Electro-Fenton Method Oxidation of Salicylic Acid in Aqueous Solution with Graphite Electrodes

Abstract

Electro-Fenton (EF) oxidation of salicylic acid (SA) from an aqueous solution was studied in a continuous stirred tank reactor with the purpose of optimizing the process parameters on a laboratory scale. A laboratory-scale continuous-flow rectangular reactor, with a predetermined 3 L liquid volume of SA in the reactor, was used for this study. Electrolytic reactions carried out for 100 mg/L of SA solution, the parameters of catalyst concentration, pH, voltage, flow rate (residence time), and electrode spacing were varied. Electrolysis was carried out for a period of 10 h to optimize these parameters. A maximum of 70% oxidation was obtained at an initial Fe²⁺ concentration of 5 mg/L, pH of 2.5, voltage of 2.5 V, flow rate of 10 mL/min, and electrode spacing of 3 cm, after 10 h of electrolysis. Results obtained show the feasibility of this continuous EF process for the oxidation of SA.

Introduction

Salicylic acid (SA) has been identified as a water pollutant, which originates from the manufacturing activities of paper milling, cosmetic industries, and as a landfill leachate (Tian et al., 2009). Although SA is largely produced from phenol, it can also be obtained from various natural sources, including trees like willow, birch, and myrtle (Hanesian and Perna, 2000). Although SA is known for its superb ability to ease aches and pains and reduce fevers, it has proven to be very dangerous to human beings at a high concentrations (Wang et al., 2012). SA reacts with proteins, and can stimulate skin and mucous membrane, resulting in various health issues like tinnitus, qualm, naupathia, and electrolytic turbulence (Chen et al., 2010). It is known that SA can cause fetal abnormalities and central nervous system depression if swallowed, inhaled, or absorbed through the skin (Tian et al., 2009). Wastewater containing SA originates mainly from the industries synthesizing SA and its derivatives (Goi et al., 2008). These multiple sources of SA contamination call for faster development of innovative water purification technologies.

Recently, much attention has been paid to separate the source of the refractory or toxic effluent and treat it by advanced oxidation processes (AOPs) (Mohamed et al., 2009). Recently, many studies have been concentrated on the possibility to improve an economic aspect of these AOPs. Among AOPs, the Fenton process procured great attention in recent years (Balci et al., 2009). The major reaction mechanism of the Fenton process can be described by means of the following reactions [Eqs. (1 )–(5)] (Pignatello et al., 2006): \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{2 + } + { \rm H}_2{ \rm O}_2 \rightarrow { \rm Fe}^{3 + } + { \rm OH}^{ - } + { \rm HO}^{ \bullet} \tag{1}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{2 + } + { \rm OH}^{ \bullet} \rightarrow { \rm Fe}^{3 + } + { \rm OH}^{ - } \tag{2}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{3 + } + { \rm H}_2{ \rm O}_2 \rightarrow { \rm Fe}^{2 + } + { \rm HO}_2^{ \;\;\bullet} + { \rm H}^{ + } \tag{3}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm OH}^{ \bullet} + { \rm organics} \rightarrow { \rm products} \tag{4}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2{ \rm O}_2 + { \rm OH}^{ \bullet} \rightarrow { \rm H}_2{ \rm O} + { \rm HO}_2^{ \;\;\bullet} \tag{5}\end{align*} \end{document}

The main drawbacks of the Fenton process are the use of a large quantity of chemical reagents, large production of ferric hydroxide sludge, and very slow catalysis of the ferrous ion generation (Özcan et al., 2008a). The electrochemical technique is one of the promising schemes that has been currently applied to eliminate or lessen the disadvantages of AOPs (Anotai et al., 2006; Brillas et al., 2009; Nidheesh and Gandhimathi, 2012). The main advantages of electrochemical techniques over other AOPs are its low operating cost and high mineralization efficiency (Oturan, 2000; Oturan et al., 2001). The electro-Fenton (EF) process is one of such electrochemical approaches (Harrington and Pletcher, 1999).

