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Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices
Accepted Manuscript Title: Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices Authors: Sofia Dimitriadou, Zacharias Frontistis, Athanasia Petala, Georgios Bampos, Dionissios Mantzavinos PII: DOI: Reference:
S0920-5861(19)30058-6 https://doi.org/10.1016/j.cattod.2019.02.025 CATTOD 11977
To appear in:
Catalysis Today
Received date: Revised date: Accepted date:
28 November 2018 1 February 2019 12 February 2019
Please cite this article as: Dimitriadou S, Frontistis Z, Petala A, Bampos G, Mantzavinos D, Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices, Catalysis Today (2019), https://doi.org/10.1016/j.cattod.2019.02.025 This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.
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Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices
Sofia Dimitriadou1, Zacharias Frontistis2, Athanasia Petala1,3*, Georgios Bampos1, Dionissios Mantzavinos1,3
Department of Chemical Engineering, University of Patras, Caratheodory 1, University Campus,
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GR-26504 Patras, Greece
Department of Environmental Engineering, University of Western Macedonia, GR-50100 Kozani,
Greece 3
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INVALOR: Research Infrastructure for Waste Valorization and Sustainable Management,
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Caratheodory 1, University Campus, GR-26504 Patras, Greece
*
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Corresponding author: E-mail: [email protected]
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CB
OH
SPS 1.0 0.8
C/C0
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Graphical abstract
0.6 0.4
SPS CB CB/SPS
0.2 0.0
0
5
10
15
20
Time (min)
25
30
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Highlights
Carbon black shows good activity as a low-cost persulfate activator for diclofenac oxidation
Complex water matrix interactions determine degradation rates
Diclofenac adsorption on carbon black surface is pH-sensitive, but oxidation is not
Coupling carbon black with light or ultrasound enhances degradation synergistically
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Αbstract
In this study, carbon black (CB) was employed as a heterogeneous activator for the conversion of
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sodium persulfate (SPS) to reactive species for the degradation of drug diclofenac (DCF).
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Experiments were performed at DCF concentrations between 0.5 and 4 mg/L, CB concentrations between 25 and 75 mg/L and SPS concentrations between 25 and 200 mg/L. Degradation rates, based
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on pseudo-first order kinetics, generally increased with decreasing DCF and increasing CB concentrations. The rate also increased with increasing SPS concentration up to 50 mg/L and decreased at higher values due to scavenging effects. Besides experiments in ultrapure water (UPW),
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real matrices were tested (i.e. bottled water (BW), surface water (SW), secondary treated wastewater (WW)), as well as UPW spiked with bicarbonate (50-500 mg/L), chloride (100-500 mg/L) or humic
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acid (10-20 mg/L). Degradation rates decreased with increasing matrix complexity, while the addition of chloride or humic acid was detrimental to the process; on the contrary, bicarbonate at 500 mg/L enhanced DCF degradation rate nearly five-fold. The effect of initial solution pH was also studied in
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the range 3-9.5 showing that degradation was not pH-sensitive. Experiments were also performed activating SPS by simulated solar radiation or 20 kHz ultrasound with or without CB. Coupling activators (i.e. CB with solar light or CB with ultrasound) favored DCF degradation in a synergistic way, with the level of synergy being 45-50%. Keywords: carbocatalysis; kinetics; solar light; synergy; ultrasound; water matrix Introduction
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Pharmaceuticals and personal care products constitute a large group of synthetic organic chemicals inextricably linked to everyday life that have been categorized as contaminants of global concern [1]. These concerns derive from the fact that trace levels of the above-mentioned compounds have been quantified in surface and wastewaters, pointing out the inefficiency of classical physicochemical and biological treatment methods used in conventional wastewater treatment plants [2]. Diclofenac (DCF) is a nonsteroidal anti-inflammatory drug contained in medications under a variety of trade names that has been detected in water bodies [3]. In particular, European Union included DCF in the list of
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priority substances (2013/39/EU) and set the acceptable environmental limits to 100 ng/L for inland and 10 ng/L for coastal waters [4]. Adverse effects to aqueous ecosystems have already been reported due to prolonged exposure to DCF [5], thus creating major concerns for the impact on human health.
