table 2
Experimental data for determining the orders of HCO3– and Co2+ in RB19 degradation.
[HCO3−]lnHCO3−k1lnk1[Co2+]ln[Co2+]k1lnk150.6990.004−2.4200.01−4.6050.0026−5.952101.0000.012−1.9210.02−3.9120.0046−5.382151.1760.018−1.7500.04−3.2190.0137−4.290251.3980.052−1.2860.06−2.8130.0211−3.858301.4770.080−1.095lnk1 = 1.7 ln[HCO3–]−3.6lnk1 = 1.2 ln[Co2+]−0.5


#### 3.2.4. Effect of UV Irradiation on the RB19 Degradation

The degradation of RB19 with the selected oxidation system (H2O2−HCO3−−Co2+) was carried out without and with UV irradiation (trials 5 and 28) compared to other systems (trials 6–8 and 23–27). The results are shown in Table 3. With UV irradiation, the RB19 degradation efficiency significantly increased (ca. 90%) in all other systems compared to the ones without UV irradiation (less than 5%). This is due to the fact that UV radiation plays a crucial role in creating the radical O•H from H2O2 for the oxidation systems according to the photocatalytic mechanism as proposed in the literature [16]. Meanwhile, the highest RB19 degradation efficiency (97.6%) was obtained with the H2O2−HCO3−−Co2+ system, slightly increased compared to the one without UV irradiation (79.9%). This strongly confirmed the degradation efficiency of the H2O2−HCO3−−Co2+ system even without UV light.Table 3
The effect of UV irradiation on the RB19 degradation.
Oxidation systemRB19 degradation efficiency at 30 min reaction time (%)Non-UV irradiationUV irradiationRB190.03.7 ± 0.3RB19-H2O20.083.6 ± 3.6RB19-H2O2-Co2+1.8 ± 0.291.1 ± 2.9RB19-H2O2-HCO3–4.7 ± 2.796.7 ± 2.3RB19-H2O2-HCO3–-Co2+79.9 ± 3.597.6 ± 3.1


### 3.3. Products of the RB19 Degradation

The COD values of the initial and final reaction solutions in which RB19 (0.1 g/L) was decomposed by theH2O2−HCO3−−Co2+ system (after a reaction time of 60 min) were determined as 315 and 12.5 mg O2/L, respectively. The TOC (total organic carbon) value was obtained from the subtraction of TC (total carbon) and TIC (total inorganic carbon) values measured for the final solution, giving the result of 15.3 mg/L which is in accordance with the COD value (Table 4).Table 4
Determination of COD and TOC of RB19 degradation.
COD (mg O2/L)TC (mg/L)TIC (mg/L)TOC (mg/L)RB19−H2O2−HCO3−−Co2+ initial315 ± 0.5———RB19−H2O2−HCO3−−Co2+ final12.5 ± 0.3141.5 ± 0.9126.2 ± 0.915.3 ± 0.2These results suggest a good mineralization of RB19 by theH2O2−HCO3−−Co2+ system with a COD removal efficiency of 96%.Moreover, the RP-HPLC chromatograms of the RB19 degradation solutions usingH2O2−HCO3−−Co2+ system at reaction times of 10, 30, and 60 min were recorded at 280 nm (Figure 8(b)). It can be seen clearly that the intermediate products of RB19 degradation at a reaction time of 10 min and 30 min are less polar and eluted before the remained RB19 (zoom in the lower inset in Figure 8(b) compared with the chromatogram of pure 1 mg/L RB19 solution shown in Figure 8(a) with a retention time of 8.2 minutes). The unretained components are eluted at the time corresponding to the travel time of the mobile phase of 1.5 min. Most of these products are not observed after the reaction time of 30 min, and the chromatogram obtained for the reaction solution at 60 min is flattened (the upper inset). This result strongly supports the effectiveness of the RB19 degradation to a complete mineralization by the H2O2−HCO3−−Co2+ system.Figure 8
(a) The chromatogram of pure RB19 solution (1 mg/L); (b) the chromatograms of the RB19 degradation solution by theH2O2−HCO3−−Co2+ system at different reaction times (t = 10, 30, and 60 min); the upper inset: the chromatogram recorded at a reaction time of 60 min; the lower inset: zoom-in view of the chromatogram recorded at a reaction time of 10 min.
(a)(b)