In the EF process, hydroxyl radicals generate continuously in the electrolytic system by the reaction between externally added ferrous or ferric ions and in situ generated hydrogen peroxide. H₂O₂ can be electrogenerated by the reduction of dissolved oxygen and Fe²⁺ by the reduction of Fe³⁺ [Eqs. (6) and (7)] (Qiang et al., 2003). Simultaneous Fe²⁺ regeneration will occur at the cathode according to the following equations [Eqs. (3), (8), and (9)] (Oturan et al., 2010; Nidheesh et al., 2013). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm O}_2 + 2{ \rm H}^{ + } + 2{ \rm e}^{ - } \rightarrow { \rm H}_2{ \rm O}_2 \tag{6}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{3 + } + { \rm e}^{ - } \rightarrow { \rm Fe}^{2 + } \tag{7}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{3 + } + { \rm HO}_2^{ \;\;\bullet} \rightarrow { \rm Fe}^{2 + } + { \rm H}^{ + } + { \rm O}_2 \tag{8}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{3 + } + { \rm R}^{ \bullet} \rightarrow { \rm Fe}^{2 + } + { \rm R}^{ + } \tag{9}\end{align*} \end{document}

However, a part of ^•OH is lost by its direct reaction with Fe²⁺ [Eq. (2)] and with H₂O₂ [Eq. (5)] (Buxton et al., 1988; Sun and Pignatello, 1993). The rate of the Fenton reaction [Eq. (2)] can also decay by the oxidation of Fe²⁺ in the medium with HO₂^• [Eq. (10)] (Kwan and Voelker, 2002), as well as at the anode from the reaction [Eq. (11)], when an undivided cell is used. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{2 + } + { \rm HO}_2^{ \;\;\bullet} \rightarrow { \rm Fe}^{3 + } + { \rm HO}_2^{ - } \tag{10}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Fe}^{2 + } \rightarrow { \rm Fe}^{3 + } + { \rm e}^{ - } \tag{11}\end{align*} \end{document}

Cathode material has a major role in the efficiency of the EF process. A wide variety of electrode materials such as carbon felt electrodes (Sirés et al., 2007; Oturan et al., 2008), graphite (Zhang et al., 2008; Nidheesh and Gandhimathi, 2013a, 2013b) or gas diffusion (Boye et al., 2003), carbon sponge (Özcan et al., 2008a), and activated carbon fiber (Wang et al., 2008) have been used as electrodes in various EF studies. In this study, the effectiveness of graphite electrode on the percentage of SA oxidation was investigated.

Usually, most of the wastewater treatment applications are conducted in a continuous mode (Rosales et al., 2009). However, very few studies have been conducted in a continuous mode for EF studies. To the best of our knowledge, there is no application of continuous mode EF technology in the oxidation of SA. Therefore, this study is conducted to find the effect of continuous mode EF technology in the oxidation of SA. The experiments were carried out in a laboratory scale for the synthetically prepared SA solution. The percentage of oxidation was evaluated in view of various operational parameters such as catalyst concentrations, pH, voltage, flow rate, and distance between sacrificial electrodes.

Materials and Methods

Materials

All chemicals used were analytical grade. The 99.8% pure SA (C₇H₆O₃) was purchased from Merck. FeSO₄·7H₂O (Merck) was used as the source of Fe²⁺ catalyst in EF studies. The Orion EA940 Ion analyzer (Thermo) was employed for pH measurements and, pH of the working solution was adjusted using H₂SO₄ (Merck). Graphite plates were purchased from Anabond Sainergy Fuel Cell India Private Ltd., Chennai, India. DC power supply (0–5 A and 0–30 V) was used as the source of direct current. SA oxidation was determined by ultraviolet–visible light spectrophotometer (Lambda 25; PerkinElmer) using a quartz cell of path length 1 cm at 298 nm wavelength.

Experimental procedure

Electrolysis of the SA solution was carried out in a 30 cm×15 cm single-compartment electrolytic flow cell of 3000 mL working volume (Fig. 1). The electrolyte was prepared by dissolving 100 mg/L SA and 5 mg/L Fe²⁺ in distilled water. The initial pH of the solution was adjusted with 2 N (normality) H₂SO₄. An inlet tank was provided to hold the untreated SA solution to be passed into the system and the influent was passed through a 1-cm-diameter pipe. A flow control valve was used to control the flow of water entering the reactor. The experimental setup is shown in Fig. 1. Electrolyses were performed with DC power supplies using graphite electrodes under controlled electrolysis conditions. Six graphite plates of geometric area 40 cm² were used as anodes and cathodes. Anodes and cathodes were arranged vertically as well as parallel with equal interelectrode spacing. Oxygen was supplied to the bottom of the cathodes by three fish aerators. The treated SA solution was collected in an outlet tank. During the experiments, samples were drawn from the reactor at predetermined time intervals and analyzed in a UV visible spectroscope.

