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DPV-assisted understanding of TiO2 photocatalytic decomposition of aspirin by identifying the role of produced reactive species
Applied Catalysis B: Environmental 266 (2020) 118646
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DPV-assisted understanding of TiO2 photocatalytic decomposition of aspirin by identifying the role of produced reactive species
T
Abbas Valia,b, Hesam Zamankhan Malayeria, Mohammadmehdi Azizib, Hyeok Choia,* a b
Department of Civil Engineering, The University of Texas at Arlington, 416 Yates Street, Arlington, TX 76019-0308, USA Department of Chemistry and Biochemistry, The University of Texas at Arlington, 700 Planetarium Place, Arlington, TX 76019-0065, USA
ARTICLE INFO
ABSTRACT
Keywords: TiO2 photocatalysis Acetylsalicylic acid Reactive species Reaction pathway Differential pulse voltammetry
A TiO2/UV photocatalytic process generates various reactive species via different oxidation and reduction routes and thus the contribution of such reactive species to the decomposition of organic chemicals in water has been unclear. In this study, decomposition of acetylsalicylic acid (ASA) as a probe was performed to find the specific role of charge carriers (ecb− and hvb+) and primary active species (such as %O2−, %OH, %HO2, and H2O2). Differential pulse voltammetry (DPV) was introduced as an in-situ fast electrochemical method to identify produced intermediates. The study revealed that hydroxyl radicals (%OH) produced by reduction of oxygen with ecb− and subsequent primary reactions (reduction pathway) play the major role in the photocatalytic decomposition of ASA, in comparison to direct oxidation of OH− or ASA by hvb+ (oxidation pathway). The in-situ DPV analysis probed that sequential %OH addition to the ASA aromatic ring is the favorable oxidation mechanism.
1. Introduction Heterogeneous photocatalysis using semiconducting materials has been of great interest [1]. In particular, the photocatalytic decomposition mechanisms of many organic contaminants in water have been most extensively studied with titanium dioxide (TiO2) [1–5]. The process starts with generation of charge carriers at TiO2 under ultraviolet (UV) radiation, i.e., electrons in conduction band (ecb−) and holes in valence band (hvb+) [1]. Some of the charge carriers migrate to the surface of TiO2 and thus are available for chemical reactions in water. In fact, the charge carriers themselves are a strong reductant or oxidant so that organic chemicals can be directly reduced by ecb- (e.g., benzoquinone) and oxidized by hvb+ (e.g., 2-propanol) [2,6]. The photogenerated ecb− is also trapped in vacancies nearby the surface to reduce electron acceptors such as oxygen (O2) abundant in water to superoxide radical anion (%O2−) while the photogenerated hvb+ takes electrons from electron donors such as hydroxide ion (OH−) abundant in water to produce mostly hydroxyl radicals (%OH) [1]. Consequently, the reaction between the charge carriers and electron accepters and donors generates primary reactive species including %O2−, %OH, %HO2, and H2O2 (Table S1), which are, then as either oxidants or reductants, also involved in decomposition of contaminants in water, so-called secondary reactions (Table S2). Organic chemicals are decomposed via reductive pathways by ecb−
⁎
and/or %O2− while they are also decomposed via oxidative pathways by hvb+ and/or %OH. In particular, %OH, which is a strong oxidizing species, is believed to significantly contribute to decomposition of many organic chemicals [7–9]. As summarized in Fig. 1 (also Table S1), •OH can be formed through the direct oxidation of OH− by hvb+ (i.e., oxidation pathway R2) and/or the reduction of oxygen by ecb− followed by subsequent reactions (i.e., reduction pathway R3-R7) [1,7,10–13]. Previous studies showed different contributions of the oxidation pathway and the reduction pathway to the photocatalytic decomposition of aromatic chemicals (Table S3) [9,14–18]. Since the reduction pathway to produce %OH occurs in the presence of O2, those studies compared reaction kinetics in the presence (i.e., both reduction and oxidation pathways) and the absence of O2 (i.e., oxidation pathway). Salicylic acid (SA) showed a great decrease in its first order reaction rate constant from 120 × 10-3 min-1 with O2 to 2.5 × 10-3 min-1 without O2 (98 % decrease) while 2,3-dichlorophenol showed a slight decrease in its first order reaction rate constant from 1.2 × 10-3 min-1 with O2 to 1.0 × 10-3 min-1 without O2 (only 17 % decrease) [15,18]. The results indicate that the reduction pathway to generate •OH is significant for SA but not for 2,3-dichlorophenol, implying that the contribution of each of the pathways to the decomposition of organic chemicals seems to be case-specific and other reactive species than •OH might also be involved in the decomposition reaction. It is not clear which reaction pathway is more favorable for the generation of %OH
Corresponding author. E-mail address: [email protected] (H. Choi).
https://doi.org/10.1016/j.apcatb.2020.118646 Received 15 October 2019; Received in revised form 8 January 2020; Accepted 16 January 2020 Available online 16 January 2020 0926-3373/ © 2020 Elsevier B.V. All rights reserved.
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Fig. 1. Generation of reactive species from charge carriers (R1) via the primary reduction and oxidation pathways (R2-R7) on TiO2 under UV radiation. Through either the reduction pathway or the oxidation pathway, strong hydroxyl radicals are formed.
and thus for the decomposition of a specific organic chemical. In addition, as summarized in Table S2, it is less known how each of the reactive species including ecb−, hvb+, %O2−, %OH, %HO2, and H2O2 contributes to the decomposition of organic chemicals. As a result, the objective of this study is to clarify the role of the primary reactive species in the photocatalytic decomposition of an organic chemical, through systematic experimental design to quench certain reactions producing specific reactive species. Acetylsalicylic acid (ASA, known as aspirin) was carefully chosen as a probe chemical because ASA is not electroactive with hvb+ and ecb− and thus the direct role of the charge carriers in the decomposition of ASA can be largely ignored, making interpretation of ASA decomposition results less complicated [19,20]. Along with mass spectrometry (MS), differential pulse voltammetry (DPV) was introduced as an in-situ fast electrochemical method to identify produced reaction intermediates and thus propose photocatalytic decomposition mechanisms of ASA.
its hydrolysis is negligible [22]. 2.3. Electrochemical analysis In-situ DPV analysis was deployed to probe the photocatalytic decomposition intermediates of ASA. A single compartment, three-electrode cell setup, and a potentiostat (Model CHI720C) were used for the measurement. Glassy carbon electrode with 0.07 cm2 surface area was used as a working electrode. Platinum wire and Ag/AgCl/4 M KCl were used as a counter electrode and a reference electrode, respectively. Operation parameters of DPV were: initial potential = −0.4 V, final potential = 1.0 V, incremental potential = 0.004 V, amplitude = 0.05 V, pulse width = 0.05 s, sample width = 0.0167 s, and pulse period = 0.2 s. For electrochemical measurement of remaining concentration of ASA which is not electroactive during photocatalytic decomposition, an indirect method was used where ASA was hydrolyzed quickly at high pH (> 13) to SA (see reaction R18) which is electroactive [19,20]. At each time interval, 0.4 mL of solution was withdrawn and mixed with 0.4 mL of 0.4 M NaOH to indirectly measure concentration of ASA by DPV (Fig S1). This measured concentration by DPV is total concentration of ASA and SA. At the same time, concentration of SA was measured at pH 3.5 by in-situ DPV method. Remaining concentration of ASA was calculated by subtracting the concentration of SA from the total concentration. Similarly, for electrochemical measurement of remaining concentration of ASA which is not electroactive during hydrolysis at pH 3.5, in-situ DPV method was also deployed. According to reaction R18, moles of hydrolyzed ASA are equal to moles of produced SA during hydrolysis reaction. Accordingly, concentration of SA during hydrolysis of ASA was first measured by in-situ DPV method at pH 3.5 (Fig. S2). Then, remaining concentration of ASA (i.e., unhydrolyzed ASA) was calculated by subtracting the measured concentration of SA from the initial concentration of ASA. The same electrochemical set up was deployed for cyclic voltammetry analysis, where initial potential = −0.3 V, high potential = 1.2 V, final potential = -0.3 V, scan rate =5.0 mV/s, and rest time = 2.0 s.
