Introduction

The effluent that was discharged from the dyeing industry was an important source of environmental pollution (Salinas et al. 2018; Wang et al. 2018). Due to the existences of dyes, the effluent was usually strongly colored. Discharging dye wastewater into receiving water bodies without effective treatment would affect the balance of the aquatic ecosystem, and even cause damage to human health because of the mutagenic, carcinogenic, and toxic effects of the dyes (Jain and Gogate 2018; Long et al. 2017). In particular, azo dyes and triphenylmethane dyes attracted special attention due to the extensive use in textile, printing, food, and cosmetic industries. Moreover, both types of dyes were characterized by the complex and steady molecular structures making them quite resistant to biological degradation and exert long-term adverse effects on the aquatic environment (Babendure et al. 2003; Sun et al. 2012; Yang et al. 2012).

With the increasing concentrations of dyes occurring in effluents, the development of degrading dyes from textile dyeing wastewater had been more and more crucial than ever. Among versatile biological, physical, and chemical technologies in pollutant treatments (Falås et al. 2016; Ganzenko et al. 2014; Li et al. 2017; Mouele et al. 2015; Rozas et al. 2016; Teh et al. 2016), advanced oxidation processes (AOPs) were one of the most attractive technologies due to the rate in pollutant removing (Benzaquén et al. 2015; Li et al. 2011; Pignatello et al. 2006; Wang et al. 2011). Several AOPs, such as TiO2-mediated photocatalysis (Li et al. 2011), Fenton (Pignatello et al. 2006) or photo-Fenton reaction (Benzaquén et al. 2015), and Co2+/ozone process (Wang et al. 2011), had been examined for the degradation of dyes. AOPs based on the activation of hydrogen peroxide (H2O2) produce highly powerful hydroxyl radical (·OH) and non-selectively degrade large amounts of organic pollutants. The process was considered to be one of the most economical, simple, and effective methods (And and Gallard 1999; Asghar et al. 2015; Benitez et al. 2001; Bokare and Choi 2014, Detomaso et al. 2003, Gemeay et al. 2010, Li et al. 2013, Long et al. 2012, Nidheesh et al. 2013, Rivas et al. 2015, Sun et al. 2009). The activation of H2O2 with transition metal ions (Ling et al. 2010), electrolysis (Ltaïef et al. 2018), microwave irradiation (Yuan and Hong 2012), ultrasonic irradiation (Kang et al. 1999), and UV irradiation (Anipsitakis and Dionysiou 2004) had been widely used in wastewater industry to degrade the toxic organic pollutants. Nevertheless, the traditional Fenton reagent (Fe2+/H2O2) requires strict acidic conditions (pH < 4) while many wastewaters were characterized by neutral or slightly alkaline (Bokare and Choi 2014). As for the other activation methods such as electrolysis, microwave irradiation, ultrasonic irradiation, and UV irradiation, they all required extra power which would add additional cost in practice. Therefore, a simple, efficient, and low-cost activation method is still needed for degrading organic pollutants.

Recent studies showed that H2O2 under alkaline conditions would be ionized to form hydroperoxy anion (HO2−) and decomposed to generate reactive oxygen species such as hydroxyl radical (·OH), superoxide radical (O2·−), and singlet oxygen (1O2), which were highly efficient for the degradation of organic pollutants (Fragoso et al. 2009; Gould 1985; Hayase et al. 1984; Katafias et al. 2010; Thasilu and Karthikeyan 2016; Wright and Abbot 1993; Xu et al. 2011a; Yuan and Hong 2012). Moreover, the process of alkali-activated H2O2 was easily implementable since the extra reagents required in this system were only H2O2 and sodium hydroxide. Meanwhile, it was well known that many dye wastewaters were commonly characterized by high alkalinity (pH > 12.0) (Li et al. 2018; Prisciandaro et al. 2005; Rao et al. 2007). Thus, the process of alkali-activated H2O2 was highly feasible and superior for the treatment of alkaline dye wastewaters.