FIG. 1.

Experimental setup for continuous study. 1, cathode; 2, anode; 3, DC power supply; 4, aerator; 5, collection tank; 6, inlet tank; 7, outlet; 8, inlet; 9, electrotytic cell.

Results and Discussion

Effect of Fe²⁺ concentration

Fe²⁺ concentration is an important parameter that has great control over EF reactions (Ramirez et al., 2005). Since the generation of hydrogen peroxide depends on the applied voltage, cathode material, solution pH, etc., the efficiency of the EF process depends more on ferrous ion concentrations. Excess or shortage of any of these two reagents results in the occurrence of scavenging reactions [Eqs. (5) and (2)] (Tang and Huang, 1997; Lopez et al., 2004). Investigations have shown that best oxidation is attained [Eq. (12)] when neither H₂O₂ nor Fe²⁺ is overdosed to make maximum hydroxyl radicals available for the oxidation of organic compounds (Pignatello et al., 2006; Brillas et al., 2009). In the EF process, the electrolytic generation rate of H₂O₂ at constant current density, inner electrode spacing, solution pH, and electrode area is constant. Therefore, the generation of hydroxyl radical and thus the efficiency of the EF process depend more on ferrous ion concentrations. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm OH}^{ \bullet} + { \rm RH} \rightarrow { \rm H}_2{ \rm O} + { \rm R}^{ \bullet} \rightarrow \hbox{further oxidation} \tag{12}\end{align*} \end{document}

To find the effect of FeSO₄ catalyst concentrations on SA oxidation in a continuous stirred tank reactor (CSTR) rector, electrolysis was carried out for a period of 10 h by varying the Fe²⁺ concentration from 2.5 to 15 mg/L. A 100 mg/L SA solution with pH of 2.5, electrode distance of 3 cm, flow rate of 10 mL/min, and voltage of 2.5 V was used for all the experiments. The results are shown in Fig. 2: 41%, 57%, 43%, and 37% oxidation was obtained for 2.5, 5, 10, and 15 mg/L catalyst concentrations after 10 h of electrolysis, respectively. As can be seen from Fig. 2, up to 5 mg/L, the percentage of SA oxidation increased with an increase in catalyst concentrations. This is due to the increase in production of ^•OH at the surface of the cathode [Eq. (1)] with increase in ferrous ion concentrations leading to the enhancement in SA oxidation with time. Oturan et al. (2000) reported that 0.5 moles of oxygen is required for two moles of hydroxyl radical in the EF system. Since the production of hydroxyl radical depends on the concentration of hydrogen peroxide, very less concentration of oxygen molecules is required for the production of hydrogen peroxide. This is mainly due to the excess production of oxygen molecules at the anode by water oxidation (Oturan et al., 2000). Therefore, at lower concentration of ferrous ions, the excess of hydrogen peroxide in the system will cause hydroxyl radical scavenging reactions as in Equation (5) (Panda et al., 2011).

FIG. 2.

Effect of initial Fe²⁺ concentrations on SA oxidation. Solution pH: 2.5; inner electrode spacing: 3 cm; inlet flow rate: 10 mL/min; applied voltage: 2.5 V. SA, salicylic acid.

However, a further increase in catalyst concentrationss decreased the percent oxidation of the system. This is due to the fact that excess Fe²⁺ molecules compete with the SA molecules for the available ^•OH [Eq. (2)] (Benitez et al., 2001) and thus decreases the oxidation efficiency. Another cause for this is the deposition of Fe³⁺ ions on the surface of the cathode that will inhibit the electro generation of H₂O₂ (Wang et al., 2010). Thus, a suitable concentration of Fe²⁺ was an important parameter in the EF reaction. A concentration of 5 mg/L of Fe²⁺ was selected for further experiments.