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To address this problem, the development of effective methods, such as the advanced oxidation processes (AOPs), for DCF and other pharmaceuticals removal has gained significant interest. AOPs are a family of similar but not identical technologies relying on the formation of various oxidative
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species and, especially, hydroxyl radicals (•OH), which are characterized by an oxidation potential of 2.8 V vs. NHE at pH ~ 6 [6].
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Recently, the activity of sulfate radicals (𝑆𝑂4•− ), which are characterized by an equally high
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oxidation potential (2.6 V vs NHE at pH~6) has been studied in detail; sulfate radicals often lead to
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a more efficient decomposition mechanism [7,8], while some of their advantageous characteristics include longer half-life [9] and reactivity over a wide pH range [10].
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Sulfate radicals can be produced after activation of compounds belonging to the group of persulfates. Generally, persulfates are of comparable cost to hydrogen peroxide in the order of 1 $/kg but they are easier to handle due to their crystalline form, thus limiting problems related to storage
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and transportation [11]. Common activation methods include heat [12] and ultraviolet (UV) radiation [13], that can provide energy for the fission of O-O bond in the persulfate structure and the formation
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of two sulfate radicals according to Equation 1 [14]. In addition, persulfates can be activated via redox reactions with transition metals in both heterogeneous and homogeneous systems leading to the formation of one sulfate radical (Equation 2) [15]. In both cases, 𝑆𝑂4•− may further react resulting in the formation of additional oxidizing species such as •OH (Equation 3). (1)
S2O82- +Mn+ SO42- +SO-•4 +M(n+1)+
(2)
SO-•4 +OH- SO42- + • OH
(3)
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hv/heat S2O82- 2SO-•4
Αctivation by transition metals reduces significantly the energy requirements of the process and consequently the whole cost, making this activation method more attractive for practical applications [14]. Moreover, special emphasis has been given towards the development of heterogeneous,
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transition metal-based activation systems, thus avoiding secondary pollution due to the occurrence of metal ions in aqueous solutions [16]. In particular, copper- and cobalt-based catalysts have been used for heterogeneous activation resulting in successful degradation of many persistent micro-pollutants [17,18]. However, the reported activation mechanisms include irreversible changes in the oxidation state of the metal, thus challenging the true heterogeneous catalytic nature of the process and limiting the privileges of collection and reuse of the catalyst [19]. In order to avoid such phenomena, metal-free carbonaceous materials have been proposed as
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alternative persulfate activators. Activated carbon [20], carbon nanotube [21], reduced graphene oxide [22] and graphene [23], multi-walled carbon nanotubes [24] and biochars [25,26] are some characteristic examples. Considering carbonaceous-based activation systems, several studies propose
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a non-radical oxidation pathway for organic degradation, proceeding through the formation of reactive complexes between activators and the persulfate and subsequent electron abstraction by these complexes [24,27,28]. This mechanism can be described according to the following equation (Εqn 2
S2O8 [OP] 2SO2-4 +CC
(4)
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CC OP CC(e )
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4), where CC is the carbocatalyst and OP the organic pollutant.
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Carbon black (CB), mainly consisting of fine particles of carbon, has recently been proposed as a low-cost alternative to expensive carbon materials, such as graphene, in the field of electrochemistry
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showing great performance as an electric conductive agent [29]. In particular, the current worldwide CB production exceeds 8 Mt/y, with 90% used in rubber applications. This has a positive impact on
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its commercial price, being several orders of magnitude lower than other carbonaceous materials. A complicated blend of carbon chemistry and surface energy makes it ideal for a variety of applications
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explaining its ever-increasing popularity and huge utility. In the present work, the efficiency of CB for SPS activation and the subsequent degradation of DCF is tested. To the best of our knowledge,
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only o handful of studies report DCF removal from water matrices with heterogeneous persulfate activation systems, such as metal-doped, manganese oxide octahedral molecular sieve catalysts [30], perovskite oxides [31] and graphene oxide-TiO2 nanosheets [32], whereas this is the first work using CB instead of high-priced carbonaceous materials for SPS activation. The effects of CB, SPS and
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DCF concentrations on DCF degradation kinetics are investigated. Emphasis is given on the effect of water matrix, while possible synergistic effects arising from the simultaneous use of solar or ultrasound irradiation are also discussed.