## 3.1. Formation of Peroxymonocarbonate

The peroxymonocarbonate ion was producedin situ by the reaction of HCO3− and H2O2 according to the following reversible reaction [13]: (1)HCO3−+H2O2⇌HCO4−+H2OTo determine the maximum formation ofHCO4− formed in situ in the system, the variation of HCO4– concentrations with time at a HCO3−: H2O2 molar ratio of 1 : 2 ([HCO3−] = 312.5 mM) has been studied and the result is given in Figure 2. It is revealed that HCO4− was formed at the beginning, reached a maximum concentration at around 50 minutes, and then slowly decreased. This phenomenon can be explained by the decomposition of HCO4− after formation.Figure 2
The variation of peroxymonocarbonate concentrations with time.Therefore, the dye decolorization experiments by theHCO4− agent in this study were carried out by mixing H2O2 with HCO3− 50 minutes before adding to the dye solution.


## 3.2. RB19 Degradation byH2O2−HCO3− System

### 3.2.1. Effect of Metallic Ion Catalysts

The effect of metallic ion catalysts (i.e., Ni2+, Mn2+, Zn2+, and Co2+) was investigated by performing trials 1–5 (Table 1). The result shown in Figure 3 reveals the most significant effect of Co2+ on the degradation of RB19 by the H2O2−HCO3− system. RB19 was degraded ca. 90% after 200 min in the presence of Co2+, while the similar systems with other metallic ions such as Mn2+, Zn2+, and Ni2+ and the bare system (without metallic ions) gave the degradation efficiency of less than 10%.Figure 3
Degradation efficiency of RB19vs. time without and with different metallic ion catalysts (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Ni2+/Mn2+/Zn2+/Co2+; pH 8).To clarify the role of Co2+ in the RB19 degradation by the HCO4− agent, a comparison of RB19 degradation efficiencies by different systems (H2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+, trials 6–9) is presented in Figure 4. It clearly showed the remarkable effectiveness of Co2+ on the RB19 degradation of the H2O2-HCO3– system. It can be seen from Figure 4 that the RB19 degradation efficiency is ca. 80% at a reaction time of 40 min with the H2O2−HCO3−−Co2+ system, while three others (H2O2, H2O2−Co2+, and H2O2−HCO3−) gave the degradation efficiencies of less than 10%. This phenomenon can be explained by the great capability of Co2+ in catalyzing oxidation reactions. In peroxyacid solutions, Co2+ enhances the decomposition of peroxy compounds to generate radicals [14, 15]. Applying the same mechanism, radicals are proposed to be formed by the following reactions:(2)HCO4−+CO2+⟶CO32−+Co3++OH∗(3)HCO4−+Co2+⟶CO3•−+Co3++OH−Figure 4
Comparison of the degradation of RB19 byH2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+ systems.Co3+ is then reduced to regenerate Co2+:(4)HCO4−+Co3+⟶CO4•+Co2++H+These radicals then react with the organic compounds through many steps and eventually form CO2 and H2O. The oxidation process of organic compounds occurs continuously with a crucial role of the Co2+/Co3+ redox couple.


### 3.2.2. Effect of pH

The effect of pH was investigated by varying pH values from 7 to 10 and keeping constant other parameters (trials 10–13). The result shown in Figure5 revealed that the increase in pH value causes faster dye degradation. However, the H2O2−HCO3− system acts as a buffer solution of pH ca. 8. Therefore, the pH value used during this study was adjusted to 8 in order to be the closest to the intrinsic pH of the solution.Figure 5
Degradation efficiency of RB19vs. time at different pH values (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Co2+; pH 7, 8, 9, and 10).