Effect of pH

Efficiency of the EF process depends more on solution pH than other key operating parameters, because H₂O₂ production and relative concentration of ferrous ion in the Fenton electrolytic system depends more on solution pH. Equation (6) indicates that pH value of solution is one of the important factors affecting H₂O₂ generation because it converts the dissolved oxygen to hydrogen peroxide by consuming the protons in the electrolyte at acidic conditions (Mingyu et al., 2005).

The effect of pH was investigated by treating a 100 mg/L SA solution with 5 mg/L Fe²⁺ of catalyst concentration, 3 cm electrode distance, 10 mL/min flow rate, and 2.5 V voltage by varying the pH values from 2 to 3.5 in a CSTR reactor. The results are shown in Fig. 3. The percent oxidation that corresponds to pH values 2, 2.5, 3, and 3.5 was 59%, 70%, 57%, and 25%, respectively. When the pH decreased to 2.0 from 3, SA oxidation efficiency of the EF process was decreased. Two mechanisms might explain such a trend. First, the formation of oxonium ions [Eq. (13)] led to the deactivation of H₂O₂, which in turn reduced the production of hydroxyl radicals (Zhou et al., 2007). Second, increase in the scavenging effect of H⁺ ions and hydroxyl radicals led to the reduction of SA oxidation (Tang and Huang, 1997). It was observed that, the SA oxidation efficiency reduced when the pH of the solution was increased from 2.5. This is mainly due to the formation of excess ferric hydroxo complexes at higher pH values. This leads to decrease in ferrous ion concentrations in the solution and decrease in the Fenton reaction frequency. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2{ \rm O}_2 + { \rm H}^{ + } \rightarrow { \rm H}_3{ \rm O}_2^{ + } \tag{13}\end{align*} \end{document}

FIG. 3.

Effect of pH on SA oxidation. Inner electrode spacing: 3 cm; ferrous concentration: 5 mg/L; inlet flow rate: 10 mL/min; applied voltage: 2.5 V.

Most of the studies reported that the Fenton process has an optimal at pH of 3 (Nidheesh et al., 2013). For SA oxidation, a significant difference has been observed between pH 3 and 2.5. This may be due to the difference in speciation of ferrous ions at various pH values. Wells and Salam (1965, 1968) reported that the majority of ferrous ions will be in the form of Fe²⁺ at pH less than 3. This helps the rapid degradation of SA at pH 2.5, by the Fenton reaction. Complex formation of SA with ferric ion is a well-known phenomenon. This complex formation has an optimal pH around 2.6–2.8 (Mehlig, 1938). Therefore, at pH 3, the amount of complex formation with ferric ions (produced from Fenton reaction) and SA is higher compared with pH 2.5. This reduces the effective concentration of iron available for EF at pH 3 and results in less SA oxidation rate.

Effect of voltage

Voltage is another important parameter influencing the efficiency of EF reactions. It can influence the generation of H₂O₂ [Eq. (6)], which in turn enhances the production of hydroxyl radicals [Eq. (1)] (Panizzaa and Oturan, 2011). To study the effect of voltage on the oxidation of SA, experiments were conducted at different voltages varying from 1.5 to 3.5 V. A catalyst concentration of 5 mg/L, electrode spacing of 3 cm, flow rate of 10 mL/min, and a pH of 2.5 was maintained for all the experiments. The results are shown in Fig. 4. Oxidation efficiency increased from 28% to 70% with the increase in voltage from 1.5 to 2.5 V. This enhancement of oxidation is due to the production of greater amounts of ^•OH from Fenton reaction [Eq. (1)] by the faster production of H₂O₂ at the cathode [Eq. (6)] (Ruiz et al., 2011). The increase of the applied current enhances hydrolysis of Fe³⁺ ions and formation of Fe(OH)₃(s), which consequently inhibits both the regeneration of Fe²⁺ ions and the production of OH^• radicals (Qiang et al., 2003; Hsueh et al., 2005). The formation of ferric–hydroxyl complexes not only decreases iron catalyst activity and Fe²⁺ regeneration, but also retards dissolved oxygen mass transfer and H₂O₂ formation by its presence in the solution and formation of a coating on the electrode surface (Zhang et al., 2009). On the other hand, it was observed that the percent oxidation was dropped dramatically from 70% to 39% when the applied voltage increased from 2.5 to 3.5 V (Fan et al., 2010). The reduction in SA oxidation is mainly due to the increase in hydrogen gas evolution at cathode as in Equation (14), formation of water as in Equation (15), and hydrogen peroxide decomposition (Zhang et al., 2006; Özcan et al., 2008b). In addition, the temperature of the solution increases with an increase in applied voltage. This decreases the solubility of oxygen and enhances the H₂O₂ decomposition rate in the electrolytic system (Özcan et al., 2008a). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}2{ \rm H}^{ + } + 2{ \rm e}^{ - } \rightarrow { \rm H_2} \tag{14}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}0.5{\rm O}_2 + 2{ \rm H}^{ + } + 2{ \rm e}^{ - } \leftrightarrow { \rm H}_2{ \rm O} \tag{15}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}2{ \rm H}_2{ \rm O}_2 \rightarrow 4{ \rm H}^{ + } + { \rm O}_2 + 4{ \rm e}^{ - } \tag{16}\end{align*} \end{document}