2. Experimental methods 2.1. Chemicals and reagents ASA, SA, paracetamol (PA), 2,3-dihydroxy benzoic acid (2,3DHBA), 2,5-dihydroxy benzoic acid (2,5-DHBA), monosodium phosphate, disodium phosphate, 2-propanol, fluorine doped tin oxide (FTO), ethanol (all from Sigma Aldrich), sodium hydroxide, phosphoric acid, ammonium oxalate monohydrate (all from Alfa Aesar), and TiO2 (P25; Degussa) were used in this study, as received. Double distilled water was used for all the experiments. 2.2. Photocatalytic decomposition TiO2 (P25, Degussa) nanoparticles with 25 ± 5 nm particle size, specific surface area of 50 m2/g, and mixture of 80 % anatase and 20 % rutile was used. Concentrations of ASA and TiO2 were 1.0 mM and 1.0 g/L, respectively. Photocatalytic reactor volume was 25 mL and effective volume was 15 mL. Solution was mixed by using a magnetic stirrer. Temperature was maintained at 25 ± 1.0 °C. Solution was initially kept in dark condition for 30 min to achieve adsorption equilibrium, and then irradiated with UV for 420 min. Two 15 W low pressure mercury UV lamps emitting 365 nm were used. Lamps were mounted at 10 cm above the reactor, giving light intensity of 0.4 mW/ cm2 measured by a photonics power meter (Ophir). For experiments requiring oxygen, air was purged into the reactor while for those requiring no oxygen, nitrogen gas was purged. Purging rate was 40 mL/ min and purging was initially applied for 10 min under dark condition and then continued over UV irradiation. Buffer solution was used to prevent pH change and thus to make it easy to interpret ASA decomposition mechanism. In particular, 0.1 M phosphate buffer was used as a supporting electrolyte for in-situ electrochemical analysis of decomposition intermediates. Monosodium phosphate, disodium phosphate, and phosphoric acid with specific molar ratio were mixed to prepare 0.1 M phosphate buffer with pH of 3.5, 5.7, and 7.0 [21]. In particular, pH 3.5 was selected as a standard condition for most of experiments because direct oxidation of OH− by hvb+ to produce •OH can be significantly minimized under the acidic condition and ASA has the highest chemical stability at pH 3.5 so that
2.4. Photoelectrochemistry setting The same potentiostat (Model CHI720C) used for the measurement of the decomposition intermediates of ASA was employed. For preparation of TiO2 photoelectrode, 20 mg of TiO2 P25 was dispersed in 25 mL of ethanol by sonication for 1 min, and then FTO (sheet resistance: ∼7 Ω/sq) substrate with 2.5 cm2 surface area was placed at the bottom of TiO2 suspension for 5 min. Deposited TiO2 on FTO was washed with water and dried for 30 min at room temperature. The prepared TiO2 photoelectrode on FTO substrate was served as a working electrode. Platinum foil and Ag/AgCl/4 M KCl were used as a counter electrode and a reference electrode, respectively. The electrochemical cell was placed at 20 cm away from a xenon arc lamp. A 360 nm UV cut-off glass was placed in front of the electrochemical cell to prevent photoactivation of FTO. The potential was swept from -0.3 V to +1.0 V with low scan rate of 5.0 mV/s. Light was manually chopped in 5 s interval. Light intensity was 200 mW/cm2, which was measured using a radiant power meter (Newport 70260). 2
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2.5. Mass spectrometry
ASA, SA, and PA under different oxygen and media conditions labeled with E1-E9. In comparison between E1 and E2, it was found that there is no photocatalytic decomposition of ASA in phosphate buffer at pH 3.5 in the absence of oxygen (Table 1). It is obvious that the primary engine reaction R3 and the subsequent primary reactions R4-R7 are quenched by purging the reaction solution with nitrogen. Even the oxidation pathway was negligible, i.e., hvb+ did not directly oxidize OH− (R2) and ASA (R11). The reduction pathway in the presence of O2 was significant where ecb− plays the main role in the photocatalytic decomposition of ASA by reducing oxygen (R3) and driving subsequent reactions (R4-R7) to produce %OH. The same experiment was conducted with SA and PA (E3-E6), given that ASA is not electroactive while SA with hydroxyl group and PA with hydroxyl and amide groups might be subject to direct electrochemical oxidation. However, the results with SA and PA were very similar to those with ASA, as also reported elsewhere for SA [15]. The results indicated that ASA, SA, and PA do not have any functional groups vulnerable to direct oxidation by hvb+. In fact, several chemicals including oxalate, formate, methanol, ethanol, and 2-propanol are directly oxidized by hvb+, known as hole scavengers [23–25]. For example, methanol molecules in water are adsorbed well onto TiO2 surface and then attacked by hvb+. The main parameter for the direct oxidation of organic chemicals by hvb+ is charge-transfer rate between photocatalyst surface and organic chemicals, which is function of oxidation standard potential and adsorption behavior [1]. Although the absence of O2 can prevent oxygen-dependent primary reactions (R3-R7) and related secondary reactions (R13-R17), there is one possible secondary reaction to decompose ASA, which is direct reduction of ASA by ecb− (R12). However, the observed negligible decomposition of ASA suggests that even R12 did not occur, which is well expected, given that there is no reductable functional group in ASA to participate in a direct electron transfer reaction.
In order to further identify reaction intermediates of ASA, samples were also introduced to ion trap-time off flight MS (Shimadzu, IT-TOF), where negative electrospray ionization mode was used. Size of tube was 38" × 1/16" OD × 0.005" ID. Water with flow rate of 0.1 mL/min was used as eluent and injection volume was 2.0 μL. Other parameters of data acquisition included: range of m/z acquisition = 100 − 220 Da, acquisition time =5 min, ion accumulation time =10 ms, purge time =1 min with water, and auto-sampler rinsing volume =100 μL. 3. Results and discussion Photocatalytic decomposition of ASA occurs via two steps: (i) generation of charge carriers (ecb− and hvb+) and reactive species (•O2−, • OH, •HO2, and H2O2), so-called primary reaction (R1-R10 in Table S1) and (ii) interaction of the primary species with ASA and its reaction intermediates, so-called secondary reaction (R11-R18 in Table S2). The following sections describe modification of photocatalytic reaction conditions to reveal the contribution of the produced charge carriers and reactive species as oxidants or reductants to decomposition of ASA. 3.1. Hydrolysis of ASA To quantify the direct hydrolysis of ASA to SA and acetic acid (AcAc) (R18), the hydrolysis rate of ASA was measured at pH 3.5 in phosphate buffer (Fig. S3). The first order rate constant of its hydrolysis was 8.66 × 10−5 min-1, which is negligible, as also reported elsewhere [22]. The photocatalytic decomposition of ASA is shown in Fig. 2. Result on dark adsorption without UV indicated no significant adsorption of ASA onto TiO2. Direct photolysis of ASA without TiO2 was also negligible. ASA was significantly removed only via photocatalytic decomposition mechanism and first order reaction rate constant was at 1.77 × 10-2 min-1, which is much higher than its hydrolysis rate constant of 8.66 × 10−5 min-1. The absence of the direct hydrolysis of ASA can make it less complicated to interpret the decomposition mechanism of ASA.
3.3. Removal of buffer and oxygen One possible reason that hvb+ did not directly oxidize OH− (R2) and ASA (R11) is unfavorable adsorption of OH− and/or ASA onto the surface of TiO2. Both phosphate buffer and oxygen were removed to find the effect of phosphate buffer on the adsorption process (E7). Interestingly, ASA was decomposed in this case. However, the decomposition kinetics in E7 was two orders of magnitude slower than that in E1. Nonetheless, removing phosphate buffer caused favorable adsorption of OH− and/or ASA which can be oxidized by hvb+. In general, the results from E1-E7 indicated that the oxidation pathway to produce •OH is significantly slower than the reduction pathway. The standard potentials of associated reactions can explain the finding. Oxidation of OH− to produce %OH (Eq. 1) competes with oxygen evolution reaction (Eq. 2). Since the standard potential of OH−/%OH is higher than OH−/O2 + H2O in any pHs (note Fig. S4), oxygen evolution reaction is thermodynamically more favorable than OH−/%OH oxidation reaction. Meanwhile, oxygen reduction (Eq. 3) competes with hydrogen reduction reaction (Eq. 4). Since the standard potential of oxygen reduction at pHs above 2.71 is less than hydrogen reduction reaction (Fig. S4), oxygen reduction is more favorable than hydrogen reduction reaction.
3.2. Removal of oxygen Both the oxidation (R2) and reduction (R3-R7) pathways produce OH in presence of oxygen while only the oxidation pathway (R2) produces •OH in absence of oxygen (Fig. 1 and Table S1). Table 1 summarizes the photocatalytic decomposition reaction rate constants of
•
OH− + hvb+ ⇆ •OH or OH− ⇆ •OH + e−; E0 = +1.695 eV vs. NHE (1) 4OH− + 4hvb+ ⇆ O2 + 2H2O or 4OH− ⇆ O2 + 2H2O + 4e- ; E0 = +0.401 eV vs. NHE (2) O2 + ecb− ⇆ •O2−; E0 = −0.160 V vs. NHE 2H Fig. 2. UV photolytic (-◼-) and UV/TiO2 photocatalytic decomposition of 1.0 mM ASA (..►..) at pH 3.5.
+
+
2ecb−
0
⇆ H2; E = 0.0 V vs. NHE
(3) (4)
In addition, the conduction band electron trapping by oxygen is much faster than valence band hole trapping by OH− because the 3
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Table 1 Photocatalytic decomposition of ASA, SA, and PA under different oxygen and media conditions. Experiment No.
Chemical
O2 (mg/L)
Media
Reaction rate constant (min−1)a
E1 E2 E3 E4 E5 E6 E7 E8
ASA ASA SA SA PA PA ASA ASA
8.4 0 8.4 0 8.4 0 0 8.4
1.77 NDb 1.81 ND 2.31 ND 1.86 4.81
E9
ASA
8.4
0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water Water 0.1 M phosphate buffer + 0.1 M ammonium oxalate in water 0.1 M phosphate buffer + 0.1 M 2-propanol in water
a b
± 0.01 × 10−2 ± 0.01 × 10−2 ± 0.02 × 10−2 ± 0.01 × 10−4 ± 0.02 × 10−3
ND
Note that pH 3.5 was used for all experiments. Not determined: Reaction within the time frame was negligible to determine the rate constant.
conduction band electron has lower effective mass than valence band hole, as verified elsewhere [26–29]. However, the reasons mentioned above might be more reasonable when organic chemicals are less favorable for adsorption onto the surface of TiO2 than OH−. If organic chemicals adsorb well onto TiO2 and have low oxidation potential, the direct oxidation of organic chemicals by hvb+ should be considered.
3.5. Addition of %OH and hole scavenger The reactions (R3-R7) produce the primary active specious (%O2−, OH, %HO2, and H2O2), which directly react with ASA (R13-R17). To find the contribution of the primary species to ASA decomposition, 2propanol was added to phosphate buffer at pH 3.5 (E9) because 2propnaol is a strong %OH scavenger to prevent R16 and R17 [9,31]. In addition, 2-propanol is a hvb+ scavenger to inhibit R2 and R11 [23–25]. 2-proanol is directly oxidized to acetone (Eq. 6). As shown in Fig. 3, the presence of 2-propanol produced almost 1.5 times higher photocurrent than ammonium oxalate, indicating that hvb+ reacts with 2-propanol faster than ammonium oxalate. No decomposition of ASA was observed in the presence of 2-propanol (Table 1). It seems that 2-propanol successfully acted as a hole and •OH scavenger to prevent ASA decomposition.
%
3.4. Addition of hole scavenger Previously, in order to find contribution of the oxidation pathway, the reduction pathway to produce %OH was killed with no O2, concluding that %OH production by the reduction pathway in the presence of O2 is significant. Inversely, the oxidation pathway (R2) was quenched (E8) by adding 0.1 M ammonium oxalate as a hole scavenger to confirm the results of E1-E7 [30]. As shown in Fig. 3, the significant increase in photocurrent by adding ammonium oxalate into phosphate buffer at pH 3.5 confirmed that hvb+ can immediately oxidize oxalate (Eq.5). Consequently, the observed photocatalytic decomposition of ASA (reaction rate constant at 4.81 × 10−3 min-1) even in the presence of the hole scavenger indicated that ecb- drives the photocatalytic decomposition of ASA.