However, an important and interesting phenomenon had been observed that the primary reactive species were found different for degrading organic pollutants with the process of alkali-activated H2O2. In the research reported by Yuan et al., ·OH was regarded as the oxidizing species in microwave irradiation/H2O2 system under alkaline circumstance for the degradation of Rhodamine B (Yuan and Hong 2012). However, Katafias et al. found HO2− was the dominant oxidizing species in the reaction for the degradation of methylene blue (Katafias et al. 2010). And in the study reported by Hayase et al., HO2− was also suggested as the primary oxidant for the degradation of melanoidins in alkaline circumstance (Hayase et al. 1984). Moreover, Gould and co-workers found that both ·OH and O2·− were the primary oxidizing species in the delignification process of agricultural residues (Gould 1985). Therefore, identifying the primary oxidant responsible for organic pollutant degradation in the process of alkali-activated H2O2 was highly necessary. In the present study, a typical azo dye of Orange G (OG) and a common triphenylmethane dye of Rhodamine B (RhB) were chosen as the target organic pollutants, and the aims of this study were (i) to investigate the degradation efficiencies of OG and RhB by alkali-activated H2O2, (ii) to study the influences of operational parameters on OG and RhB degradation by alkali-activated H2O2, and (iii) to identify the primary oxidant responsible for OG and RhB degradation by alkali-activated H2O2 using quenching experiments.

Materials and methods

Materials

Chemicals including RhB, t-butanol, and furfuryl alcohol were purchased from Shanghai Aladdin Bio-Chem Technology Co., Ltd. OG, hydrogen peroxide (H2O2), sodium hydroxide (NaOH), ascorbic acid, methanol, isopropanol, and nitro blue tetrazolium (NBT) were obtained from Sinopharm Chemical Reagent Co., Ltd. All of chemicals used were analytical grade without any further purification. Stock solutions were prepared with ultrapure water (18.2 MΩ cm) produced from an ultrapure water system (Milli-Q Biocel, USA).

Experimental procedures

All batch experiments were performed in 150-mL conical beakers under light and at 25 °C. The concentrations of both OG and RhB stock solutions were 10 mM. A 10-mM OG stock solution was obtained by dissolving 0.452 g OG in 100 mL of ultrapure water. A 10-mM RhB stock solution was obtained by dissolving 0.479 g RhB in 100 mL of ultrapure water. The concentration of H2O2 stock solution was 900 mM. In a typical degradation experiment, a desired volume of ultrapure water and dye stock solution was initially added into the beaker, and then NaOH solution was used to adjust solution pH to the desired value. Finally, H2O2 stock solution was added to the above solution to initiate the reaction (H2O2 stock solution was prepared by the dilution of 30% H2O2 (9.76 M)). The total volume of the reaction solution was 100 mL. Control experiments were conducted without the addition of H2O2. The quenching experiments employed ascorbic acid, methanol, t-butanol, isopropanol, furfuryl alcohol, and NBT as quenchers, and then the desired dosages of scavengers were separately added into the beaker before dosing H2O2 stock solution.

At a given reaction time, 1 mL of aqueous solution was taken from the bottle by the pipette. Since H2O2 could not degrade OG (pKa = 11.5) and RhB (pKa = 3.0–3.7) under neutral pH conditions, reaction solutions in quartz cuvettes were then adjusted to neutral pH by adding 0.5 M of pH 6.0 phosphate buffer solution at a volume ratio of 1:1 to quench the degradation reactions and avoid the impact of pH on the measurements of OG and RhB.

Analysis and calculations

The absorption spectra from 200 to 800 nm of the dyes were recorded using an UV-Vis spectrophotometer (TU-1901, Persee Instruments, China). The absorption values of dyes were measured at the maximum absorbance wavelength (λmax) where OG was 478 nm and RhB was 555 nm. The dye degradation efficiency and the corresponding apparent rate constant (kobs) in a sample were calculated according to Eqs. (1) and (2), respectively:

$$ \mathrm{DE}\left(\%\right)=\frac{C_0-{C}_t}{C_0}\times 100 $$
(1)
$$ \mathrm{Ln}\left({C}_0/{C}_t\right)={k}_{\mathrm{obs}}\times t $$
(2)

where:

DE:

the dye degradation efficiency (%)

C 0 :

the initial dye concentration (mM)

C t :

the residual dye concentration at reaction time of t (min) (mM)

k obs :

the corresponding apparent rate constant (min−1)

Results and discussion

General observations

The effects of H2O2 on the degradation of OG and RhB were investigated at alkaline pH, and the results are shown in Fig. 1. Control experiments were conducted in the absence of H2O2, and the outcomes showed little degradation of OG and RhB. However, the degradation efficiencies of OG and RhB were tremendously accelerated by the addition of H2O2. About 74% of OG and 99% of RhB were respectively decomposed within 60 min at pH 11.5 and pH 13.0. Additionally, it was found that solution pH was nearly unchanged during the reaction (results not shown), which was due to the strong buffering capacity of the highly alkaline solution. Consequently, the process of alkali-activated H2O2 was highly effective for degrading OG and RhB. Additionally, the changes of TOC concentrations were also measured during the degradation of OG and RhB with alkaline-activated H2O2, and the TOC removals for OG and RhB after 60 min were respectively 10.4% and 8.3% as shown in Fig. S1.