### 3.2.3. Effect ofHCO3− and Catalyst Concentrations: Kinetic Study of the RB19 Degradation Catalyzed by Co2+

To study the effect ofHCO3− and Co2+ catalyst concentrations on the RB19 degradation, as well as to study reaction kinetics, the RB19 degradation was carried out either at different HCO3– concentrations (trials 14–18) and at different Co2+ concentrations (trials 19–22). The RB19 concentration decreased as a function of time.It can be seen from Figure6(a) that RB19 degradation efficiency and rate remarkably increase with an increase in HCO3− concentration from 5 mM to 30 mM while keeping constant other parameters (H2O2 and Co2+ concentrations, pH). It is because the increase in HCO3− concentration leads to the increase in HCO4− concentration according to the equilibrium reaction:(5)HCO3−+H2O2⇌HCO4−+H2OFigure 6
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different HCO3− concentrations and 40 mM H2O2, 0.1 mg/L Co2+, and pH 8.
(a)(b)Similarly, Figure7(a) shows a proportional relationship between the Co2+ concentration and the RB19 degradation efficiency and reaction rate. The reason for this may be the increase in Co2+ complexes as mentioned above.Figure 7
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different Co2+ concentrations and 10 mM HCO3−, 20 mM H2O2, and pH 8.
(a)(b)The plots oflnCo/Ctvs. time shown in Figures 6(b) and 7(b) reveal that the kinetic degradation of RB19 is well fitted with the first-order kinetic model given by the following equation: (6)lnCoCt=k1t.The reaction rate of RB19 degradation was assumed to be thepseudo-first-order kinetics with respect to RB19 as the following expression: (7)v=−dCtdt=k0RB19HCO3−n1H2O2n2Co2+n3H+n4=k1RB19,where(8)k1=k0HCO3−n1H2O2n2Co2+n3H+n4is constant.Integrating (7) gives (6): lnCo/Ct=k1t.Thepseudo-first-order rate constants, k1 (min−1), were calculated from the slope of the plots of ln (Co/Ct) vs. time t and subsequently used to calculate the experimental order of HCO3− and Co2+.(1) Order ofHCO3−. To determine the order of HCO3−, experiments were conducted at different HCO3− concentrations when [H2O2], [Co2+], and [H+] were kept constant.From equation (8), (9)k1=k2HCO3−n1,where k2=k0H2O2n2Co2+n3H+n4.Taking the logarithm of both sides of (9) gives(10)lnk1=n1lnHCO3−+lnk2.(2) Order of Co2+. Similarly, the order of Co2+ was determined when changing [Co2+] and keeping concentrations of other species constant.(11)whenH2O2,HCO3−,andH+are constant,k1=k3Co2+n3,where k3=k0HCO3−n1H2O2n2H+n4.Taking the logarithm of both sides of (11) gives(12)lnk1=n3lnCo2++lnk3.The experimental orders ofHCO3−, n1, and Co2+, n3, were determined from the slope of the plot of ln(k1) vs. lnHCO3− and ln(k1) vs. ln[Co2+], respectively. The calculated data are shown in Table 2, and the result gives the experimental orders of HCO3− and Co2+ of 1.7 and 1.2.Table 2
Experimental data for determining the orders of HCO3– and Co2+ in RB19 degradation.
[HCO3−]lnHCO3−k1lnk1[Co2+]ln[Co2+]k1lnk150.6990.004−2.4200.01−4.6050.0026−5.952101.0000.012−1.9210.02−3.9120.0046−5.382151.1760.018−1.7500.04−3.2190.0137−4.290251.3980.052−1.2860.06−2.8130.0211−3.858301.4770.080−1.095lnk1 = 1.7 ln[HCO3–]−3.6lnk1 = 1.2 ln[Co2+]−0.5