C3H6O + 2H+ + 2e− ⇆ C3H8O or C3H6O + 2H+ ⇆ C3H8O + 2hvb+; E0 = +0.13 V vs. NHE (6) 3.6. DPV-assisted understanding of ASA decomposition mechanism Based on the results from E1-E9, it can be concluded that •OH exclusively produced by reduction of oxygen by ecb− (R3) and subsequent primary reactions (R4-R7), so-called reduction pathway, was responsible for the photocatalytic decomposition of ASA. To identify reaction intermediates and thus to propose ASA decomposition mechanisms by •OH, in-situ electrochemical analysis DPV was performed along with MS analysis. Fig. 4 shows DPV analysis of intermediates formed during the photocatalytic decomposition of ASA and SA for 90 min, exhibiting clear five anodic peaks at 0.166, 0.342, 0.442, 0.842, and 0.982 V. In addition, MS analysis of the same samples showed m/z values of 137, 153, 167, 169, 179, and 195 (Table S4). The peak at 0.982 V was attributed to the presence of SA because the reference DPV voltamomogram of SA confirmed the peak (Fig. S5). The m/z at 137 also confirmed the presence of SA as one of reaction intermediates from ASA. As mentioned early, direct hydrolysis of ASA to SA was negligible at pH 3.5 and thus the formation of SA can be explained by %OH attack to ASA, in particular the single bond between oxygen and carbon (C10O11, note Fig. 5a) in ASA with the lowest bond dissociation energy at 78.0 kcal/mol [32]. All DPV peaks for ASA and SA were the same except for the peak at +0.842 V only in ASA while all MS m/z peaks for ASA and SA were the same except for the peaks at 179 and 195 only in ASA. The m/z at 179 indicated ASA and thus m/z at 195 suggested the presence of another noticeable intermediate from ASA. •OH addition to aromatic groups is an electrophilic substitution (note carbons 3, 4, 5, and 6 in Fig. 5a). Since the ester functional group in ASA is an activator (i.e., donating electron to the ring), it tends to drive incoming %OH to the ortho or para
2CO2 + 2e− ⇆ C2O42- or 2CO2 ⇆ C2O42- + 2hvb+; E0 = +0.51 V vs. NHE (5)
Fig. 3. Photocurrent vs. potential at pH = 3.5 for (i) 0.1 M phosphate buffer, (ii) 0.1 M phosphate buffer and 0.1 M ammonium oxalate, and (iii) 0.1 M phosphate buffer and 0.1 M 2-propanol. 4
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DPV peaks (e.g., 0.166, 0.342, and 0.442 V in Fig. 4) and m/z peaks (e.g., 153, 167, and 169 in Table S4) in ASA and SA, •OH attack mechanisms to SA and HASA were proposed. Since hydroxyl group is an activator and carboxylic group is deactivator and thus both functional groups in SA direct •OH to carbon 3 or 5 (Fig. 5c), formation of 2,5DHBA and 2,3-DHBA is highly possible (Fig. 5e and f) [33]. A comparison between the in-situ voltammogram in Fig. 4 and reference voltammograms of 2,5-DHBA and 2,3-DHBA (Fig. S5) confirmed the presence of 2,5-DHBA and 2,3-DHBA formed from SA. The m/z peak at 153 also verified the presence of 2,5-DHBA and 2,3-DHBA. Meanwhile, formation of 2,5-DHBA and 2,3-DHBA can also be originated from 5HASA and 3-HASA, respectively because the single bond C10-O11 in 3HASA (Fig. 5d) and 5-HASA (Fig. 5b) is vulnerable to %OH attack [32]. To explain the m/z peak 169 (167 later), trihydroxybenzoic acid (THBA; MW of 170, Fig. 5g) formed by •OH addition to 2,5-DHBA and 2,3-DHBA was proposed. Cyclic voltammograms of 2,3-DHBA and 2,5DHBA at pH 3.5 are shown in Fig. 7. In the first cycle, one oxidation peak, A1 for 2,5-DHBA and A2 for 2,3-DHBA, was revealed. The counterpart reduction peaks for A1 and A2 are C1 and C2, respectively, indicating that 2,5-DHBA and 2,3-DHBA can be oxidized to counterpart oxidized forms, as shown in Fig. 8, i.e., 8f (2,5-DHBA) to 8 g and 8a (2,3-DHBA) to 8b in oxidation vice versa in reduction. However, oxidation of 2,5-DHBA and 2,3-DHBA is not reversible at pH 3.5 because there is an additional cathodic peak, C3 for 2,5-DHBA and C4 for 2,3DHBA. The peaks C3 and C4 disappeared gradually by increasing pHs from 3.5–7.0 (note Fig. S6), as also observed elsewhere [34]. Disappearing C3 and C4 peaks at high pHs imply that 8b and 8 g promptly react with H+ and H2O under such a low pH of 3.5 to produce chemicals 8d and 8i. Meanwhile, the oxidized counterparts of C3 and C4 are A3 and A4 in the second cycle. The peak potentials of C3/C4 couple and A3/A4 couple were the same, suggesting that possibly oxidation product of 2,5-DHBA and 2,3-DHBA is the same as 2,3,5-THBA (8e). The m/z peaks in MS analysis corresponded to the chemicals 8d and 8i (m/z 167) and 2,3,5-THBA (m/z 169). DPV analysis to prove time evolution of primary intermediates (HASA and SA) and secondary intermediates (2,3-DHBA, 2,5-DHBA, and 2,3,5-THBA) is shown in Fig. 9 while concentration-time evolution of primary and secondary reaction intermediates formed by the photocatalytic decomposition of ASA is summarized in Fig. 10. The primary intermediates, SA and HASA, reached the highest concentrations after around 60 min and they were completely decomposed after 360 min
Fig. 4. In-situ electrochemical DPV analysis of intermediates from the photocatalytic decomposition of 1.0 mM ASA and 1.0 mM SA at pH = 3.5 for 90 min.
position (carbon 3 or 5 in Fig. 5a). However, the carboxylic acid moiety as an electron withdrawing group directs incoming %OH to the meta position (carbon 3 or 5 in Fig. 5a). As a result, both functional groups (i.e., carboxyl and ester) direct incoming %OH to the carbon 3 or 5 in ASA molecule to form 3-hydroxyacetylsalicylic acid (3-HASA shown in Fig. 5d) or 5-hydroxyacetylsalicylic acid (5-HASA shown in Fig. 5b). The m/z peak at 195 indicated the presence of such HASAs. The DPV peak at +0.842 V could be ascribed to the HASAs. Fig. 6 shows DPV analysis of primary intermediates (SA and HASA) formed by the photocatalytic decomposition of ASA at pH = 3.5, 5.7, and 7.0. SA and HASA (3-HASA or 5-HASA) peaks are shifted to low potentials over pH increase with a slope of -51.1 mV/pH for SA and -56.5 mV/pH for HASA. The relationship between peak potential and pH is shown in Eq. 7, where m and n are the numbers of protons and electrons, respectively. When the numbers of protons and electrons are identical, the theoretical slope is 59 mV/pH, which is very close to the results at 51.1 and 56.5. As a result, the number of involved protons was almost the same as that of electrons for oxidation of SA and HASA. Ep = E p
(pH=0)
– (0.059 m/n) pH
(7)
To identify more reaction intermediates corresponding to common
Fig. 5. Proposed photocatalytic decomposition pathways of ASA at pH 3.5 by hydroxyl radical attack mechanisms. 5
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Fig. 6. (a) In-situ electrochemical DPV analysis of primary intermediates (SA and HASA) formed by the photocatalytic decomposition of 1.0 mM ASA in 60 min at pH = 3.5, 5.7, and 7.0 and b) DPV peak potential shift of SA and HASA over pHs.
4. Conclusions This study provided a systematic experimental design to quench certain reactions producing specific reactive species and thus to find the role of such reactive species in TiO2 photocatalytic decomposition of organic chemicals in water. ASA was selected as a probe chemical. The results revealed that %OH production by reduction of oxygen with ecb− and subsequent primary reactions (reduction pathway) played the major role in the photocatalytic decomposition of ASA as a probe chemical, in comparison to direct oxidation of OH− or ASA by hvb+ (oxidation pathway). In-situ electrochemical DPV analysis greatly helped to understand the decomposition mechanisms and pathways of ASA and thus to probe that sequential •OH addition to the ASA aromatic ring is the favorable oxidation mechanism, forming SA, 3-HASA, 5HASA, 2,3-DHBA, 2,5-DHBA, and 2,3,5-THBA. Findings from this study would help to set a strategy to control the TiO2/UV operating parameters for generating effective reactive species, leading to favorable reaction pathways and fast decomposition of organic contaminants in water.
Fig. 7. Cyclic voltammogram of 0.1 mM 2,3-DHBA (1 st and 2nd cycles) and 0.1 mM 2,5-DHBA (1 st and 2nd cycles) at 25 mV/s.
CRediT authorship contribution statement
while the secondary intermediates reached the highest concentrations after around 180 min for 2,3-DHBA and 2,5-DHBA and 240 min for 2,3,5-THBA and they were completely decomposed after 420 min. Although we have not investigated in detail beyond the point, further oxidation of 2,3,5-THBA may cause aromatic ring cleavage to yield carboxylic acids such as malic acid and succinic acid.
Abbas Vali: Conceptualization, Data curation, Investigation, Writing - original draft. Hesam Zamankhan Malayeri: Methodology. Mohammadmehdi Azizi: Methodology. Hyeok Choi: Supervision, Writing - review & editing.
Fig. 8. Possible oxidation mechanism of 2,3-DHBA and 2,5-DHBA at pH 3.5. A1-A4 and C1-C4 are oxidation and reduction CV peaks, respectively, which are shown in Fig. 7.