Fig. 1
figure 1

Degradation of OG and RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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Effect of solution pH

Solution pH is an important factor in AOPs based on the activation of H2O2, since it can significantly affect the degradation of dye pollutants (Chahbane et al. 2007; Huang et al. 2003). Figure 2 showed the influence of initial solution pH on the degradation of OG and RhB in the process of alkali-activated H2O2.

Fig. 2
figure 2

Effect of initial solution pH on the degradation of a OG and b RhB by alkaline-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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As shown in Fig. 2a, with the addition of 36 mM H2O2, the degradation efficiency of OG within 60 min was significantly increased from about 7% to 74% when solution pH was ranged from 10.0 to 11.5, and then promptly decreased to about 4% with further increasing pH to 13.0. Moreover, the kinetic of OG degradation in the pH range of 10.0–13.0 was well consistent with a second-order reaction (Fig. S2 and Tab. S2). kobs of OG degradation was increased from 0.00139 to 0.02225 min−1 and then decreased to 0.00037 min−1 when solution pH ranged from 10.0 to 13.0 (Fig. 3, Fig. S2 and Tab. S2). Consequently, the optimum pH for OG degradation in the process of alkali-activated H2O2 was 11.5.

Fig. 3
figure 3

Effect of solution pH on αHO2− and kobs. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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Figure 2b showed the effect of solution pH on RhB degradation by alkali-activated H2O2. When solution pH was increased from 10.0 to 13.0, the degradation efficiency of RhB within 60 min was continuously increased from about 5 to 99%. Similarly, the kinetic of RhB degradation in the pH range of 10.0–13.0 was also accorded with a second-order reaction (Fig. S2 and Tab. S2). kobs of RhB degradation was continuously increased from 0.00076 to 0.09703 min−1 when the solution pH was ranged from 10.0 to 13.0 (Fig. 3, Fig. S2 and Tab. S2). In other words, the degradation of RhB was continuously accelerated by increasing solution pH in the pH range of 10.0–13.0. It indicated that the effect of solution pH was obviously different for the degradation of OG and RhB in the process of alkali-activated H2O2.

Importantly, two interesting phenomena were observed in Fig. 3. One was that the optimum solution pH for OG degradation (pH = 11.5) was corresponding with the pKa of H2O2 (pKa = 11.62) (Christensen et al. 1982), and the other was that there was a positive correlation between kobs of RhB degradation and the fraction of HO2− (αHO2−) in the pH range of 10.0–13.0. The stability of H2O2 under alkaline conditions had been reported to be pH dependent in previous literatures (Buxton et al. 1988; Christensen et al. 1982; Cui et al. 2017; Fu et al. 2017; Miao et al. 2015), and the self-decomposition of H2O2 tended to be fastest in the pH range of 11–12 (Qiang et al. 2002). Moreover, the primary pathways for the decomposition of H2O2 to generate reactive oxygen species (i.e., ·OH, O2·−, and 1O2) under alkaline conditions are shown in Eqs. (3)–(14) (Buxton et al. 1988; Christensen et al. 1982; Cui et al. 2017; Fu et al. 2017; Miao et al. 2015; Qiang et al. 2002). Additionally, HO2−, the product of deprotonation of H2O2 (Eq. (3)), had been reported to be more efficient for phenothiazine degradation (Katafias et al. 2010). As a result, the assumption that the primary oxidants for OG degradation and RhB degradation were respectively reactive oxygen species and HO2− was proposed in this study.

