## 3 Results and Discussion
\begin{table}
\begin{tabular}{c c c c c c c} \hline \hline \multirow{2}{*}{
\begin{tabular}{c} **c\(\mathbf{c_{2}O_{2}}\)** \\ **(\(\mu\)g/L)** \\ \end{tabular} } & \multicolumn{2}{c}{**Kinetic Data for MB**} & \multicolumn{2}{c}{**Kinetic Data for MO**} \\ \cline{2-7}  & **Total Fe** & \multicolumn{2}{c}{**Rate Constants**} & \multicolumn{2}{c}{**Total Fe**} & \multicolumn{2}{c}{**Rate Constants**} \\ \cline{2-7}  & **Total Fe** & \multicolumn{2}{c}{**Rate Constants**} & \multicolumn{2}{c}{**Total Fe**} & \multicolumn{2}{c}{**Rate Constants**} \\ \cline{2-7}  & **(mg/L)** & **\(k_{1}\) (h\({}^{-1}\))** & **\(k_{2}\) (h\({}^{-1}\))** & **(mg/L)** & **\(k_{1}\) (h\({}^{-1}\))** & **\(k_{2}\) (h\({}^{-1}\))** \\ \hline
333 & 0.94 & 0.35 (0.892) & 1.36 (0.996) & 0.98 & 0.20 (0.936) & 0.35 (0.983) \\
666 & 1.36 & 0.10 (0.951) & 0.38 (0.995) & 1.32 & 0.23 (0.984) & 0.36 (0.990) \\
999 & 2.96 & 0.15 (0.980) & 0.24 (0.992) & 3.04 & 0.10 (0.989) & 0.33 (0.999) \\
1665 & 3.68 & 0.12 (0.996) & 0.30 (0.999) & 3.95 & 0.25 (0.988) & 0.28 (0.999) \\ \hline \hline \end{tabular}
\end{table} 
Table 2: Rate constants of the photo-Fenton process as a function of S\({}_{2}\)O\({}_{8}\)\({}^{2-}\) concentration. Figure 3: Pseudo-first order kinetic of photodegradation for (**a**): MB in the presence of H\({}_{2}\)O\({}_{2}\); (**b**): MO in the presence of H\({}_{2}\)O\({}_{2}\); (**c**): MB in the presence of S\({}_{2}\)O\({}_{8}\)\({}^{2-}\); (**d**): MO in the presence of S\({}_{2}\)O\({}_{8}\)\({}^{2-}\). Figure 3a-d indicate that the photo-Fenton degradation mechanism of both dyes occurs in two steps. The first part of the kinetic (rate constant, k1) is characterized by a slower overall degradation of the dyes for both H2O2 and S2O82, because in this step the oxidation of metallic iron Fe0 to Fe2+ and the generation of hydroxyl and sulfate radicals take place; thus, initially there is a low amount of oxidant species to initiate the photodegradation. The generation of reactive HO\(\bullet\) radicals in the presence of H2O2 and Fe2+ ions from the catalyst is depicted by the following Mechanism (4) [32; 33]:  
\[ \begin{split}\text{Fe}^{0}+\text{H}_{2}\text{O}_{2}\overset{h \
u}{=}\text{Fe}^{2+}+2\text{OH}^{-}\\ \text{Fe}^{2+}+\text{H}_{2}\text{O}_{2}\overset{h\
u}{=}\text{Fe}^{3+ }+\text{OH}^{-}+\text{HO}\bullet\\ \text{Fe}^{0}+2\text{Fe}^{3+}\overset{h\
u}{=}3\text{Fe}^{2+} \end{split}  \tag{4}\]  
The persulfate ions can decompose in a first slow, ongoing step to highly reactive sulfate radicals (SO4-\(\bullet\)), which oxidize the metallic iron to produce Fe2+ ions (i.e., the iron is \"activated\"), which can further initiate the production of sulfate radicals on their own. Avoiding surface passivation in this step is crucial for achieving a high photodegradation yield [32]. The sulfate radicals can react with water to generate hydroxyl radicals, so for the persulfate oxidant, the dyes can be oxidized by both type of radicals [29; 34; 35], according to the following Mechanism (5):  
\[ \begin{split}\text{S}_{2}\text{O}_{8}^{2-\overset{h\
u}{=}}& 2\text{SO}_{4}^{-}\bullet\\ 2\text{SO}_{4}^{-}\bullet+\text{Fe}^{0}+2\text{H}_{2}\text{O} \overset{h\
u}{=}\text{Fe}^{2+}+2\text{H}\text{SO}_{4}^{-}+2\text{HO}^{-}\\ \text{Fe}^{2+}+\text{S}_{2}\text{O}_{8}^{2-\overset{h\
u}{=}} \text{Fe}^{3+}+\text{SO}_{4}^{-}\bullet+\text{SO}_{4}^{2-}\\ \text{SO}_{4}^{-}\bullet+\text{H}_{2}\text{O}\overset{h\
u}{=} \text{HO}\bullet+\text{H}\text{SO}_{4}^{-}\end{split}  \tag{5}\]  
Regarding the dye photodegradation in this step, these could be oxidized to higher molecular mass intermediates (possibly hydroxylated compounds or carboxylic acids) according to Mechanism (6), where DyeH denotes the methylene blue/methylorange dyes in their non-oxidized form. The dyes can also be directly oxidized by Fe3+ ions [35; 36]. \[