FIG. 4.

Effect of voltage on SA oxidation. Inner electrode spacing: 3 cm; ferrous concentration: 5 mg/L; inlet flow rate: 10 mL/min; solution pH: 2.5.

Effect of flow rate

The effect of flow rate on the oxidation of SA was examined by electrolyzing 100 mg/L of SA solution of pH 2.5 at 5, 10, and 15 mL/min. The corresponding SA oxidations are depicted in Fig. 5. As can be seen in Fig. 5, the oxidation corresponding to 5, 10, and 15 mL/min flow rates are 75.00%, 70.29%, and 56.65%, respectively. The experimental results show that the oxidation of SA from an aqueous solution increases with a decrease in the flow rate. Percent oxidation of SA was found to be high for very slow flow rates, since the residence time of the electrolyte solution inside the electrolytic cell was very high. Thus, for low residence time of the electrolyte solution inside the cell, the percentage oxidation of SA was found to be decreased with high flow rates (Prabhakaran et al., 2009). To balance the treatment capacity and oxidation, a 10 mL/min flow rate was selected for further experiments.

FIG. 5.

Effect of flow rate on SA oxidation. Solution pH: 2.5; inner electrode spacing: 3 cm; ferrous concentration: 5 mg/L; applied voltage: 2.5 V.

Effect of electrode spacing

To find the effect of electrode spacing on SA oxidation in a CSTR reactor, electrolysis was carried out for a period of 10 h by varying the electrode space from 2 to 5 cm. A catalyst concentration of 5 mg/L, voltage of 2.5 V, flow rate of 10 mL/min, and a pH of 2.5 were maintained for all the experiments. The results are shown in Fig. 6. It was found that the electrode spacing significantly affect the efficiency of this system. The oxidation of SA was 43%, 70%, 44%, and 40% when the electrode spacings were 2, 3, 4, and 5 cm, respectively. The oxidation of SA increased with in electrode spacing from 2 to 3 cm. Such promotion could be attributed to the enhancement of ion transfer at short electrode spacing, which resulted in the rapid regeneration of Fe²⁺ ions at the cathode surface. Further, it was observed that the percent oxidation was reduced from 70% to 40% when the electrode spacing increased from 3 to 5 cm. The negative effect of the higher electrode spacing on the SA oxidation can be explained by the decrease of the mass transfer of Fe³⁺ions between electrodes, which is necessary for the regeneration of Fe²⁺ ions [Eq. (7)] (Fockedey and Lierde, 2002). The rate of electro regeneration of Fe²⁺ decreases with the increase in inner electrode gap due to the reduction of the mass transfer of Fe³⁺ to the cathode surface (Zhang et al., 2006), leading to the reduction in the rate of generation of hydroxyl radical. Electro-regenerated Fe²⁺ will get oxidized to Fe³⁺ at a shorter electrode distance as in Equation (11) (Zhang et al., 2006). Fenton reactions get inhibited by this oxidation and result in a decrease in the SA oxidation rate. In addition, consumption of current was increased with a decrease in electrode spacing due to scavenging reactions as given in Equation (11), leading to the increase in cost of operation.

FIG. 6.

Effect of electrode spacing on SA oxidation. Applied voltage: 2.5 V; solution pH: 2.5; ferrous concentration: 5 mg/L; inlet flow rate: 10 mL/min.