### 3.2.4. Effect of UV Irradiation on the RB19 Degradation

The degradation of RB19 with the selected oxidation system (H2O2−HCO3−−Co2+) was carried out without and with UV irradiation (trials 5 and 28) compared to other systems (trials 6–8 and 23–27). The results are shown in Table 3. With UV irradiation, the RB19 degradation efficiency significantly increased (ca. 90%) in all other systems compared to the ones without UV irradiation (less than 5%). This is due to the fact that UV radiation plays a crucial role in creating the radical O•H from H2O2 for the oxidation systems according to the photocatalytic mechanism as proposed in the literature [16]. Meanwhile, the highest RB19 degradation efficiency (97.6%) was obtained with the H2O2−HCO3−−Co2+ system, slightly increased compared to the one without UV irradiation (79.9%). This strongly confirmed the degradation efficiency of the H2O2−HCO3−−Co2+ system even without UV light.Table 3
The effect of UV irradiation on the RB19 degradation.
Oxidation systemRB19 degradation efficiency at 30 min reaction time (%)Non-UV irradiationUV irradiationRB190.03.7 ± 0.3RB19-H2O20.083.6 ± 3.6RB19-H2O2-Co2+1.8 ± 0.291.1 ± 2.9RB19-H2O2-HCO3–4.7 ± 2.796.7 ± 2.3RB19-H2O2-HCO3–-Co2+79.9 ± 3.597.6 ± 3.1


## 3.2.1. Effect of Metallic Ion Catalysts

The effect of metallic ion catalysts (i.e., Ni2+, Mn2+, Zn2+, and Co2+) was investigated by performing trials 1–5 (Table 1). The result shown in Figure 3 reveals the most significant effect of Co2+ on the degradation of RB19 by the H2O2−HCO3− system. RB19 was degraded ca. 90% after 200 min in the presence of Co2+, while the similar systems with other metallic ions such as Mn2+, Zn2+, and Ni2+ and the bare system (without metallic ions) gave the degradation efficiency of less than 10%.Figure 3
Degradation efficiency of RB19vs. time without and with different metallic ion catalysts (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Ni2+/Mn2+/Zn2+/Co2+; pH 8).To clarify the role of Co2+ in the RB19 degradation by the HCO4− agent, a comparison of RB19 degradation efficiencies by different systems (H2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+, trials 6–9) is presented in Figure 4. It clearly showed the remarkable effectiveness of Co2+ on the RB19 degradation of the H2O2-HCO3– system. It can be seen from Figure 4 that the RB19 degradation efficiency is ca. 80% at a reaction time of 40 min with the H2O2−HCO3−−Co2+ system, while three others (H2O2, H2O2−Co2+, and H2O2−HCO3−) gave the degradation efficiencies of less than 10%. This phenomenon can be explained by the great capability of Co2+ in catalyzing oxidation reactions. In peroxyacid solutions, Co2+ enhances the decomposition of peroxy compounds to generate radicals [14, 15]. Applying the same mechanism, radicals are proposed to be formed by the following reactions:(2)HCO4−+CO2+⟶CO32−+Co3++OH∗(3)HCO4−+Co2+⟶CO3•−+Co3++OH−Figure 4
Comparison of the degradation of RB19 byH2O2, H2O2−Co2+, H2O2−HCO3−, and H2O2−HCO3−−Co2+ systems.Co3+ is then reduced to regenerate Co2+:(4)HCO4−+Co3+⟶CO4•+Co2++H+These radicals then react with the organic compounds through many steps and eventually form CO2 and H2O. The oxidation process of organic compounds occurs continuously with a crucial role of the Co2+/Co3+ redox couple.


## 3.2.2. Effect of pH

The effect of pH was investigated by varying pH values from 7 to 10 and keeping constant other parameters (trials 10–13). The result shown in Figure5 revealed that the increase in pH value causes faster dye degradation. However, the H2O2−HCO3− system acts as a buffer solution of pH ca. 8. Therefore, the pH value used during this study was adjusted to 8 in order to be the closest to the intrinsic pH of the solution.Figure 5
Degradation efficiency of RB19vs. time at different pH values (100 mg/L RB19, 10 mM HCO3−, 20 mM H2O2, and 0.1 mg/L Co2+; pH 7, 8, 9, and 10).