6
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Fig. 9. In-situ electrochemical DPV analysis on time evolution of (a) primary intermediates (HASA and SA) and (b) secondary intermediates (2,3-DHBA, 2,5-DHBA, and 2,3,5THBA) formed by the photocatalytic decomposition of 1.0 mM ASA at pH = 3.5 for (a1) 0−60 min and (a2) 60−360 min and (b1) 0−180 min and (b2) 240−420 min. Please note two different time slots were purposely used to show concentration increase by up to around 60 min for (a) and 180 min for (b) and then concentration decrease.
Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.apcatb.2020.118646. References [1] J. Schneider, M. Matsuoka, M. Takeuchi, J. Zhang, Y. Horiuchi, M. Anpo, D.W. Bahnemannt, Understanding TiO2 photocatalysis: mechanisms and materials, Chem. Rev. 114 (2014) 9919–9986, https://doi.org/10.1021/cr5001892. [2] M.R. Hoffmann, S.T. Martin, W. Choi, D.W. Bahnemannt, Environmental applications of semiconductor photocatalysis, Chem. Rev. 95 (1995) 69–96, https://doi. org/10.1021/cr00033a004. [3] K. Hashimoto, H. Irie, A. Fujishima, TiO2 photocatalysis: a historical overview and future prospects, Japan J. Appl. Phys. 44 (2005) 8269–8285, https://doi.org/10. 1143/JJAP.44.8269. [4] A. Mills, S.L. Hunte, An overview of semiconductor photocatalysis, J. Photochem. Photobiol. A 108 (1997) 1–35, https://doi.org/10.1016/S1010-6030(97)00118-4. [5] Y. Nosaka, A.Y. Nosaka, Generation and detection of reactive oxygen species in photocatalysis, Chem. Rev. 117 (2017) 11302–11336, https://doi.org/10.1021/acs. chemrev.7b00161. [6] M.A. Henderson, M. Shen, Electron-scavenging chemistry of benzoquinone on TiO2 (110), Top. Catal. 60 (2017) 440–445, https://doi.org/10.1007/s11244-0160707-7. [7] Y. Nosaka, A. Nosaka, Understanding hydroxyl radical (%OH) generation processes in photocatalysis, ACS Energy Lett. 1 (2016) 356–359, https://doi.org/10.1021/ acsenergylett.6b00174. [8] K.I. Ishibashia, A. Fujishima, T. Watanabea, K. Hashimoto, Quantum yields of active oxidative species formed on TiO2 photocatalyst, J. Photochem. Photobiol. A 134 (2000) 139–142, https://doi.org/10.1016/S1010-6030(00)00264-1. [9] L. Yang, L.E. Yu, M.B. Ray, Photocatalytic oxidation of paracetamol: dominant reactants, intermediates, and reaction mechanisms, Environ. Sci. Technol. 43 (2009) 460–465, https://doi.org/10.1021/es8020099. [10] P.M. Wood, The two redox potentials for oxygen reduction to superoxide, Trends Biochem. Sci. 12 (1987) 250–251, https://doi.org/10.1016/0968-0004(87) 90123-X. [11] P.S. Rao, E. Hayon, Redox potentials of free radicals. IV. superoxide and hydroperoxy radicals•O2− and •HO2, J. Phys. Chem. 79 (1975) 397–402, https://doi.org/ 10.1021/j100571a021. [12] D.R. Lide, CRC Handbook, 84th ed., CRC Press, Boca Raton, Florida, 2003, pp. 5:1–5:2 5:10. [13] J.A. Dean, Lange’s Handbook of Chemistry, 11th ed., McGraw-Hill, New York, 1979 9:4-9:128.
Fig. 10. Concentration-time evolution of primary and secondary reaction intermediates formed by the photocatalytic decomposition of 1.0 mM ASA at pH = 3.5.
Declaration of Competing Interest The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper. Acknowledgements This research was supported by the University of Texas at Arlington through new faculty startup funds. The authors thank Dr. Krishnan Rajeshwar for allowing us to use his electrochemistry laboratory at the University of Texas at Arlington.
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A. Vali, et al. [14] L. Yang, L.E. Yu, M.B. Ray, Degradation of paracetamol in aqueous solutions by TiO2 photocatalysis, Water Res. 42 (2008) 3480–3488, https://doi.org/10.1016/j. watres.2008.04.023. [15] C. Karunakaran, B. Naufal, P. Gomathisankar, Efficient photocatalytic degradation of salicylic acid by bactericidal ZnO, J. Korean Chem. Soc. 56 (2012) 108–114, https://doi.org/10.5012/jkcs.2012.56.1.108. [16] S. Zheng, Y. Cai, K.E. O’Shea, TiO2 photocatalytic degradation of phenylarsonic acid, J. Photochem. Photobiol. A 210 (2010) 61–68, https://doi.org/10.1016/j. jphotochem.2009.12.004. [17] D.-H. Tseng, L.-C. Juang, H.-H. Huang, Effect of oxygen and hydrogen peroxide on the photocatalytic degradation of monochlorobenzene in TiO2 aqueous suspension, Int J Photoenergy 1 (2012) 1–9, https://doi.org/10.1155/2012/328526. [18] H.C. Liang, X.Z. Li, Y.H. Yang, K.H. Sze, Effects of dissolved oxygen, pH, and anions on the 2,3-dichlorophenol degradation by photocatalytic reaction with anodic TiO2 nanotube films, Chemosphere 73 (2008) 805–812, https://doi.org/10.1016/j. chemosphere.2008.06.007. [19] E.R. Sartori, R.A. Medeiros, R.C. Rocha-Filho, O. Fatibello-Filho, Square-wave voltammetric determination of acetylsalicylic acid in pharmaceutical formulations using a boron-doped diamond electrode without the need of previous alkaline hydrolysis step, J. Braz. Chem. Soc. 20 (2009) 360–366, https://doi.org/10.1590/ S0103-50532009000200022. [20] C. Cofan, C. Radovan, Anodic determination of acetylsalicylic acid at a mildly oxidized boron-doped diamond electrode in sodium sulphate medium, Int. J. Electrochem. Sci. 1 (2011) 1–9, https://doi.org/10.4061/2011/451830. [21] Council of Europe, European Pharmacopoeia Commission, The European Pharmacopoeia, 6th ed., Conseil de l’Europe, Strasburg, France, 2007, pp. 508–514. [22] C.A. Kelly, Determination of the decomposition of aspirin, J. Pharm. Sci. 59 (1979) 1053–1079, https://doi.org/10.1002/jps.2600590802. [23] Y. Tamaki, A. Furube, M. Murai, K. Hara, R. Katoh, M. Tachiya, Direct observation of reactive trapped holes in TiO2 undergoing photocatalytic oxidation of adsorbed alcohols: evaluation of the reaction rates and yields, J. Am. Chem. Soc. 128 (2006) 416–417, https://doi.org/10.1021/ja055866p. [24] D.A. Panayotov, S.P. Burrows, J.R. Morris, Photooxidation mechanism of methanol on rutile TiO2 nanoparticles, J. Phys. Chem. C 116 (2012) 6623–6635, https://doi. org/10.1021/jp209215c.
[25] C. Wang, H. Groenzin, M.J. Shultz, Direct observation of competitive adsorption between methanol and water on TiO2: an in situ sum-frequency generation study, J. Am. Chem. Soc. 126 (2004) 8094–8095, https://doi.org/10.1021/ja048165l. [26] S. Leytner, J.T. Hupp, Evaluation of the energetics of electron trap states at the nanocrystalline titanium dioxide/aqueous solution interface via time-resolved photoacoustic spectroscopy, Chem. Phys. Lett. 330 (2000) 231–236, https://doi. org/10.1016/S0009-2614(00)01112-X. [27] D.P. Colombo Jr, K.A. Roussel, J. Saeh, D.E. Skinner, J.J. Cavaleri, R.M. Bowman, Femtosecond study of the intensity dependence of electron-hole dynamics in TiO2 nanoclusters, Chem. Phys. Lett. 232 (1995) 207–214, https://doi.org/10.1016/ 0009-2614(94)01343-T. [28] G. Rothenbergei, J. Moser, M. Grátzel, N. Serpone, D.K. Sharma, Charge carrier trapping and recombination dynamics in small semiconductor particles, J. Am. Chem. Soc. 107 (1985) 8054–8059, https://doi.org/10.1021/ja00312a043. [29] R.F. Howe, M. Grátzel, EPR study of hydrated anatase under UV irradiation, J. Phys. Chem. 91 (1987) 3906–3909, https://doi.org/10.1021/j100298a035. [30] T. Xian, H. Yang, L. Di, J. Ma, H. Zhang, J. Dai, Photocatalytic reduction synthesis of SrTiO3-graphene nanocomposites and their enhanced photocatalytic activity, Nanoscale Res. Lett. 9 (2014) 1–9, https://doi.org/10.1186/1556-276X-9-327. [31] M. Pelaez, P. Falaras, V. Likodimos, K. O’Shea, A.A. de la Cruz, P.S.M. Dunlop, J.A. Byrne, D.D. Dionysiou, Use of selected scavengers for the determination of NFTiO2 reactive oxygen species during the degradation of microcystin-LR under visible light irradiation, J. Mol. Catal. A: Chem. 425 (2016) 183–189, https://doi. org/10.1016/j.molcata.2016.09.035. [32] L. He, X. Sun, F. Zhu, S. Ren, S. Wang, OH-initiated degradation and hydrolysis of aspirin in AOPs system: DFT and experimental studies, Sci. Total Environ. 592 (2017) 33–40, https://doi.org/10.1016/j.scitotenv.2017.03.041. [33] C. Adan, J.M. Coronado, R. Bellod, J. Soria, H. Yamaoka, Photochemical and photocatalytic degradation of salicylic acid with hydrogen peroxide over TiO2/SiO2 fibers, Appl. Catal. A Gen. 303 (2006) 199–206, https://doi.org/10.1016/j.apcata. 2006.02.006. [34] H. Beginejad, D. Nematollahi, F. Varmaghani, H. Shayani-Jam, Investigation of the electrochemical behavior of some dihydroxybenzoic acids in aqueous solution, Monatsh. Chem. 144 (2013) 1481–1488, https://doi.org/10.1007/s00706-0131031-6.