2. Experimental 2.1. Chemicals and water matrices
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Carbon black (Vulcan XC72R) was supplied by Cabot. Diclofenac sodium salt (C14H10Cl12NNaO2, CAS:15307-79-6) and sodium persulfate (Na2S2O8, CAS: 7775-27-1) were purchased from Sigma Aldrich. Potassium iodide (KI, CAS:7681-11-0) for SPS consumption measurements and for radical quenching experiments was purchased from Carlo Erba reagents. Hydrogen peroxide (H2O2, CAS: 7722-84-1), sodium chloride (NaCl, CAS:7647-14-5), sodium bicarbonate (NaHCO3, CAS: 144-558), humic acid (HA, CAS:1415-93-6), methanol (CH3OH, CAS: 67-56-1), ethanol (CH3CH2OH, CAS: 64-17-5) sodium hydroxide (NaOH, CAS: 1310-73-2), sulfuric acid (H2SO4, CAS: 7664-93-9)
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and acetonitrile (CH3CN, CAS: 75-05-8, for HPLC analysis) were also purchased from Sigma Aldrich.
Ultrapure water (pH=6, EASYpureRF-Barnstead/Thermolyne, USA) was employed in most of the
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experiments presented in this study. Moreover, bottled water (BW) (pH=7.8, 0.4 mS/cm conductivity, 182 mg/L HCO3- , 2 mg/L Cl- , 14 mg/L SO24 , 8.6 mg/L NO3- and 76 mg/L of various metal ions), surface water (SW) taken from a river in the region of Athens, Greece (pH=7.5, 491 μS/cm
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conductivity, 2.7 mg/L total organic carbon, 274 mg/L SO24 , 5 mg/L Cl- ) and secondary treated
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wastewater (WW) (pH=8, 21 mg/L chemical oxygen demand) were also used in order to study the
2.2 Catalyst characterization
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effect of water matrix.
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CB was characterized with respect to its specific surface area according to the Brunaueur-EmmettTeller (BET) method with the use of a Micromeritics (Gemini III 2375), whereas Brucker D8 Advance (Cu Kα) was used for X-Ray diffraction pattern of the sample. The zeta potential
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measurements of CB were performed via laser doppler micro-electrophoresis employing a Malvern Zetasizer (Nano-ZS). Briefly, small amounts of CB powder were ultrasonically dispersed in UPW.
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The pH of these suspensions was adjusted using H2SO4 or NaOH solutions. The velocity of the CB particles of each suspension was measured implementing a patented laser interferometric technique (phase analysis light scattering), enabling the calculation of electrophoretic mobility of the particles
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and, thus, the estimation of the zeta potential values. Fourier transform infrared (FTIR) spectroscopy was performed using a Perkin Elmer Spectrum RX FTIR system. The measurement range was 4000– 450 cm-1. Vulcan (about 0.5%) and KBr were sieved and pressurized to produce a homogeneous disk.
2.3 Experimental procedure
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Degradation experiments were carried out in a pyrex vessel of 250 mL capacity open to the atmosphere. In a typical run, 120 mL of an aqueous solution with the desired concentration of DCF (1 mg/L in most cases) was loaded in the vessel followed by the addition of CB and SPS. The solution was under continuous magnetic stirring. Samples of 1.2 mL were withdrawn from the vessel at desired time periods followed by the addition of 0.3 mL of methanol in order to quench the reaction. Samples were then filtered (0.22 μm PVDF filters) to remove CB particles and DCF concentration
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was measured with the use of high performance liquid chromatography (HPLC).