## 3.2.3. Effect ofHCO3− and Catalyst Concentrations: Kinetic Study of the RB19 Degradation Catalyzed by Co2+

To study the effect ofHCO3− and Co2+ catalyst concentrations on the RB19 degradation, as well as to study reaction kinetics, the RB19 degradation was carried out either at different HCO3– concentrations (trials 14–18) and at different Co2+ concentrations (trials 19–22). The RB19 concentration decreased as a function of time.It can be seen from Figure6(a) that RB19 degradation efficiency and rate remarkably increase with an increase in HCO3− concentration from 5 mM to 30 mM while keeping constant other parameters (H2O2 and Co2+ concentrations, pH). It is because the increase in HCO3− concentration leads to the increase in HCO4− concentration according to the equilibrium reaction:(5)HCO3−+H2O2⇌HCO4−+H2OFigure 6
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different HCO3− concentrations and 40 mM H2O2, 0.1 mg/L Co2+, and pH 8.
(a)(b)Similarly, Figure7(a) shows a proportional relationship between the Co2+ concentration and the RB19 degradation efficiency and reaction rate. The reason for this may be the increase in Co2+ complexes as mentioned above.Figure 7
Dependence of (a) RB19 concentration and (b)lnCo/Ctvs. time at different Co2+ concentrations and 10 mM HCO3−, 20 mM H2O2, and pH 8.
(a)(b)The plots oflnCo/Ctvs. time shown in Figures 6(b) and 7(b) reveal that the kinetic degradation of RB19 is well fitted with the first-order kinetic model given by the following equation: (6)lnCoCt=k1t.The reaction rate of RB19 degradation was assumed to be thepseudo-first-order kinetics with respect to RB19 as the following expression: (7)v=−dCtdt=k0RB19HCO3−n1H2O2n2Co2+n3H+n4=k1RB19,where(8)k1=k0HCO3−n1H2O2n2Co2+n3H+n4is constant.Integrating (7) gives (6): lnCo/Ct=k1t.Thepseudo-first-order rate constants, k1 (min−1), were calculated from the slope of the plots of ln (Co/Ct) vs. time t and subsequently used to calculate the experimental order of HCO3− and Co2+.(1) Order ofHCO3−. To determine the order of HCO3−, experiments were conducted at different HCO3− concentrations when [H2O2], [Co2+], and [H+] were kept constant.From equation (8), (9)k1=k2HCO3−n1,where k2=k0H2O2n2Co2+n3H+n4.Taking the logarithm of both sides of (9) gives(10)lnk1=n1lnHCO3−+lnk2.(2) Order of Co2+. Similarly, the order of Co2+ was determined when changing [Co2+] and keeping concentrations of other species constant.(11)whenH2O2,HCO3−,andH+are constant,k1=k3Co2+n3,where k3=k0HCO3−n1H2O2n2H+n4.Taking the logarithm of both sides of (11) gives(12)lnk1=n3lnCo2++lnk3.The experimental orders ofHCO3−, n1, and Co2+, n3, were determined from the slope of the plot of ln(k1) vs. lnHCO3− and ln(k1) vs. ln[Co2+], respectively. The calculated data are shown in Table 2, and the result gives the experimental orders of HCO3− and Co2+ of 1.7 and 1.2.Table 2
Experimental data for determining the orders of HCO3– and Co2+ in RB19 degradation.
[HCO3−]lnHCO3−k1lnk1[Co2+]ln[Co2+]k1lnk150.6990.004−2.4200.01−4.6050.0026−5.952101.0000.012−1.9210.02−3.9120.0046−5.382151.1760.018−1.7500.04−3.2190.0137−4.290251.3980.052−1.2860.06−2.8130.0211−3.858301.4770.080−1.095lnk1 = 1.7 ln[HCO3–]−3.6lnk1 = 1.2 ln[Co2+]−0.5