Conclusions

The oxidation of SA in an aqueous medium by the EF process in a CSTR under various operating conditions was investigated. Continuous experiments in saturated air conditions determined that the optimal values for the oxidation of SA were 5 mg/L Fe²⁺, pH 2.5, voltage 2.5 V, and electrode spacing 3 cm. Under these optimal conditions, a maximum SA oxidation of 70% was obtained. It is obvious that the CSTR was able to operate without operational problems, and attaining good SA oxidation depended on the residence time without the addition of other oxidizing agents.

Footnotes

Author Disclosure Statement

No competing financial interests exist.

References

Anotai

, Lub

M.C.

, and Chewpreecha

(2006). Kinetics of aniline degradation by Fenton and electro-Fenton processes. Water Res., 40, 1841.

Balci

, Oturan

, Cherrier

, and Oturan

M.A.

(2009). Degradation of atrazine in aqueous medium by electrocatalytically generated hydroxyl radicals. A kinetic and mechanistic study. Water Res., 43, 1924.

Benitez

F.J.

, Acero

J.L.

, Real

F.J.

, Rubio

F.J.

, and Leal

A.I.

(2001). The role of hydroxyl radicals for the decomposition of p-hydroxy phenylacetic acid in aqueous solutions. Water Res., 35, 1338.

Boye

, Dieng

M.M.

, and Brillas

(2003). Anodic oxidation, electro-Fenton and photoelectro-Fenton treatments of 2,4,5-trichlorophenoxyacetic acid. J. Electroanal. Chem., 557, 135.

Brillas

, Sires

, and Oturan

M.A.

(2009). Electro-Fenton process and related electrochemical technologies based on Fenton's reaction chemistry. Chem. Rev., 109, 6570.

Buxton

G.U.

, Greenstock

C.L.

, Helman

W.P.

, and Ross

A.B.

(1988). Critical review of rate constant for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Data Ref., 17, 513.

Chen

X.M.

, da Silva

D.R.

, and Martínez-Huitle

C.A.

(2010). Application of advanced oxidation processes for removing salicylic acid from synthetic wastewaters. Chin. Chem. Lett., 21, 101.

Fan

, Zhihui

, and Zhang

(2010). Design of an electro-Fenton system with a novel sandwich film cathode for wastewater treatment. J. Hazard. Mater., 176, 678.

Fockedey

, and Lierde

A.V.

(2002). Coupling of anodic and cathodic reactions for phenol electro-oxidation using three-dimensional electrodes. Water Res., 36, 4169.

10.

Goi

, Veressinina

, and Trapido

(2008). Degradation of salicylic acid by Fenton and modified Fenton treatment. Chem. Eng. J., 143, 1.

11.

Hanesian

, and Perna

A.J.

(2000). Process Analysis and Design of a Manufacturing Facility Using Hazardous Material in a Residential Community (The Manufacture of Aspirin). Workbook, New Jersey Institute of Technology, Newark, NJ.

12.

Harrington

, and Pletcher

(1999). The removal of low levels of organics from aqueous solutions using Fe(II) and hydrogen peroxide formed in situ at gas diffusion electrodes. J. Electrochem. Soc., 146, 2983.

13.

Hsueh

C.L.

, Huang

Y.H.

, Wang

C.C.

, and Chen

C.Y.

(2005). Degradation of azo dyes using low iron concentration of Fenton and Fenton-like system. Chemosphere, 58, 1409.

14.

Kwan

W.P.

, and Voelker

B.M.

(2002). Decomposition of hydrogen peroxide and organic compounds in the presence of dissolved iron and ferrihydrite. Environ. Sci. Technol., 36, 1467.

15.

Lopez

, Pagano

, Volpe

, and Di Pinto

(2004). Fenton's pre-treatment of mature landfill leachate. Chemosphere, 54, 1000.

16.

Mehlig

J.P.

(1938). Colorimetric determination of iron with salicylic acid: A spectrophotometric study. Ind. Eng. Chem. Anal. Ed., 10, 136.

17.

Mingyu

, Lin

, Yunyun

, Na

, Yuanming

, and Hua

(2005). Studies on photo-electro-chemical catalytic degradation of Acid Scarlet 3R dye. Sci. China Ser. B Chem., 48, 297.