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DPV-assisted understanding of TiO2 photocatalytic decomposition of aspirin by identifying the role of produced reactive species
T
Abbas Valia,b, Hesam Zamankhan Malayeria, Mohammadmehdi Azizib, Hyeok Choia,* a b
Department of Civil Engineering, The University of Texas at Arlington, 416 Yates Street, Arlington, TX 76019-0308, USA Department of Chemistry and Biochemistry, The University of Texas at Arlington, 700 Planetarium Place, Arlington, TX 76019-0065, USA
ARTICLE INFO
ABSTRACT
Keywords: TiO2 photocatalysis Acetylsalicylic acid Reactive species Reaction pathway Differential pulse voltammetry
A TiO2/UV photocatalytic process generates various reactive species via different oxidation and reduction routes and thus the contribution of such reactive species to the decomposition of organic chemicals in water has been unclear. In this study, decomposition of acetylsalicylic acid (ASA) as a probe was performed to find the specific role of charge carriers (ecb− and hvb+) and primary active species (such as %O2−, %OH, %HO2, and H2O2). Differential pulse voltammetry (DPV) was introduced as an in-situ fast electrochemical method to identify produced intermediates. The study revealed that hydroxyl radicals (%OH) produced by reduction of oxygen with ecb− and subsequent primary reactions (reduction pathway) play the major role in the photocatalytic decomposition of ASA, in comparison to direct oxidation of OH− or ASA by hvb+ (oxidation pathway). The in-situ DPV analysis probed that sequential %OH addition to the ASA aromatic ring is the favorable oxidation mechanism.
1. Introduction Heterogeneous photocatalysis using semiconducting materials has been of great interest [1]. In particular, the photocatalytic decomposition mechanisms of many organic contaminants in water have been most extensively studied with titanium dioxide (TiO2) [1–5]. The process starts with generation of charge carriers at TiO2 under ultraviolet (UV) radiation, i.e., electrons in conduction band (ecb−) and holes in valence band (hvb+) [1]. Some of the charge carriers migrate to the surface of TiO2 and thus are available for chemical reactions in water. In fact, the charge carriers themselves are a strong reductant or oxidant so that organic chemicals can be directly reduced by ecb- (e.g., benzoquinone) and oxidized by hvb+ (e.g., 2-propanol) [2,6]. The photogenerated ecb− is also trapped in vacancies nearby the surface to reduce electron acceptors such as oxygen (O2) abundant in water to superoxide radical anion (%O2−) while the photogenerated hvb+ takes electrons from electron donors such as hydroxide ion (OH−) abundant in water to produce mostly hydroxyl radicals (%OH) [1]. Consequently, the reaction between the charge carriers and electron accepters and donors generates primary reactive species including %O2−, %OH, %HO2, and H2O2 (Table S1), which are, then as either oxidants or reductants, also involved in decomposition of contaminants in water, so-called secondary reactions (Table S2). Organic chemicals are decomposed via reductive pathways by ecb−
⁎
and/or %O2− while they are also decomposed via oxidative pathways by hvb+ and/or %OH. In particular, %OH, which is a strong oxidizing species, is believed to significantly contribute to decomposition of many organic chemicals [7–9]. As summarized in Fig. 1 (also Table S1), •OH can be formed through the direct oxidation of OH− by hvb+ (i.e., oxidation pathway R2) and/or the reduction of oxygen by ecb− followed by subsequent reactions (i.e., reduction pathway R3-R7) [1,7,10–13]. Previous studies showed different contributions of the oxidation pathway and the reduction pathway to the photocatalytic decomposition of aromatic chemicals (Table S3) [9,14–18]. Since the reduction pathway to produce %OH occurs in the presence of O2, those studies compared reaction kinetics in the presence (i.e., both reduction and oxidation pathways) and the absence of O2 (i.e., oxidation pathway). Salicylic acid (SA) showed a great decrease in its first order reaction rate constant from 120 × 10-3 min-1 with O2 to 2.5 × 10-3 min-1 without O2 (98 % decrease) while 2,3-dichlorophenol showed a slight decrease in its first order reaction rate constant from 1.2 × 10-3 min-1 with O2 to 1.0 × 10-3 min-1 without O2 (only 17 % decrease) [15,18]. The results indicate that the reduction pathway to generate •OH is significant for SA but not for 2,3-dichlorophenol, implying that the contribution of each of the pathways to the decomposition of organic chemicals seems to be case-specific and other reactive species than •OH might also be involved in the decomposition reaction. It is not clear which reaction pathway is more favorable for the generation of %OH
Corresponding author. E-mail address: [email protected] (H. Choi).
https://doi.org/10.1016/j.apcatb.2020.118646 Received 15 October 2019; Received in revised form 8 January 2020; Accepted 16 January 2020 Available online 16 January 2020 0926-3373/ © 2020 Elsevier B.V. All rights reserved.
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Fig. 1. Generation of reactive species from charge carriers (R1) via the primary reduction and oxidation pathways (R2-R7) on TiO2 under UV radiation. Through either the reduction pathway or the oxidation pathway, strong hydroxyl radicals are formed.
and thus for the decomposition of a specific organic chemical. In addition, as summarized in Table S2, it is less known how each of the reactive species including ecb−, hvb+, %O2−, %OH, %HO2, and H2O2 contributes to the decomposition of organic chemicals. As a result, the objective of this study is to clarify the role of the primary reactive species in the photocatalytic decomposition of an organic chemical, through systematic experimental design to quench certain reactions producing specific reactive species. Acetylsalicylic acid (ASA, known as aspirin) was carefully chosen as a probe chemical because ASA is not electroactive with hvb+ and ecb− and thus the direct role of the charge carriers in the decomposition of ASA can be largely ignored, making interpretation of ASA decomposition results less complicated [19,20]. Along with mass spectrometry (MS), differential pulse voltammetry (DPV) was introduced as an in-situ fast electrochemical method to identify produced reaction intermediates and thus propose photocatalytic decomposition mechanisms of ASA.
its hydrolysis is negligible [22]. 2.3. Electrochemical analysis In-situ DPV analysis was deployed to probe the photocatalytic decomposition intermediates of ASA. A single compartment, three-electrode cell setup, and a potentiostat (Model CHI720C) were used for the measurement. Glassy carbon electrode with 0.07 cm2 surface area was used as a working electrode. Platinum wire and Ag/AgCl/4 M KCl were used as a counter electrode and a reference electrode, respectively. Operation parameters of DPV were: initial potential = −0.4 V, final potential = 1.0 V, incremental potential = 0.004 V, amplitude = 0.05 V, pulse width = 0.05 s, sample width = 0.0167 s, and pulse period = 0.2 s. For electrochemical measurement of remaining concentration of ASA which is not electroactive during photocatalytic decomposition, an indirect method was used where ASA was hydrolyzed quickly at high pH (> 13) to SA (see reaction R18) which is electroactive [19,20]. At each time interval, 0.4 mL of solution was withdrawn and mixed with 0.4 mL of 0.4 M NaOH to indirectly measure concentration of ASA by DPV (Fig S1). This measured concentration by DPV is total concentration of ASA and SA. At the same time, concentration of SA was measured at pH 3.5 by in-situ DPV method. Remaining concentration of ASA was calculated by subtracting the concentration of SA from the total concentration. Similarly, for electrochemical measurement of remaining concentration of ASA which is not electroactive during hydrolysis at pH 3.5, in-situ DPV method was also deployed. According to reaction R18, moles of hydrolyzed ASA are equal to moles of produced SA during hydrolysis reaction. Accordingly, concentration of SA during hydrolysis of ASA was first measured by in-situ DPV method at pH 3.5 (Fig. S2). Then, remaining concentration of ASA (i.e., unhydrolyzed ASA) was calculated by subtracting the measured concentration of SA from the initial concentration of ASA. The same electrochemical set up was deployed for cyclic voltammetry analysis, where initial potential = −0.3 V, high potential = 1.2 V, final potential = -0.3 V, scan rate =5.0 mV/s, and rest time = 2.0 s.