$$ {\mathrm{H}}_2{\mathrm{O}}_2\leftrightarrow {\mathrm{H}\mathrm{O}}_2^{-}+{\mathrm{H}}^{+}\kern0.75em \mathrm{p}{K}_{\mathrm{a}}=11.62 $$
(3)
$$ {\mathrm{H}\mathrm{O}}_2^{-}+{\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{H}\mathrm{O}}_2^{\cdotp }+\cdotp \mathrm{OH}+{\mathrm{O}\mathrm{H}}^{-}\kern0.75em k={3\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(4)
$$ {\mathrm{H}\mathrm{O}}_2^{\cdotp }+\cdotp \mathrm{OH}\to {\mathrm{H}}_2\mathrm{O}+{\mathrm{O}}_2\kern0.75em k=7.5\times {10}^9{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(5)
$$ {\mathrm{H}\mathrm{O}}_2^{\cdotp }+{\mathrm{O}}_2^{\cdotp -}\to {\mathrm{H}}_2{\mathrm{O}}_2+{\mathrm{O}}_2\kern0.75em k=9.7\times {10}^7{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(6)
$$ \cdotp \mathrm{OH}+{\mathrm{O}}_2^{\cdotp -}\to {\mathrm{O}\mathrm{H}}^{-}+{\mathrm{O}}_2\kern0.75em k=1.01\times {10}^{10}{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(7)
$$ {\mathrm{H}\mathrm{O}}_2^{\cdotp }+{\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{H}}_2\mathrm{O}+\cdotp \mathrm{OH}+{\mathrm{O}}_2\kern0.75em k=0.5\pm 0.09{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(8)
$$ {\mathrm{H}\mathrm{O}}_2^{\cdotp}\leftrightarrow {\mathrm{O}}_2^{\cdotp -}+{\mathrm{H}}^{+}\kern0.75em \mathrm{p}{K}_{\mathrm{a}}=4.8 $$
(9)
$$ \cdotp \mathrm{OH}+{\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{H}}_2\mathrm{O}+{\mathrm{O}}_2^{\cdotp -}\kern0.75em k=2.7\times {10}^7{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(10)
$$ \cdotp \mathrm{OH}+{\mathrm{H}\mathrm{O}}_2^{-}\to {\mathrm{H}}_2\mathrm{O}+{\mathrm{O}}_2^{\cdotp -}\kern0.75em k=7.5\times {10}^9{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(11)
$$ \cdotp \mathrm{OH}+\cdotp \mathrm{OH}\to {\mathrm{H}}_2{\mathrm{O}}_2\kern0.75em k=5.5\times {10}^9{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(12)
$$ \cdotp \mathrm{OH}+{\mathrm{O}}_2^{\cdotp -}\to {\mathrm{O}\mathrm{H}}^{-}+{{}^1\mathrm{O}}_2\kern0.75em k=8\times {10}^9{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(13)
$$ {\mathrm{H}\mathrm{O}}_2^{\cdotp }+{\mathrm{H}\mathrm{O}}_2^{\cdotp}\to {\mathrm{H}}_2{\mathrm{O}}_2+{{}^1\mathrm{O}}_2\kern0.75em k=8.3\times {10}^5{\mathrm{M}}^{-1}{\mathrm{s}}^{-1} $$
(14)

Effect of H2O2 concentration

Since H2O2 was the source of reactive oxygen species generation and HO2− formation (Buxton et al. 1988, Christensen et al. 1982, Cui et al. 2017, Fu et al. 2017, Miao et al. 2015, Qiang et al. 2002), the effect of H2O2 concentration on the degradation of OG and RhB was studied. As shown in Fig. S3(a), the degradation efficiency of OG was increased rapidly from about 24 to 91% when H2O2 concentration increased from 9 to 72 mM. Correspondingly, kobs of OG degradation was almost linearly increased from 0.00484 to 0.04331 min−1 when H2O2 concentration increased from 9 to 72 mM (Fig. 4, Fig. S2 and Tab. S2). As shown in Fig. S3(b), the RhB degradation efficiency was augmented from 73% to near 100% when H2O2 concentration was increased from 9 to 72 mM. The corresponding kobs was increased from 0.02077 to 0.18078 min−1 with the increasing H2O2 concentration (Fig. 4, Fig. S2 and Tab. S2). Hence, it could be inferred that more reactive oxygen species and HO2− were generated with higher H2O2 concentration, resulting in higher degradation efficiencies of OG and RhB by alkali-activated H2O2.

Fig. 4
figure 4

Effect of H2O2 concentration on the degradation of OG and RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, temperature (T) = 25 ± 2 °C

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Effect of reaction temperature

Figure S4 showed the effect of reaction temperature on the degradation of OG and RhB degradation by alkali-activated H2O2. The degradation efficiencies of OG and RhB within 60 min were respectively about 38% and 87% when reaction temperature was 10 °C. Higher degradation efficiencies were observed when reaction temperature increased to 40 °C. The major reason was that higher temperature could enhance reactive oxygen species generation and HO2− formation, further enhancing the degradation of OG and RhB. In addition, the degradation of OG and RhB between 10 and 40 °C followed the second-order kinetics well (Fig. S2 and Tab. S2). kobs of OG degradation and RhB degradation by alkali-activated H2O2 could be also calculated with Arrhenius’ equation (Eq. (15)) (Sun et al. 2009).