## 3.2.4. Effect of UV Irradiation on the RB19 Degradation

The degradation of RB19 with the selected oxidation system (H2O2−HCO3−−Co2+) was carried out without and with UV irradiation (trials 5 and 28) compared to other systems (trials 6–8 and 23–27). The results are shown in Table 3. With UV irradiation, the RB19 degradation efficiency significantly increased (ca. 90%) in all other systems compared to the ones without UV irradiation (less than 5%). This is due to the fact that UV radiation plays a crucial role in creating the radical O•H from H2O2 for the oxidation systems according to the photocatalytic mechanism as proposed in the literature [16]. Meanwhile, the highest RB19 degradation efficiency (97.6%) was obtained with the H2O2−HCO3−−Co2+ system, slightly increased compared to the one without UV irradiation (79.9%). This strongly confirmed the degradation efficiency of the H2O2−HCO3−−Co2+ system even without UV light.Table 3
The effect of UV irradiation on the RB19 degradation.
Oxidation systemRB19 degradation efficiency at 30 min reaction time (%)Non-UV irradiationUV irradiationRB190.03.7 ± 0.3RB19-H2O20.083.6 ± 3.6RB19-H2O2-Co2+1.8 ± 0.291.1 ± 2.9RB19-H2O2-HCO3–4.7 ± 2.796.7 ± 2.3RB19-H2O2-HCO3–-Co2+79.9 ± 3.597.6 ± 3.1


## 3.3. Products of the RB19 Degradation

The COD values of the initial and final reaction solutions in which RB19 (0.1 g/L) was decomposed by theH2O2−HCO3−−Co2+ system (after a reaction time of 60 min) were determined as 315 and 12.5 mg O2/L, respectively. The TOC (total organic carbon) value was obtained from the subtraction of TC (total carbon) and TIC (total inorganic carbon) values measured for the final solution, giving the result of 15.3 mg/L which is in accordance with the COD value (Table 4).Table 4
Determination of COD and TOC of RB19 degradation.
COD (mg O2/L)TC (mg/L)TIC (mg/L)TOC (mg/L)RB19−H2O2−HCO3−−Co2+ initial315 ± 0.5———RB19−H2O2−HCO3−−Co2+ final12.5 ± 0.3141.5 ± 0.9126.2 ± 0.915.3 ± 0.2These results suggest a good mineralization of RB19 by theH2O2−HCO3−−Co2+ system with a COD removal efficiency of 96%.Moreover, the RP-HPLC chromatograms of the RB19 degradation solutions usingH2O2−HCO3−−Co2+ system at reaction times of 10, 30, and 60 min were recorded at 280 nm (Figure 8(b)). It can be seen clearly that the intermediate products of RB19 degradation at a reaction time of 10 min and 30 min are less polar and eluted before the remained RB19 (zoom in the lower inset in Figure 8(b) compared with the chromatogram of pure 1 mg/L RB19 solution shown in Figure 8(a) with a retention time of 8.2 minutes). The unretained components are eluted at the time corresponding to the travel time of the mobile phase of 1.5 min. Most of these products are not observed after the reaction time of 30 min, and the chromatogram obtained for the reaction solution at 60 min is flattened (the upper inset). This result strongly supports the effectiveness of the RB19 degradation to a complete mineralization by the H2O2−HCO3−−Co2+ system.Figure 8
(a) The chromatogram of pure RB19 solution (1 mg/L); (b) the chromatograms of the RB19 degradation solution by theH2O2−HCO3−−Co2+ system at different reaction times (t = 10, 30, and 60 min); the upper inset: the chromatogram recorded at a reaction time of 60 min; the lower inset: zoom-in view of the chromatogram recorded at a reaction time of 10 min.
(a)(b)


## 4. Conclusions

In the present work, the degradation efficiency of RB19 reactive dye has been proved to be strongly affected by theHCO3− concentration and the presence of metallic ion catalysts, especially Co2+. This improvement can be explained by an enhancement in radicals generated by the Co2+/Co3+ redox couple. The in situ formation of peroxymonocarbonate ion HCO4− is responsible for the oxidation efficiency of the system. This HCO4− ion was determined to reach a maximum concentration after a H2O2−HCO3− mixing time of 40 min and remain stable for ca. 10 min. This kind of information is essential for practical applications of the system. The degradation kinetics of RB19 was determined to be the first order in RB19 dye while reaction orders of HCO3– and Co2+ catalyst were 1.7 and 1.2, respectively. The products of RB19 degradation were investigated by RP-HPLC analyses and COD and TOC measurements, which suggest a complete mineralization of RB19 by the H2O2−HCO3−−Co2+ system.



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*Source: 6696600-2021-03-04.xml*