2. Experimental methods 2.1. Chemicals and reagents ASA, SA, paracetamol (PA), 2,3-dihydroxy benzoic acid (2,3DHBA), 2,5-dihydroxy benzoic acid (2,5-DHBA), monosodium phosphate, disodium phosphate, 2-propanol, fluorine doped tin oxide (FTO), ethanol (all from Sigma Aldrich), sodium hydroxide, phosphoric acid, ammonium oxalate monohydrate (all from Alfa Aesar), and TiO2 (P25; Degussa) were used in this study, as received. Double distilled water was used for all the experiments. 2.2. Photocatalytic decomposition TiO2 (P25, Degussa) nanoparticles with 25 ± 5 nm particle size, specific surface area of 50 m2/g, and mixture of 80 % anatase and 20 % rutile was used. Concentrations of ASA and TiO2 were 1.0 mM and 1.0 g/L, respectively. Photocatalytic reactor volume was 25 mL and effective volume was 15 mL. Solution was mixed by using a magnetic stirrer. Temperature was maintained at 25 ± 1.0 °C. Solution was initially kept in dark condition for 30 min to achieve adsorption equilibrium, and then irradiated with UV for 420 min. Two 15 W low pressure mercury UV lamps emitting 365 nm were used. Lamps were mounted at 10 cm above the reactor, giving light intensity of 0.4 mW/ cm2 measured by a photonics power meter (Ophir). For experiments requiring oxygen, air was purged into the reactor while for those requiring no oxygen, nitrogen gas was purged. Purging rate was 40 mL/ min and purging was initially applied for 10 min under dark condition and then continued over UV irradiation. Buffer solution was used to prevent pH change and thus to make it easy to interpret ASA decomposition mechanism. In particular, 0.1 M phosphate buffer was used as a supporting electrolyte for in-situ electrochemical analysis of decomposition intermediates. Monosodium phosphate, disodium phosphate, and phosphoric acid with specific molar ratio were mixed to prepare 0.1 M phosphate buffer with pH of 3.5, 5.7, and 7.0 [21]. In particular, pH 3.5 was selected as a standard condition for most of experiments because direct oxidation of OH− by hvb+ to produce •OH can be significantly minimized under the acidic condition and ASA has the highest chemical stability at pH 3.5 so that
2.4. Photoelectrochemistry setting The same potentiostat (Model CHI720C) used for the measurement of the decomposition intermediates of ASA was employed. For preparation of TiO2 photoelectrode, 20 mg of TiO2 P25 was dispersed in 25 mL of ethanol by sonication for 1 min, and then FTO (sheet resistance: ∼7 Ω/sq) substrate with 2.5 cm2 surface area was placed at the bottom of TiO2 suspension for 5 min. Deposited TiO2 on FTO was washed with water and dried for 30 min at room temperature. The prepared TiO2 photoelectrode on FTO substrate was served as a working electrode. Platinum foil and Ag/AgCl/4 M KCl were used as a counter electrode and a reference electrode, respectively. The electrochemical cell was placed at 20 cm away from a xenon arc lamp. A 360 nm UV cut-off glass was placed in front of the electrochemical cell to prevent photoactivation of FTO. The potential was swept from -0.3 V to +1.0 V with low scan rate of 5.0 mV/s. Light was manually chopped in 5 s interval. Light intensity was 200 mW/cm2, which was measured using a radiant power meter (Newport 70260). 2
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2.5. Mass spectrometry
ASA, SA, and PA under different oxygen and media conditions labeled with E1-E9. In comparison between E1 and E2, it was found that there is no photocatalytic decomposition of ASA in phosphate buffer at pH 3.5 in the absence of oxygen (Table 1). It is obvious that the primary engine reaction R3 and the subsequent primary reactions R4-R7 are quenched by purging the reaction solution with nitrogen. Even the oxidation pathway was negligible, i.e., hvb+ did not directly oxidize OH− (R2) and ASA (R11). The reduction pathway in the presence of O2 was significant where ecb− plays the main role in the photocatalytic decomposition of ASA by reducing oxygen (R3) and driving subsequent reactions (R4-R7) to produce %OH. The same experiment was conducted with SA and PA (E3-E6), given that ASA is not electroactive while SA with hydroxyl group and PA with hydroxyl and amide groups might be subject to direct electrochemical oxidation. However, the results with SA and PA were very similar to those with ASA, as also reported elsewhere for SA [15]. The results indicated that ASA, SA, and PA do not have any functional groups vulnerable to direct oxidation by hvb+. In fact, several chemicals including oxalate, formate, methanol, ethanol, and 2-propanol are directly oxidized by hvb+, known as hole scavengers [23–25]. For example, methanol molecules in water are adsorbed well onto TiO2 surface and then attacked by hvb+. The main parameter for the direct oxidation of organic chemicals by hvb+ is charge-transfer rate between photocatalyst surface and organic chemicals, which is function of oxidation standard potential and adsorption behavior [1]. Although the absence of O2 can prevent oxygen-dependent primary reactions (R3-R7) and related secondary reactions (R13-R17), there is one possible secondary reaction to decompose ASA, which is direct reduction of ASA by ecb− (R12). However, the observed negligible decomposition of ASA suggests that even R12 did not occur, which is well expected, given that there is no reductable functional group in ASA to participate in a direct electron transfer reaction.
In order to further identify reaction intermediates of ASA, samples were also introduced to ion trap-time off flight MS (Shimadzu, IT-TOF), where negative electrospray ionization mode was used. Size of tube was 38" × 1/16" OD × 0.005" ID. Water with flow rate of 0.1 mL/min was used as eluent and injection volume was 2.0 μL. Other parameters of data acquisition included: range of m/z acquisition = 100 − 220 Da, acquisition time =5 min, ion accumulation time =10 ms, purge time =1 min with water, and auto-sampler rinsing volume =100 μL. 3. Results and discussion Photocatalytic decomposition of ASA occurs via two steps: (i) generation of charge carriers (ecb− and hvb+) and reactive species (•O2−, • OH, •HO2, and H2O2), so-called primary reaction (R1-R10 in Table S1) and (ii) interaction of the primary species with ASA and its reaction intermediates, so-called secondary reaction (R11-R18 in Table S2). The following sections describe modification of photocatalytic reaction conditions to reveal the contribution of the produced charge carriers and reactive species as oxidants or reductants to decomposition of ASA. 3.1. Hydrolysis of ASA To quantify the direct hydrolysis of ASA to SA and acetic acid (AcAc) (R18), the hydrolysis rate of ASA was measured at pH 3.5 in phosphate buffer (Fig. S3). The first order rate constant of its hydrolysis was 8.66 × 10−5 min-1, which is negligible, as also reported elsewhere [22]. The photocatalytic decomposition of ASA is shown in Fig. 2. Result on dark adsorption without UV indicated no significant adsorption of ASA onto TiO2. Direct photolysis of ASA without TiO2 was also negligible. ASA was significantly removed only via photocatalytic decomposition mechanism and first order reaction rate constant was at 1.77 × 10-2 min-1, which is much higher than its hydrolysis rate constant of 8.66 × 10−5 min-1. The absence of the direct hydrolysis of ASA can make it less complicated to interpret the decomposition mechanism of ASA.
3.3. Removal of buffer and oxygen One possible reason that hvb+ did not directly oxidize OH− (R2) and ASA (R11) is unfavorable adsorption of OH− and/or ASA onto the surface of TiO2. Both phosphate buffer and oxygen were removed to find the effect of phosphate buffer on the adsorption process (E7). Interestingly, ASA was decomposed in this case. However, the decomposition kinetics in E7 was two orders of magnitude slower than that in E1. Nonetheless, removing phosphate buffer caused favorable adsorption of OH− and/or ASA which can be oxidized by hvb+. In general, the results from E1-E7 indicated that the oxidation pathway to produce •OH is significantly slower than the reduction pathway. The standard potentials of associated reactions can explain the finding. Oxidation of OH− to produce %OH (Eq. 1) competes with oxygen evolution reaction (Eq. 2). Since the standard potential of OH−/%OH is higher than OH−/O2 + H2O in any pHs (note Fig. S4), oxygen evolution reaction is thermodynamically more favorable than OH−/%OH oxidation reaction. Meanwhile, oxygen reduction (Eq. 3) competes with hydrogen reduction reaction (Eq. 4). Since the standard potential of oxygen reduction at pHs above 2.71 is less than hydrogen reduction reaction (Fig. S4), oxygen reduction is more favorable than hydrogen reduction reaction.
3.2. Removal of oxygen Both the oxidation (R2) and reduction (R3-R7) pathways produce OH in presence of oxygen while only the oxidation pathway (R2) produces •OH in absence of oxygen (Fig. 1 and Table S1). Table 1 summarizes the photocatalytic decomposition reaction rate constants of
•
OH− + hvb+ ⇆ •OH or OH− ⇆ •OH + e−; E0 = +1.695 eV vs. NHE (1) 4OH− + 4hvb+ ⇆ O2 + 2H2O or 4OH− ⇆ O2 + 2H2O + 4e- ; E0 = +0.401 eV vs. NHE (2) O2 + ecb− ⇆ •O2−; E0 = −0.160 V vs. NHE 2H Fig. 2. UV photolytic (-◼-) and UV/TiO2 photocatalytic decomposition of 1.0 mM ASA (..►..) at pH 3.5.
+
+
2ecb−
0
⇆ H2; E = 0.0 V vs. NHE
(3) (4)
In addition, the conduction band electron trapping by oxygen is much faster than valence band hole trapping by OH− because the 3
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Table 1 Photocatalytic decomposition of ASA, SA, and PA under different oxygen and media conditions. Experiment No.
Chemical
O2 (mg/L)
Media
Reaction rate constant (min−1)a
E1 E2 E3 E4 E5 E6 E7 E8
ASA ASA SA SA PA PA ASA ASA
8.4 0 8.4 0 8.4 0 0 8.4
1.77 NDb 1.81 ND 2.31 ND 1.86 4.81
E9
ASA
8.4
0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water 0.1 M phosphate buffer in water Water 0.1 M phosphate buffer + 0.1 M ammonium oxalate in water 0.1 M phosphate buffer + 0.1 M 2-propanol in water
a b
± 0.01 × 10−2 ± 0.01 × 10−2 ± 0.02 × 10−2 ± 0.01 × 10−4 ± 0.02 × 10−3
ND
Note that pH 3.5 was used for all experiments. Not determined: Reaction within the time frame was negligible to determine the rate constant.
conduction band electron has lower effective mass than valence band hole, as verified elsewhere [26–29]. However, the reasons mentioned above might be more reasonable when organic chemicals are less favorable for adsorption onto the surface of TiO2 than OH−. If organic chemicals adsorb well onto TiO2 and have low oxidation potential, the direct oxidation of organic chemicals by hvb+ should be considered.
3.5. Addition of %OH and hole scavenger The reactions (R3-R7) produce the primary active specious (%O2−, OH, %HO2, and H2O2), which directly react with ASA (R13-R17). To find the contribution of the primary species to ASA decomposition, 2propanol was added to phosphate buffer at pH 3.5 (E9) because 2propnaol is a strong %OH scavenger to prevent R16 and R17 [9,31]. In addition, 2-propanol is a hvb+ scavenger to inhibit R2 and R11 [23–25]. 2-proanol is directly oxidized to acetone (Eq. 6). As shown in Fig. 3, the presence of 2-propanol produced almost 1.5 times higher photocurrent than ammonium oxalate, indicating that hvb+ reacts with 2-propanol faster than ammonium oxalate. No decomposition of ASA was observed in the presence of 2-propanol (Table 1). It seems that 2-propanol successfully acted as a hole and •OH scavenger to prevent ASA decomposition.
%
3.4. Addition of hole scavenger Previously, in order to find contribution of the oxidation pathway, the reduction pathway to produce %OH was killed with no O2, concluding that %OH production by the reduction pathway in the presence of O2 is significant. Inversely, the oxidation pathway (R2) was quenched (E8) by adding 0.1 M ammonium oxalate as a hole scavenger to confirm the results of E1-E7 [30]. As shown in Fig. 3, the significant increase in photocurrent by adding ammonium oxalate into phosphate buffer at pH 3.5 confirmed that hvb+ can immediately oxidize oxalate (Eq.5). Consequently, the observed photocatalytic decomposition of ASA (reaction rate constant at 4.81 × 10−3 min-1) even in the presence of the hole scavenger indicated that ecb- drives the photocatalytic decomposition of ASA.