2.4 Analytical methods
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An HPLC system (Waters Alliance 2695) consisting of a photodiode array detector (Waters 2996), a gradient pump (Waters 2695) and a Kinetex column (C18 100A, 150 mm × 3 mm, 2.6 μm particle size) maintained at 45oC was employed. The mobile phase consisted of 68:32 water:acetonitrile,
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while the injection volume was 40 μL.
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Persulfate was measured spectrophotometrically based on its reaction with potassium iodide [33]. A solar simulator (Oriel, model LCS-100) equipped with a 100 W Xenon (ozone free) lamp and a
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Branson 450 horn-type digital sonifier operating at a frequency of 20 kHz and an actual power density
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of 36 W/L were employed to study the combined activation of persulfate by CB/solar light or
3. Results
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3.1 Characterization of CB
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CB/ultrasound. Details regarding the solar simulator and the sonifier can be found elsewhere [25].
The XRD pattern of CB is presented in Figure 1A. It consists of reflections at 24o and 43.5o attributed to graphitic regions (JCPDS 2-456) [34]. Measurements of zeta potential (Figure 1B) show
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that the CB surface is negatively charged at pH>2 with the negative charge increasing at more alkaline conditions. The specific surface area of the sample was measured equal to 216 m2/g. The FTIR spectrum of CB is shown in Figure 1C. The peak at 3450 cm–1 is assigned to stretching vibration of
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the –OH groups of CB. As can be seen, the surface of CB exhibits a significant number of –OH groups. The absence of a peak at about 1100 cm–1 denotes that there are no C-O-C bonds. The peaks at about 1640 and 1380 cm–1 correspond to C=C and C-OH groups.
3.2 Screening experiments
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Figure 2A shows normalized DCF concentration-time profiles for experiments where CB and SPS were employed either alone or together. CB (25 mg/L) alone led to ca 20% decrease in DCF concentration after 30 min presumably due to adsorption phenomena. A comparable concentration decrease was observed when SPS (100 mg/L) was used alone due its mild oxidizing ability. On the contrary, the combined application of CB and SPS led to complete degradation of 1 mg/L DCF after 15 min, indicating the successful activation of SPS by CB. In addition, Figure 2B shows the conversion of 100 mg/L SPS in the presence of 25 mg/L CB and 1 mg/L DCF. It is observed that ca
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30% of SPS was consumed within the first 5 min of reaction, while its concentration remained almost constant thereafter.
The efficiency of SPS for DCF degradation was compared to that of H2O2. As seen in Figure 3,
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SPS at 100 mg/L was more effective than H2O2 at 740-2960 mg/L for the degradation of 4 mg/L DCF
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in the presence of 25 mg/L CB; this implies that CB is not a suitable agent to activate H2O2.
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3.3 Effect of DCF, SPS and CB concentrations on DCF degradation
The impact of increasing DCF concentration in the range 0.5-4 mg/L on its degradation is
rate=-
𝑑[SMX] dt
=k app [SMX]
(5)
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reaction rate can be expressed as follows [35]:
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presented in Figure 4A. Considering that DCF degradation follows pseudo-first-order kinetics, the
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Apparent rate constants (kapp) were calculated from the linearized form of Equation (5) and are shown in brackets. As expected, increasing DCF concentration results in longer time periods for its
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complete degradation. This is clearly reflected to the apparent rate constants, which decrease from 0.52 to 0.015 min-1 increasing DCF concentration from 0.5 to 4 mg/L.
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Figure 4B shows the effect of varying CB concentration between 0-75 mg/L. It is observed that increasing CB concentration from 25 to 50 mg/L halved the time needed for the complete degradation of 1 mg/L DCF, as more active sites were available for SPS activation; kapp increased from 0.22 min1
at 25 mg/L CB to 0.667 min-1 at 50 mg/L. However, further increase of CB concentration to 75
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mg/L had practically no impact on kapp, implying that the reaction rate reached a plateau. The existence of such a plateau is common and is associated with the number of the catalyst’s active centers available relative to SPS concentration [25]. Experiments were also performed varying SPS concentration (0-200 mg/L) in the presence of 25 mg/L CB. As seen in Figure 4C, an increase in SPS concentration from 25 to 50 mg/L led to an increase in DCF removal from 30% to 90% after 5 min. Further increase of SPS to 100 and 200 mg/L had an adverse effect on degradation rate with the apparent rate constant decreasing from 0.341 min-
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at 50 mg/L to 0.201 min-1 at 100 mg/L. This may be attributed to the fact that excess amount of SPS
scavenges the highly reactive species forming species of lower reactivity [26,36].