$$ {k}_{\mathrm{obs}}=A\exp \left(\frac{-E}{RT}\right) $$
(15)

where:

A :

the pre-exponential factor

E :

the activation energy for the reaction (J mol−1)

R :

universal gas constant (8.314 J mol−1 K−1)

T :

the absolute temperature (K)

The plots of kobs versus T and the Arrhenius plots of ln kobs versus T−1 are shown in Fig. 5. There was a well linear relationship between ln kobs and T−1 (R2 > 0.999). The apparent activation energy of the degradation of OG and RhB by alkali-activated H2O2 was then calculated to be 44.48 kJ mol−1 and 47.37 kJ mol−1, respectively. It was important to notice that the apparent activation energy of the thermal degradation of organic pollutants was usually more than 60 kJ mol−1 (Chen and Zhu 2007). This implied that the degradation of two dyes in aqueous solution by alkali-activated H2O2 was quite easy to achieve.

Fig. 5
figure 5

Effect of reaction temperature on the degradation of a OG and b RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM

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Identification of the primary reactive species

Several research groups have reported the decomposition of H2O2 under alkaline conditions for in situ generation of a variety of reactive oxygen species, including ·OH (Bokare and Choi 2014), O2·− (Li et al. 2018), and 1O2 (Zhou et al. 2015), which were reactive for pollutant degradation (Eqs. (3)–(14)). To identify whether the abovementioned reactive oxygen species were the primary oxidants for the degradation of OG and RhB in alkali-activated H2O2, ascorbic acid, a common radical scavenger, was added (Zhou et al. 2013). As shown in Fig. 6, the degradation of OG was completely inhibited with the addition of 72 mM ascorbic acid whereas the degradation of RhB was only slightly inhibited. It could be inferred that the abovementioned reactive oxygen species (i.e., ·OH, O2·−, and/or 1O2) were the primary oxidants for OG degradation, but not for RhB degradation. On the contrary, HO2−, the product of deprotonation of H2O2 (Eq. (3)), was the primary oxidant responsible for RhB degradation by alkali-activated H2O2.

Fig. 6
figure 6

Influence of ascorbic acid on the degradation of a OG and b RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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Moreover, methanol, t-butanol, and isopropyl alcohol were commonly recognized powerful ·OH scavengers, having the second-order rate constants of 9.7 × 108M−1s−1, 6 × 108M−1s−1, and 1.9 × 109M−1s−1, respectively with ·OH (Li et al. 2018; Zhou et al. 2013; Zhou et al. 2015). Figure 7 and Fig. S5 showed the inhibition of methanol, t-butanol, and isopropyl alcohol on the degradation of OG and RhB. Obviously, adding these three ·OH scavengers did not affect the degradation of OG and RhB at all, indicating that ·OH was not the primary oxidant for the degradation of OG and RhB by alkali-activated H2O2.

Fig. 7
figure 7

Influence of t-butanol on the degradation of a OG and b RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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Furfuryl alcohol was generally assumed as an effective quencher for 1O2 with the rate constant of 1.2 × 108M−1s−1 (Zhou et al. 2015). As demonstrated in Fig. 8, adding furfuryl alcohol did not affect OG and RhB degradation at all, indicating that 1O2 was also not the primary oxidant for the degradation of OG and RhB by alkali-activated H2O2.

Fig. 8
figure 8

Influence of furfuryl alcohol on the degradation of a OG and b RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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NBT had been reported for detecting O2·− because it did not react with other active oxygen species except O2·− and ·OH (Peng et al. 2016; Xu et al. 2011b). As it had been demonstrated that ·OH did not play significant role in the degradation of OG and RhB (Fig. 7 and Fig. S5), NBT was used to identify O2·−.