C3H6O + 2H+ + 2e− ⇆ C3H8O or C3H6O + 2H+ ⇆ C3H8O + 2hvb+; E0 = +0.13 V vs. NHE (6) 3.6. DPV-assisted understanding of ASA decomposition mechanism Based on the results from E1-E9, it can be concluded that •OH exclusively produced by reduction of oxygen by ecb− (R3) and subsequent primary reactions (R4-R7), so-called reduction pathway, was responsible for the photocatalytic decomposition of ASA. To identify reaction intermediates and thus to propose ASA decomposition mechanisms by •OH, in-situ electrochemical analysis DPV was performed along with MS analysis. Fig. 4 shows DPV analysis of intermediates formed during the photocatalytic decomposition of ASA and SA for 90 min, exhibiting clear five anodic peaks at 0.166, 0.342, 0.442, 0.842, and 0.982 V. In addition, MS analysis of the same samples showed m/z values of 137, 153, 167, 169, 179, and 195 (Table S4). The peak at 0.982 V was attributed to the presence of SA because the reference DPV voltamomogram of SA confirmed the peak (Fig. S5). The m/z at 137 also confirmed the presence of SA as one of reaction intermediates from ASA. As mentioned early, direct hydrolysis of ASA to SA was negligible at pH 3.5 and thus the formation of SA can be explained by %OH attack to ASA, in particular the single bond between oxygen and carbon (C10O11, note Fig. 5a) in ASA with the lowest bond dissociation energy at 78.0 kcal/mol [32]. All DPV peaks for ASA and SA were the same except for the peak at +0.842 V only in ASA while all MS m/z peaks for ASA and SA were the same except for the peaks at 179 and 195 only in ASA. The m/z at 179 indicated ASA and thus m/z at 195 suggested the presence of another noticeable intermediate from ASA. •OH addition to aromatic groups is an electrophilic substitution (note carbons 3, 4, 5, and 6 in Fig. 5a). Since the ester functional group in ASA is an activator (i.e., donating electron to the ring), it tends to drive incoming %OH to the ortho or para
2CO2 + 2e− ⇆ C2O42- or 2CO2 ⇆ C2O42- + 2hvb+; E0 = +0.51 V vs. NHE (5)
Fig. 3. Photocurrent vs. potential at pH = 3.5 for (i) 0.1 M phosphate buffer, (ii) 0.1 M phosphate buffer and 0.1 M ammonium oxalate, and (iii) 0.1 M phosphate buffer and 0.1 M 2-propanol. 4
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DPV peaks (e.g., 0.166, 0.342, and 0.442 V in Fig. 4) and m/z peaks (e.g., 153, 167, and 169 in Table S4) in ASA and SA, •OH attack mechanisms to SA and HASA were proposed. Since hydroxyl group is an activator and carboxylic group is deactivator and thus both functional groups in SA direct •OH to carbon 3 or 5 (Fig. 5c), formation of 2,5DHBA and 2,3-DHBA is highly possible (Fig. 5e and f) [33]. A comparison between the in-situ voltammogram in Fig. 4 and reference voltammograms of 2,5-DHBA and 2,3-DHBA (Fig. S5) confirmed the presence of 2,5-DHBA and 2,3-DHBA formed from SA. The m/z peak at 153 also verified the presence of 2,5-DHBA and 2,3-DHBA. Meanwhile, formation of 2,5-DHBA and 2,3-DHBA can also be originated from 5HASA and 3-HASA, respectively because the single bond C10-O11 in 3HASA (Fig. 5d) and 5-HASA (Fig. 5b) is vulnerable to %OH attack [32]. To explain the m/z peak 169 (167 later), trihydroxybenzoic acid (THBA; MW of 170, Fig. 5g) formed by •OH addition to 2,5-DHBA and 2,3-DHBA was proposed. Cyclic voltammograms of 2,3-DHBA and 2,5DHBA at pH 3.5 are shown in Fig. 7. In the first cycle, one oxidation peak, A1 for 2,5-DHBA and A2 for 2,3-DHBA, was revealed. The counterpart reduction peaks for A1 and A2 are C1 and C2, respectively, indicating that 2,5-DHBA and 2,3-DHBA can be oxidized to counterpart oxidized forms, as shown in Fig. 8, i.e., 8f (2,5-DHBA) to 8 g and 8a (2,3-DHBA) to 8b in oxidation vice versa in reduction. However, oxidation of 2,5-DHBA and 2,3-DHBA is not reversible at pH 3.5 because there is an additional cathodic peak, C3 for 2,5-DHBA and C4 for 2,3DHBA. The peaks C3 and C4 disappeared gradually by increasing pHs from 3.5–7.0 (note Fig. S6), as also observed elsewhere [34]. Disappearing C3 and C4 peaks at high pHs imply that 8b and 8 g promptly react with H+ and H2O under such a low pH of 3.5 to produce chemicals 8d and 8i. Meanwhile, the oxidized counterparts of C3 and C4 are A3 and A4 in the second cycle. The peak potentials of C3/C4 couple and A3/A4 couple were the same, suggesting that possibly oxidation product of 2,5-DHBA and 2,3-DHBA is the same as 2,3,5-THBA (8e). The m/z peaks in MS analysis corresponded to the chemicals 8d and 8i (m/z 167) and 2,3,5-THBA (m/z 169). DPV analysis to prove time evolution of primary intermediates (HASA and SA) and secondary intermediates (2,3-DHBA, 2,5-DHBA, and 2,3,5-THBA) is shown in Fig. 9 while concentration-time evolution of primary and secondary reaction intermediates formed by the photocatalytic decomposition of ASA is summarized in Fig. 10. The primary intermediates, SA and HASA, reached the highest concentrations after around 60 min and they were completely decomposed after 360 min
Fig. 4. In-situ electrochemical DPV analysis of intermediates from the photocatalytic decomposition of 1.0 mM ASA and 1.0 mM SA at pH = 3.5 for 90 min.
position (carbon 3 or 5 in Fig. 5a). However, the carboxylic acid moiety as an electron withdrawing group directs incoming %OH to the meta position (carbon 3 or 5 in Fig. 5a). As a result, both functional groups (i.e., carboxyl and ester) direct incoming %OH to the carbon 3 or 5 in ASA molecule to form 3-hydroxyacetylsalicylic acid (3-HASA shown in Fig. 5d) or 5-hydroxyacetylsalicylic acid (5-HASA shown in Fig. 5b). The m/z peak at 195 indicated the presence of such HASAs. The DPV peak at +0.842 V could be ascribed to the HASAs. Fig. 6 shows DPV analysis of primary intermediates (SA and HASA) formed by the photocatalytic decomposition of ASA at pH = 3.5, 5.7, and 7.0. SA and HASA (3-HASA or 5-HASA) peaks are shifted to low potentials over pH increase with a slope of -51.1 mV/pH for SA and -56.5 mV/pH for HASA. The relationship between peak potential and pH is shown in Eq. 7, where m and n are the numbers of protons and electrons, respectively. When the numbers of protons and electrons are identical, the theoretical slope is 59 mV/pH, which is very close to the results at 51.1 and 56.5. As a result, the number of involved protons was almost the same as that of electrons for oxidation of SA and HASA. Ep = E p
(pH=0)
– (0.059 m/n) pH
(7)
To identify more reaction intermediates corresponding to common
Fig. 5. Proposed photocatalytic decomposition pathways of ASA at pH 3.5 by hydroxyl radical attack mechanisms. 5
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Fig. 6. (a) In-situ electrochemical DPV analysis of primary intermediates (SA and HASA) formed by the photocatalytic decomposition of 1.0 mM ASA in 60 min at pH = 3.5, 5.7, and 7.0 and b) DPV peak potential shift of SA and HASA over pHs.
4. Conclusions This study provided a systematic experimental design to quench certain reactions producing specific reactive species and thus to find the role of such reactive species in TiO2 photocatalytic decomposition of organic chemicals in water. ASA was selected as a probe chemical. The results revealed that %OH production by reduction of oxygen with ecb− and subsequent primary reactions (reduction pathway) played the major role in the photocatalytic decomposition of ASA as a probe chemical, in comparison to direct oxidation of OH− or ASA by hvb+ (oxidation pathway). In-situ electrochemical DPV analysis greatly helped to understand the decomposition mechanisms and pathways of ASA and thus to probe that sequential •OH addition to the ASA aromatic ring is the favorable oxidation mechanism, forming SA, 3-HASA, 5HASA, 2,3-DHBA, 2,5-DHBA, and 2,3,5-THBA. Findings from this study would help to set a strategy to control the TiO2/UV operating parameters for generating effective reactive species, leading to favorable reaction pathways and fast decomposition of organic contaminants in water.
Fig. 7. Cyclic voltammogram of 0.1 mM 2,3-DHBA (1 st and 2nd cycles) and 0.1 mM 2,5-DHBA (1 st and 2nd cycles) at 25 mV/s.
CRediT authorship contribution statement
while the secondary intermediates reached the highest concentrations after around 180 min for 2,3-DHBA and 2,5-DHBA and 240 min for 2,3,5-THBA and they were completely decomposed after 420 min. Although we have not investigated in detail beyond the point, further oxidation of 2,3,5-THBA may cause aromatic ring cleavage to yield carboxylic acids such as malic acid and succinic acid.
Abbas Vali: Conceptualization, Data curation, Investigation, Writing - original draft. Hesam Zamankhan Malayeri: Methodology. Mohammadmehdi Azizi: Methodology. Hyeok Choi: Supervision, Writing - review & editing.
Fig. 8. Possible oxidation mechanism of 2,3-DHBA and 2,5-DHBA at pH 3.5. A1-A4 and C1-C4 are oxidation and reduction CV peaks, respectively, which are shown in Fig. 7.