3.4 The water matrix effect All the experiments so far were carried out in UPW, thus ignoring possible effects associated with inorganic and organic species inherently present in real aqueous matrices. In order to investigate such
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interactions, experiments were performed in BW, SW and WW and the results are shown in Figure 5. As clearly seen, DCF degradation rate decreased as the complexity of water matrix increased. Specifically, kapp in WW was approximately 7 times lower than in UPW, showing that degradation
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was significantly hindered in secondary treated wastewater.
In an attempt to understand the impact of main water constituents on degradation, experiments were performed in UPW spiked with various concentrations of bicarbonate, chloride or humic acid
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and the results are shown in Figure 6. Interestingly, the addition of 50 mg/L 𝐻𝐶𝑂3− enhanced DCF degradation and kapp increased from 0.201 to 0.308 min-1 (Figure 6A). This positive effect became
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more pronounced for the higher concentration of 𝐻𝐶𝑂3− tested (500 mg/L) as 1 mg/L DCF was
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completely degraded after only 5 min. In most cases, the presence of 𝐻𝐶𝑂3− in similar oxidation
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systems has a negative effect on degradation, since 𝐻𝐶𝑂3− may scavenge the highly active 𝑆𝑂4•− and/or •OH producing the less reactive carbonate radicals (𝐻𝐶𝑂3•) [37]. However, beneficial effects
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have also been reported, which are attributed mainly to the enhanced formation of •OH induced by 𝐻𝐶𝑂3− [7,38,39]. Additionally, 𝐻𝐶𝑂3• is a one-electron oxidant that acts by both electron transfer and hydrogen abstraction mechanisms to produce radicals from the oxidized targets which may partly
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offset the scavenging of hydroxyl radicals [40,41]. On the other hand, the addition of Cl- (Figure 6B) retarded slightly DCF degradation with the
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effect being more evident for the higher concentration tested (500 mg/L). In accordance with 𝐻𝐶𝑂3− , [42], SO•-4 may react with chloride to form other less reactive radicals such as chlorine and dichloride radicals ( Cl• , Cl•-2 ) [43], thus lowering degradation rates. However, it should be noted that positive
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outcomes due to the addition of Cl- cannot be excluded [21,44]. Humic acid was used to evaluate the effect of natural organic matter typically found in real
matrices. As seen in Figure 6C, the addition of 10-20 mg/L HA resulted in a decrease in kapp from 0.201 to 0.062-0.036 min-1. This phenomenon may be ascribed to the competitive (i) adsorption of HA on the active sites of CB [45], and (ii) consumption of reactive species by target and non-target organics.
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In this light, an additional set of experiments was performed adding an excess (i.e. 2 g/L) of methanol, ethanol or potassium iodide as possible quenchers of •OH and 𝑆𝑂4•− radicals. Methanol and ethanol react with •OH and 𝑆𝑂4•− radicals with comparable rate constants of 1.2-2.8×109 M−1s−1 and 1.6-7.7×107 M−1s−1, respectively [46,47]. On the other hand, potassium iodide is known to be a surface-bound free radicals scavenger [48]. As seen in Figure 6D, ethanol has a seriously detrimental effect on DCF degradation with the rate constant decreasing from 0.201 to 0.039 min-1, while the effect of methanol is far less important with the rate constant becoming 0.143 min-1. Addition of KI
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practically stopped DCF degradation confirming the existence of surface bound radicals. Therefore, it is concluded that reactive radicals (𝑆𝑂4•− , •OH) seem to play a significant role in the present
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degradation system
3.5 Reusability of CB
Reusability tests were carried out to assess the activity of CB upon repeated use, as follows: the
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reaction vessel was added with 120 mL of ultrapure water containing 1 mg/L DCF, 50 mg/L CB and
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50 mg/L SPS. DCF concentration was monitored at regular time intervals for 20 min (first cycle).