Figure 9a reveals the inhibitory effect on OG degradation with the addition of NBT, which was the same with the addition of ascorbic acid. These results confirmed that O2·− was the primary oxidant responsible for OG degradation by alkali-activated H2O2. Meanwhile, the characteristic UV-Vis spectrum was used to confirm the generation of O2·− by alkali-activated H2O2. The method was based on the reaction that O2·− could react with NBT to produce diformazan which has the characteristic absorption peak at 560 nm (Chen et al. 2010). Experimental results are displayed in Fig. 10. The purple diformazan was not found in control experiments with NBT and NaOH. However, in the test with the addition of H2O2, the increased absorbance between 400 and 800 nm with the reaction proceeding indicated the formation of diformazan during the degradation process. Consequently, O2·− was indeed generated by alkali-activated H2O2 and played a vital role in OG degradation. Nevertheless, Fig. 9b showed that no obvious inhibition effect on RhB degradation was observed with the addition of NBT, which was also the same with the addition of ascorbic acid. These results implied that O2·− was involved in RhB degradation by alkali-activated H2O2, but the contribution of O2·− to the degradation of RhB was negligible in comparison with the contribution of HO2−. Moreover, although carbon tetrachloride and p-benzoquinone had been widely applied as O2·− scavengers to confirm the presence of O2·− (Fu et al. 2015; Li et al. 2018), it was inappropriate to use them to identify O2·− because of the low solubility of carbon tetrachloride and the catalytic ability of p-benzoquinone (Allian et al. 1994).

Fig. 9
figure 9

Influence of NBT on the degradation of a OG and b RhB by alkali-activated H2O2. Experimental conditions: [OG]initial = 0.2 mM or [RhB]initial = 0.2 mM, pH = 11.5 or 13.0, [H2O2]initial = 36 mM, temperature (T) = 25 ± 2 °C

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Fig. 10
figure 10

UV-Vis spectra obtained from the reaction of NBT and H2O2 under alkaline condition. Experimental conditions: [NBT]initial = 0.6 mM, [H2O2]initial = 36 mM, pH = 11.5, temperature (T) = 25 ± 2 °C

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In summary, although all the oxidants of HO2−, ·OH, O2·−, and 1O2 could be generated by alkali-activated H2O2 via Eqs. (3)–(14) (Buxton et al. 1988; Christensen et al. 1982; Cui et al. 2017; Fu et al. 2017; Miao et al. 2015; Qiang et al. 2002), the primary oxidants responsible for the degradation of OG and RhB were different. According to the significant inhibition of ascorbic acid and NBT and the negligible inhibition of methanol, t-butanol, isopropyl alcohol, and furfuryl alcohol on OG degradation, the primary oxidant responsible for OG degradation was confirmed as O2·−. On the contrary, the primary oxidant responsible for RhB degradation was confirmed as HO2− based on the effect of solution pH and the negligible inhibition of ascorbic acid, methanol, t-butanol, isopropyl alcohol, NBT, and furfuryl alcohol on RhB degradation. The difference for the primary oxidant could be mainly attributed to the chemical characteristics of OG and RhB (Tab. S1). Specifically, OG was a typical azo dye while RhB was a common triphenylmethane dye. Qi et al. (Qi et al. 2016) had reported that the degradation of AO7 was initiated by the breakdown of azo bond due to the oxidative attack of O2·− that generated by alkali-activated PMS. Moreover, in another study reported by Li et al., the degradation of AO7 in the process of the activation of hydrogen peroxide with carbonate was also supposed to be the breakdown of azo bond by the generated O2·−, rather than HO2− (Li et al. 2018). Consequently, it could be concluded that O2·−, rather than HO2−, was sensitive for the breakdown of azo bond, and was also mainly responsible for the degradation of the azo dye of OG in the process of alkali-activated H2O2. Earlier, Katafias et al. had reported that methylene blue, a phenothiazine dye containing four N-methyl groups, was decolorized by the successive nucleophilic attack of HO2− on the N-methyl group in the process of alkali-activated H2O2 (Katafias et al. 2010). Since RhB was a common triphenylmethane dye containing four N-ethyl groups at either side of the xanthene ring, it could be inferred that the deprotonated HO2−, which dominated at strongly alkaline pH solutions, could be sensitive for the cleavage of N-ethyl group, and was also mainly responsible for the degradation of the triphenylmethane dye of RhB in the process of alkali-activated H2O2. Therefore, it was very meaningful to investigate the degradation of other different dyes by alkali-activated H2O2 in our future study.

Conclusions

In this work, the process of alkali-activated H2O2 was used for degrading the typical azo dye of OG and the common triphenylmethane dye of RhB in aqueous solution. The optimum solution pH for degrading of OG and RhB were pH 11.5 and pH 13.0, respectively. Moreover, the degradation of OG and RhB was significantly accelerated with the increasing of H2O2 concentration and reaction temperature. The major reactive oxidant for the degradation of OG by alkali-activated H2O2 was confirmed as O2·−, while the HO2− anion was the primary oxidant for the degradation of RhB.