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Fig. 9. In-situ electrochemical DPV analysis on time evolution of (a) primary intermediates (HASA and SA) and (b) secondary intermediates (2,3-DHBA, 2,5-DHBA, and 2,3,5THBA) formed by the photocatalytic decomposition of 1.0 mM ASA at pH = 3.5 for (a1) 0−60 min and (a2) 60−360 min and (b1) 0−180 min and (b2) 240−420 min. Please note two different time slots were purposely used to show concentration increase by up to around 60 min for (a) and 180 min for (b) and then concentration decrease.
Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.apcatb.2020.118646. References [1] J. Schneider, M. Matsuoka, M. Takeuchi, J. Zhang, Y. Horiuchi, M. Anpo, D.W. Bahnemannt, Understanding TiO2 photocatalysis: mechanisms and materials, Chem. Rev. 114 (2014) 9919–9986, https://doi.org/10.1021/cr5001892. [2] M.R. Hoffmann, S.T. Martin, W. Choi, D.W. Bahnemannt, Environmental applications of semiconductor photocatalysis, Chem. Rev. 95 (1995) 69–96, https://doi. org/10.1021/cr00033a004. [3] K. Hashimoto, H. Irie, A. Fujishima, TiO2 photocatalysis: a historical overview and future prospects, Japan J. Appl. Phys. 44 (2005) 8269–8285, https://doi.org/10. 1143/JJAP.44.8269. [4] A. Mills, S.L. Hunte, An overview of semiconductor photocatalysis, J. Photochem. Photobiol. A 108 (1997) 1–35, https://doi.org/10.1016/S1010-6030(97)00118-4. [5] Y. Nosaka, A.Y. Nosaka, Generation and detection of reactive oxygen species in photocatalysis, Chem. Rev. 117 (2017) 11302–11336, https://doi.org/10.1021/acs. chemrev.7b00161. [6] M.A. Henderson, M. Shen, Electron-scavenging chemistry of benzoquinone on TiO2 (110), Top. Catal. 60 (2017) 440–445, https://doi.org/10.1007/s11244-0160707-7. [7] Y. Nosaka, A. Nosaka, Understanding hydroxyl radical (%OH) generation processes in photocatalysis, ACS Energy Lett. 1 (2016) 356–359, https://doi.org/10.1021/ acsenergylett.6b00174. [8] K.I. Ishibashia, A. Fujishima, T. Watanabea, K. Hashimoto, Quantum yields of active oxidative species formed on TiO2 photocatalyst, J. Photochem. Photobiol. A 134 (2000) 139–142, https://doi.org/10.1016/S1010-6030(00)00264-1. [9] L. Yang, L.E. Yu, M.B. Ray, Photocatalytic oxidation of paracetamol: dominant reactants, intermediates, and reaction mechanisms, Environ. Sci. Technol. 43 (2009) 460–465, https://doi.org/10.1021/es8020099. [10] P.M. Wood, The two redox potentials for oxygen reduction to superoxide, Trends Biochem. Sci. 12 (1987) 250–251, https://doi.org/10.1016/0968-0004(87) 90123-X. [11] P.S. Rao, E. Hayon, Redox potentials of free radicals. IV. superoxide and hydroperoxy radicals•O2− and •HO2, J. Phys. Chem. 79 (1975) 397–402, https://doi.org/ 10.1021/j100571a021. [12] D.R. Lide, CRC Handbook, 84th ed., CRC Press, Boca Raton, Florida, 2003, pp. 5:1–5:2 5:10. [13] J.A. Dean, Lange’s Handbook of Chemistry, 11th ed., McGraw-Hill, New York, 1979 9:4-9:128.
Fig. 10. Concentration-time evolution of primary and secondary reaction intermediates formed by the photocatalytic decomposition of 1.0 mM ASA at pH = 3.5.
Declaration of Competing Interest The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper. Acknowledgements This research was supported by the University of Texas at Arlington through new faculty startup funds. The authors thank Dr. Krishnan Rajeshwar for allowing us to use his electrochemistry laboratory at the University of Texas at Arlington.
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A. Vali, et al. [14] L. Yang, L.E. Yu, M.B. Ray, Degradation of paracetamol in aqueous solutions by TiO2 photocatalysis, Water Res. 42 (2008) 3480–3488, https://doi.org/10.1016/j. watres.2008.04.023. [15] C. Karunakaran, B. Naufal, P. Gomathisankar, Efficient photocatalytic degradation of salicylic acid by bactericidal ZnO, J. Korean Chem. Soc. 56 (2012) 108–114, https://doi.org/10.5012/jkcs.2012.56.1.108. [16] S. Zheng, Y. Cai, K.E. O’Shea, TiO2 photocatalytic degradation of phenylarsonic acid, J. Photochem. Photobiol. A 210 (2010) 61–68, https://doi.org/10.1016/j. jphotochem.2009.12.004. [17] D.-H. Tseng, L.-C. Juang, H.-H. Huang, Effect of oxygen and hydrogen peroxide on the photocatalytic degradation of monochlorobenzene in TiO2 aqueous suspension, Int J Photoenergy 1 (2012) 1–9, https://doi.org/10.1155/2012/328526. [18] H.C. Liang, X.Z. Li, Y.H. Yang, K.H. Sze, Effects of dissolved oxygen, pH, and anions on the 2,3-dichlorophenol degradation by photocatalytic reaction with anodic TiO2 nanotube films, Chemosphere 73 (2008) 805–812, https://doi.org/10.1016/j. chemosphere.2008.06.007. [19] E.R. Sartori, R.A. Medeiros, R.C. Rocha-Filho, O. Fatibello-Filho, Square-wave voltammetric determination of acetylsalicylic acid in pharmaceutical formulations using a boron-doped diamond electrode without the need of previous alkaline hydrolysis step, J. Braz. Chem. Soc. 20 (2009) 360–366, https://doi.org/10.1590/ S0103-50532009000200022. [20] C. Cofan, C. Radovan, Anodic determination of acetylsalicylic acid at a mildly oxidized boron-doped diamond electrode in sodium sulphate medium, Int. J. Electrochem. Sci. 1 (2011) 1–9, https://doi.org/10.4061/2011/451830. [21] Council of Europe, European Pharmacopoeia Commission, The European Pharmacopoeia, 6th ed., Conseil de l’Europe, Strasburg, France, 2007, pp. 508–514. [22] C.A. Kelly, Determination of the decomposition of aspirin, J. Pharm. Sci. 59 (1979) 1053–1079, https://doi.org/10.1002/jps.2600590802. [23] Y. Tamaki, A. Furube, M. Murai, K. Hara, R. Katoh, M. Tachiya, Direct observation of reactive trapped holes in TiO2 undergoing photocatalytic oxidation of adsorbed alcohols: evaluation of the reaction rates and yields, J. Am. Chem. Soc. 128 (2006) 416–417, https://doi.org/10.1021/ja055866p. [24] D.A. Panayotov, S.P. Burrows, J.R. Morris, Photooxidation mechanism of methanol on rutile TiO2 nanoparticles, J. Phys. Chem. C 116 (2012) 6623–6635, https://doi. org/10.1021/jp209215c.
[25] C. Wang, H. Groenzin, M.J. Shultz, Direct observation of competitive adsorption between methanol and water on TiO2: an in situ sum-frequency generation study, J. Am. Chem. Soc. 126 (2004) 8094–8095, https://doi.org/10.1021/ja048165l. [26] S. Leytner, J.T. Hupp, Evaluation of the energetics of electron trap states at the nanocrystalline titanium dioxide/aqueous solution interface via time-resolved photoacoustic spectroscopy, Chem. Phys. Lett. 330 (2000) 231–236, https://doi. org/10.1016/S0009-2614(00)01112-X. [27] D.P. Colombo Jr, K.A. Roussel, J. Saeh, D.E. Skinner, J.J. Cavaleri, R.M. Bowman, Femtosecond study of the intensity dependence of electron-hole dynamics in TiO2 nanoclusters, Chem. Phys. Lett. 232 (1995) 207–214, https://doi.org/10.1016/ 0009-2614(94)01343-T. [28] G. Rothenbergei, J. Moser, M. Grátzel, N. Serpone, D.K. Sharma, Charge carrier trapping and recombination dynamics in small semiconductor particles, J. Am. Chem. Soc. 107 (1985) 8054–8059, https://doi.org/10.1021/ja00312a043. [29] R.F. Howe, M. Grátzel, EPR study of hydrated anatase under UV irradiation, J. Phys. Chem. 91 (1987) 3906–3909, https://doi.org/10.1021/j100298a035. [30] T. Xian, H. Yang, L. Di, J. Ma, H. Zhang, J. Dai, Photocatalytic reduction synthesis of SrTiO3-graphene nanocomposites and their enhanced photocatalytic activity, Nanoscale Res. Lett. 9 (2014) 1–9, https://doi.org/10.1186/1556-276X-9-327. [31] M. Pelaez, P. Falaras, V. Likodimos, K. O’Shea, A.A. de la Cruz, P.S.M. Dunlop, J.A. Byrne, D.D. Dionysiou, Use of selected scavengers for the determination of NFTiO2 reactive oxygen species during the degradation of microcystin-LR under visible light irradiation, J. Mol. Catal. A: Chem. 425 (2016) 183–189, https://doi. org/10.1016/j.molcata.2016.09.035. [32] L. He, X. Sun, F. Zhu, S. Ren, S. Wang, OH-initiated degradation and hydrolysis of aspirin in AOPs system: DFT and experimental studies, Sci. Total Environ. 592 (2017) 33–40, https://doi.org/10.1016/j.scitotenv.2017.03.041. [33] C. Adan, J.M. Coronado, R. Bellod, J. Soria, H. Yamaoka, Photochemical and photocatalytic degradation of salicylic acid with hydrogen peroxide over TiO2/SiO2 fibers, Appl. Catal. A Gen. 303 (2006) 199–206, https://doi.org/10.1016/j.apcata. 2006.02.006. [34] H. Beginejad, D. Nematollahi, F. Varmaghani, H. Shayani-Jam, Investigation of the electrochemical behavior of some dihydroxybenzoic acids in aqueous solution, Monatsh. Chem. 144 (2013) 1481–1488, https://doi.org/10.1007/s00706-0131031-6.
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