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Fresh DCF solution was then added to bring the initial concentration to 1 mg/L and the liquid volume to 120 mL and this procedure was repeated for five cycles. Figure 7 shows the 20-min DCF removal
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for each cycle. It is observed that fresh CB (first cycle) decomposed completely 1 mg/L DCF in 20 min, while DCF removal remained practically unchanged after 3 cycles. Even after 5 cycles DCF
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removal remained at high levels (~80%) showing that CB retains its ability to activate persulfate upon repeated use.
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Summarizing, DCF degradation (see Figure 2A) takes place only in the presence of CB, proving its efficiency towards activating SPS. Moreover, CB maintains its activity (see Figure 7) upon repeated use showing the heterogeneous nature of the present system. As stated before, most studies
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dealing with carbon-based, activated persulfate systems propose a non-radical pathway, involving interactions between the reactant, carbon catalysts and persulfate, as follows: 2
S2O8 [DCF] 2SO2-4 +CB
(6)
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CB DCF CB(e )
In this work though, the inhibitory effect associated with the use of common scavengers (see Figure 6D) shows that the main mechanism involves the formation of radicals and can be described as follows: CB 2S2O82 S2O8 SO42 + SO 4
(7)
2 SO 4 H 2 O SO4 OH H
(8)
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This pathway has significant similarities with heterogeneous activation systems with metals [49]. No matter whether the mechanism is radical or non-radical, CB serves as electron mediator transferring electrons from and/or to SPS.
3.6 Effect of pH on DCF degradation In an attempt to evaluate the effect of solution pH on the process, experiments were carried out
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altering the initial inherent value of 5.5 with the addition of either H2SO4 or NaOH. All experiments were performed in UPW. Having in mind that the adsorption capacity of carbonaceous materials is highly pH-dependent [50], both adsorption (i.e. without SPS) and oxidation experiments (i.e. with
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SPS) were performed; Figure 8 shows the extent of DCF removal after 15 and 30 min at various pH values. As clearly seen, DCF removal by adsorption is favored at highly acidic conditions and this can be explained by the electrostatic forces developed between DCF and the surface of CB. The pKa
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of DCF is equal to 4.15 (at 25oC) [51], i.e. DCF is positively charged at pH< pKa, whereas negatively charged at at pH>pKa. Since the surface of CB is negatively charged at pH>2 (Figure 1B), enhanced
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adsorption is expected to occur at pH=3 owing to the electrostatic attraction between CB and DCF,
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while electrostatic repulsion leads to lower adsorption at near-neutral and basic pH. Interestingly, the
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extent of oxidation seems not to be affected by the initial solution pH, i.e. complete DCF degradation occurred after 30 min irrespective of the pH. These findings also imply that the discrepancy in reactivities of the various actual matrices shown in Figure 5 is likely to be due to their composition
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rather than their different pH values.
It should be noticed here that for the experiments carried out at initial pH values of 5.5 and 3, the
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final pH values remained almost the same (i.e. 5 and 2.9, respectively). Only for the experiment at
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pH 9.5, did the final value drop by about two units.
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3.7 Coupling CB with solar irradiation or ultrasound for SPS activation In order to explore the possible existence of synergistic effects on SPS activation associated with
the simultaneous use of CB and (i) simulated solar or (ii) ultrasound (US) radiation, additional experiments were carried out and the results are presented in Figure 9. As seen in Figure 9A, solar radiation in the absence of CB and SPS resulted in about 30% DCF removal after 30 min of reaction, presumably due to its photolytic degradation. The addition of 100 mg/L SPS had practically no impact on degradation, implying that SPS activation by simulated solar radiation was inconsiderable. On the
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contrary, the simultaneous use of solar radiation, SPS and CB resulted in complete removal of DCF in 7 min, i.e. half of the time needed for the combined CB and SPS system in the dark. Figure 9A also shows (dashed line) the theoretical DCF profile that would be derived if the individual processes (i.e. solar radiation and CB/SPS) were simply added up; this indeed is slower than the actual profile associated with the combined process (i.e. CB/SPS/solar), which implies synergy. To quantify the extent of synergy (S), Equation 9 was applied: n
1
(9)
kcombined
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S (%) 100
kcombined ki
where kcombined is the rate constant of the combined process (CB/SPS/solar) and ki is the rate constant
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of each individual process (CB/SPS and solar). The simultaneous activation of SPS by CB and solar radiation resulted in 50% synergy.