18.

Mohamed

I.B.

, Rifaat

A.W.

, and El-Kalliny

A.S.

(2009). Fenton-biological treatment processes for the removal of some pharmaceuticals from industrial wastewater. J. Hazard. Mater., 167, 567.

19.

Nidheesh

P.V.

, and Gandhimathi

(2012). Trends in electro-Fenton process for water and wastewater treatment: An overview. Desalination, 299, 1.

20.

Nidheesh

P.V.

, and Gandhimathi

(2013a). Comparative removal of rhodamine B from aqueous solution by electro-Fenton and electro-Fenton like processes. Clean Soil Air Water 17 Oct 2013 [Epub ahead of print]; DOI: 10.1002/clen.201300093.

21.

Nidheesh

P.V.

, and Gandhimathi

(2013b). Removal of rhodamine B from aqueous solution using graphite-graphite electro-Fenton system. Desalination Water Treat. 13 May 2013 [Epub ahead of print]; DOI: 10.1080/19443994.2013.790321.

22.

Nidheesh

P.V.

, Gandhimathi

, and Ramesh

S.T.

(2013). Degradation of dyes from aqueous solution by Fenton processes: A review. Environ. Sci. Pollut. Res., 20, 2099.

23.

Oturan

M.A.

(2000). An ecologically effective water treatment technique using electrochemically generated hydroxyl radicals for in situ destruction of organic pollutants. J. Appl. Electrochem., 30, 475.

24.

Oturan

M.A.

, Edelahi

M.C.

, Oturan

, kacemi

K.E.

, and Aaron

J.J.

(2010). Kinetics of oxidative degradation/mineralization pathways of the phenylurea herbicides diuron, monuron and fenuron in water during application of the electro-Fenton process. Appl. Catal. B Environ., 97, 82.

25.

Oturan

M.A.

, Oturan

, Lahite

, and Trévin

(2001). Production of hydroxyl radicals by electrochemically assisted Fenton's reagent. Application to the mineralization of an organic micropollutant, pentachlorophenol. J. Electroanal. Chem., 507, 96.

26.

Oturan

M.A.

, Peiroten

, Chartrin

, and Acher

A.J.

(2000). Complete destruction of p-nitrophenol in aqueous medium by electro-Fenton method. Environ. Sci. Technol., 34, 3474.

27.

Oturan

, Trajkovska

, Oturan

M.A.

, Couderchet

, and Aaron

J.J.

(2008). Study of the toxicity of diuron and its metabolites formed in aqueous medium during application of the electrochemical advanced oxidation process “electro-Fenton.”. Chemosphere, 73, 1550.

28.

Özcan

, Sahin

, Koparal

A.S.

, and Oturan

M.A.

(2008a). Carbon sponge as a new cathode material for the electro-Fenton process: Comparison with carbon felt cathode and application to degradation of synthetic dye Basic Blue 3 in aqueous medium. J. Electroanal. Chem., 616, 71.

29.

Özcan

, Sahin

, Koparal

A.S.

, and Oturan

M.A.

(2008b). Degradation of picloram by the electro-Fenton process. J. Hazard. Mater., 153, 718.

30.

Panda

, Sahoo

, and Mohapatra

(2011). Decolourization of Methyl Orange using Fenton-like mesoporous Fe₂O₃–SiO₂ composite. J. Hazard. Mater., 185, 359.

31.

Panizzaa

, and Oturan

M.A.

(2011). Degradation of Alizarin Red by electro-Fenton process using a graphite-felt cathode. Electrochim. Acta, 56, 7084.

32.

Pignatello

J.J.

, Oliveros

, and MacKay

(2006). Advanced oxidation processes for organic contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev. Environ. Sci. Technol., 36, 1.

33.

Prabhakaran

, Kannadasan

, and Basha

C.A.

(2009). Treatability of resin effluents by electrochemical oxidation using batch recirculation reactor. Int. J. Environ. Sci. Tech., 6, 491.

34.

Qiang

, Chang

, and Huang

(2003). Electrochemical regeneration of Fe²⁺ in Fenton oxidation processes. Water Res., 37, 1308.

35.

Ramirez

J.H.

, Costa

C.A.

, and Madeira

L.M.

