## 3 Results and Discussion
### Comparison of Hydroxyl Radical Formation Rates of Various AOPs  
Figure 4 shows the concentration of formaldehyde produced by the reaction of DMSO and hydroxyl radical.  
Additionally, the quantitative results for the hydroxyl radicals, calculated through stoichiometry and using the concentration of formaldehyde detected, are presented. GC-ECD analysis was performed to quantitatively determine the presence of formaldehyde. Ten minutes after the experiment started, the concentration of formaldehyde detected was 0.16 mg/L, while after 2 h 1.33 mg/L of formaldehyde was detected. This indicates that 0.03 mol and 0.27 mol of hydroxyl radicals were produced after 10 min and after 2 h, respectively. The detected hydroxyl radicals might have originated from the recombination of plasma-induced reactive chemicals, as the half-life of the hydroxyl radical is very short (\(\sim\)10\({}^{-9}\) s) [77].  
Figure 5 presents the results of calculating, from the amount of hydroxyl radical produced by CP, the instantaneous formation rates of hydroxyl radicals at the measured time intervals.  
After 10 min of the plasma reaction, the hydroxyl radical instantaneous formation rate was calculated to be 5.2 \(\times\) 10\({}^{-5}\) mol/s, but it was subsequently confirmed that this rate gradually decreased to 3.7 \(\times\) 10\({}^{-5}\) mol/s at 2 h. This decrease is estimated to be due to the fact that even the produced formaldehyde acts as a scavenger for plasma reactions over time and that some of the formaldehyde was oxidized.  
Many studies have demonstrated that formic acid is formed as an intermediate due to the oxidation reaction of plasma with formaldehyde [78]. Additionally, the reactions related to formaldehyde decomposition are shown in Equations ((10)-(13)) [79; 80; 81]. The degradation of formaldehyde due to its plasma reaction depends on two mechanisms, including (a) the direct degradation caused by the collision of electrons with formaldehyde  
Figure 4: Detection of formaldehyde and quantification of hydroxyl radicals. The left arrow denotes formaldehyde concentration, and the right arrow indicates hydroxyl radicals.  
and (b) the reaction between formaldehyde molecules and various oxidizing species [80]. Finally, formic acid degrades to end products including H2O, CO2, etc. [79]  
\[\text{HCHO}+\text{e}^{-}\rightarrow\text{:H}+\text{CHO}^{-} \tag{10}\]  
\[\text{HCHO}+\text{O}_{3}\rightarrow\text{H}_{2}\text{O}+\text{CO}_{2}+\text{O} \tag{11}\]  
\[\text{HCHO}+\text{\cdot OH}\rightarrow\text{H}_{2}\text{O}+\text{CHO} \tag{12}\]  
\[\text{HCHO}+\text{\cdot OH}\rightarrow\text{HCOOH}+\text{H} \tag{13}\]  
\[\text{HCOOH}+\text{\cdot OH}\rightarrow\text{H}_{2}\text{O}+\text{H}+\text{CO}_{2} \tag{14}\]  
Using the 10 min measurement time, the hydroxyl radical instantaneous formation rates of CP and other AOPs are compared and summarized, as shown in Table 6.  
It was confirmed that CP generated more hydroxyl radicals in a shorter time period compared to other AOPs. These rapidly generated hydroxyl radicals are able to efficiently decompose organic pollutants. This is why CP was able to efficiently degrade PFOS within a short period of time.  
\begin{table}
\begin{tabular}{c c c} \hline \hline
**Process** & **Instantaneous Formation Rate (mol/s)** & **Reference** \\ \hline UV-Fe\({}^{2+}\)/H\({}_{2}\)O\({}_{2}\) & 3.0 \(\times\) 10\({}^{-10}\) & [75] \\ \hline UV-H\({}_{2}\)O\({}_{2}\) & 1.1 \(\times\) 10\({}^{-10}\) & [75] \\ \hline Photocatalytic oxidation (Fe/TiO\({}_{2}\)) & 4.3 \(\times\) 10\({}^{-8}\) & [82] \\ \hline Fenton (FeOOH) & 4.2 \(\times\) 10\({}^{-5}\) & [83] \\ \hline Aeration-assisted CP & 5.2 \(\times\) 10\({}^{-5}\) & This study \\ \hline \hline \end{tabular}
\end{table}
Table 6: Instantaneous formation rates of various AOPs.  
Figure 5: CP-enabled instantaneous formation rates of hydroxyl radical.