A similar procedure was then repeated replacing solar radiation with ultrasound and the results are
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shown in Figure 9B. US alone resulted in ca 70% DCF degradation after 60 min of reaction and this increased to 100% when it was coupled with SPS, thus implying SPS activation by ultrasound. The
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combined process (i.e. CB/SPS/US) was again faster than the sum of the individual processes (i.e.
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US and CB/SPS) with the level of synergy being 45%.
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These results are in line with the work of Kemmou et al. [25], who studied SPS activation by biochar for the degradation of sulfamethoxazole; coupling biochar with solar radiation or ultrasound resulted in 23% and 35% synergy, respectively. Metheniti el al. [43], who examined the degradation
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of propyl paraben with SPS activated by a heterogeneous catalyst (i.e. iron in carbon xerogel) combined with ultrasound, reported synergy between 35% and 50% depending on the water matrix;
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4. Conclusions
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this was attributed to the enhancement of mass transfer and the alteration of the catalyst surface.
Carbon black, a low-cost carbonaceous material, was used as a heterogeneous activator of persulfate for the degradation of diclofenac. The main conclusions can be summarized as follows: 1)
CB exhibits satisfactory SPS activation capacity leading to complete degradation of DCF in
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short periods of time. Metal-free carbocatalysts are relatively new materials in persulfate oxidation systems and their use can pave the way for substituting costly activators. 2)
DCF degradation follows pseudo-first-order kinetics, which are determined by CB, SPS and
DCF concentrations, a behavior typical to the majority of persulfate-assisted oxidation systems. 3)
Complex interactions between CB, SPS and the various inorganic and organic species present
in real aqueous matrices retard DCF degradation, especially in the case of wastewater, thus requiring
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longer treatment times. However, the effects of individual species may be either positive (i.e. bicarbonate) or detrimental (i.e. humic acid), thus making the final, net effect case-specific. 4)
The CB/SPS system can successfully be combined with other activators, such as solar and
ultrasound radiation, thus resulting in synergistic phenomena and boosting degradation efficiency.
Acknowledgements
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AP and DM acknowledge support of this work by the project “INVALOR: Research Infrastructure for Waste Valorization and Sustainable Management” (MIS 5002495) which is implemented under the Action “Reinforcement of the Research and Innovation Infrastructure”, funded by the Operational
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Programme "Competitiveness, Entrepreneurship and Innovation" (NSRF 2014-2020) and co-
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financed by Greece and the European Union (European Regional Development Fund).
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Environ. Chem. 11 (2014) 51-62.
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Figure 1. (A) X-Ray diffraction pattern of CB; (B) Zeta potential of CB as a function of pH; (C)
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Figure 2. (A) Removal of DCF by SPS alone, CB alone and SPS/CB; (B) SPS conversion during the
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SPS/CB experiment. Conditions: 1 mg/L DCF, 100 mg/L SPS, 25 mg/L CB in UPW.
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Figure 3. Comparison between SPS and H2O2 for the degradation of 4 mg/L DCF with 25 mg/L CB
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Figure 4. Effect of (A) initial DCF concentration on its degradation with 25 mg/L CB and 100 mg/L SPS in UPW; (B) CB concentration on 1 mg/L DCF degradation with 100 mg/L SPS in UPW; (C) SPS concentration on 1 mg/L DCF degradation with 25 mg/L CB in UPW. Numbers in brackets show
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apparent rate constants in min-1. ND: not determined.