(2005). Experimental design to optimize the degradation of the synthetic dye Orange II using Fenton's reagent. Catal. Today, 107/108, 68.

36.

Rosales

, Pazos

, Longo

M.A.

, and Sanromán

M.A.

(2009). Electro-Fenton decoloration of dyes in a continuous reactor: A promising technology in colored wastewater treatment. Chem. Eng. J., 155, 62.

37.

Ruiz

E.J.

, Arias

, Brillas

, Ramírez

A.H.

, and Hernández

J.M.P.

(2011). Mineralization of Acid Yellow 36 azo dye by electro-Fenton and solar photoelectro-Fenton processes with a boron-doped diamond anode. Chemosphere, 82, 495.

38.

Sirés

, Centellas

, Garrido

J.A.

, Rodríguez

R.M.

, Arias

, Cabot

P.L.

, and Brillas

(2007). Mineralization of clofibric acid by electrochemical advanced oxidation processes using a boron-doped diamond anode and Fe²⁺ and UVA light as catalysts. Appl. Catal. B Environ., 72, 373.

39.

Sun

, and Pignatello

J.J.

(1993). Photochemical reactions involved in the total mineralization of 2,4-D by Fe³⁺/H₂O₂/UV. Environ. Sci. Technol., 27, 304.

40.

Tang

W.Z.

, and Huang

C.P.

(1997). Stoichiometry of Fenton's reagent in the oxidation of chlorinated aliphatic organic pollutants. Environ. Technol., 18, 13.

41.

Tian

, Adams

, Wen

, Matthew Asmussen

, and Chen

(2009). Photoelectrochemical oxidation of salicylic acid and salicylaldehyde on titanium dioxide nanotube arrays. Electrochim. Acta, 54, 3799.

42.

Wang

, Qu

, Liu

, and Ru

(2008). Mineralization of an azo dye Acid Red 14 by photoelectro-Fenton process using an activated carbon fiber cathode. Appl. Catal. B Environ., 84, 393.

43.

Wang

C.T.

, Chou

W.L.

, Chung

M.H.

, and Kuo

Y.M.

(2010). COD removal from real dyeing wastewater by electro-Fenton technology using an activated carbon fiber cathode. Desalination, 253, 129.

44.

Wang

, Deng

, Jin

, and Huang

(2012). Gallic acid modified hyper-cross-linked resin and its adsorption equilibria and kinetics toward salicylic acid from aqueous solution. Chem. Eng. J., 191, 195.

45.

Wells

C.F.

, and Salam

M.A.

(1965). Hydrolysis of ferrous ions: A kinetic method for determination of the Fe(II) species. Nature, 205, 690.

46.

Wells

C.F.

, and Salam

M.A.

(1968). The effect of pH on the kinetics of the reaction of iron(II) with hydrogen peroxide in perchlorate media. J. Chem. Soc. A, 1968, 24.

47.

Zhang

, Yang

, Gao

, Fang

, and Liu

(2008). Electro-Fenton degradation of azo dye using polypyrrole/anthraquinone disulphonate composite film modified graphite cathode in acidic aqueous solutions. Electrochim. Acta, 53, 5155.

48.

Zhang

, Yang

, and Liu

(2009). Comparative study of Fe^2+/H₂O₂ and Fe³⁺/H₂O₂ electro-oxidation systems in the degradation of amaranth using anthraquinone/polypyrrole composite film modified graphite cathode. J. Electroanal. Chem., 632, 154.

49.

Zhang

, Zhang

, and Zhou

(2006). Removal of COD from landfill leachate by electro-Fenton method. J. Hazard. Mater. B135, 106.

50.

Zhou

, Yu

, Lei

, and Barton

(2007). Electro Fenton method for the removal of methyl red in an efficient electrochemical system. Sep. Purif. Technol., 57, 380.

Electro-Fenton Method Oxidation of Salicylic Acid in Aqueous Solution with Graphite Electrodes

Abstract

Abstract

Introduction

Materials and Methods

Materials

Experimental procedure

Results and Discussion

Effect of Fe2+ concentration

Effect of pH

Effect of voltage

Effect of flow rate

Effect of electrode spacing

Conclusions

Footnotes

Author Disclosure Statement

References

Effect of Fe²⁺ concentration