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1.0 UPW (0.201) BW (0.157) SW (0.111) WW (0.030)
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Figure 5. Effect of water matrix on 1 mg/L DCF degradation with 25 mg/L CB and 100 mg/L SPS.
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Numbers in brackets show apparent rate constants in min-1.
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Figure 6.
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Figure 6. Effect of (A) bicarbonate; (B) chloride; (C) humic acid; (D) MeOH, EtOH and KI on 1 mg/L DCF degradation with 25 mg/L CB and 100 mg/L SPS. Numbers in brackets show apparent
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rate constants in min-1.
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DCF 20min-removal (%)
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Experimental runs
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Figure 7.
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Figure 7. Removal of DCF (1 mg/L) after 20 min of reaction for five consecutive runs with 50 mg/L
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Figure 8.
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Figure 8. Effect of initial solution pH on 15-min DCF removal (left axis) and 30-min DCF removal
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(right axis) in the absence (patterned columns) and in the presence (non-patterned columns) of 100
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solar (0.016) solar/SPS (0.021) CB/SPS (0.201) sum CB/SPS/solar (0.433)
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US (0.024) US/SPS (0.067) CB/SPS (0.201) sum CB/SPS/US (0.409)
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Figure 9.
Figure 9. Synergistic effect of (A) solar radiation; (B) ultrasound radiation on the degradation of 1
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S0920-5861(19)30058-6 https://doi.org/10.1016/j.cattod.2019.02.025 CATTOD 11977
To appear in:
Catalysis Today
Received date: Revised date: Accepted date:
28 November 2018 1 February 2019 12 February 2019
Please cite this article as: Dimitriadou S, Frontistis Z, Petala A, Bampos G, Mantzavinos D, Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices, Catalysis Today (2019), https://doi.org/10.1016/j.cattod.2019.02.025 This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.
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Carbocatalytic activation of persulfate for the removal of drug diclofenac from aqueous matrices
Sofia Dimitriadou1, Zacharias Frontistis2, Athanasia Petala1,3*, Georgios Bampos1, Dionissios Mantzavinos1,3
Department of Chemical Engineering, University of Patras, Caratheodory 1, University Campus,
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1
GR-26504 Patras, Greece
Department of Environmental Engineering, University of Western Macedonia, GR-50100 Kozani,
Greece 3
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2
INVALOR: Research Infrastructure for Waste Valorization and Sustainable Management,
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Caratheodory 1, University Campus, GR-26504 Patras, Greece
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Corresponding author: E-mail: [email protected]
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Graphical abstract
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Highlights
Carbon black shows good activity as a low-cost persulfate activator for diclofenac oxidation
Complex water matrix interactions determine degradation rates
Diclofenac adsorption on carbon black surface is pH-sensitive, but oxidation is not
Coupling carbon black with light or ultrasound enhances degradation synergistically
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Αbstract
In this study, carbon black (CB) was employed as a heterogeneous activator for the conversion of
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sodium persulfate (SPS) to reactive species for the degradation of drug diclofenac (DCF).
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Experiments were performed at DCF concentrations between 0.5 and 4 mg/L, CB concentrations between 25 and 75 mg/L and SPS concentrations between 25 and 200 mg/L. Degradation rates, based
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on pseudo-first order kinetics, generally increased with decreasing DCF and increasing CB concentrations. The rate also increased with increasing SPS concentration up to 50 mg/L and decreased at higher values due to scavenging effects. Besides experiments in ultrapure water (UPW),
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real matrices were tested (i.e. bottled water (BW), surface water (SW), secondary treated wastewater (WW)), as well as UPW spiked with bicarbonate (50-500 mg/L), chloride (100-500 mg/L) or humic
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acid (10-20 mg/L). Degradation rates decreased with increasing matrix complexity, while the addition of chloride or humic acid was detrimental to the process; on the contrary, bicarbonate at 500 mg/L enhanced DCF degradation rate nearly five-fold. The effect of initial solution pH was also studied in
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the range 3-9.5 showing that degradation was not pH-sensitive. Experiments were also performed activating SPS by simulated solar radiation or 20 kHz ultrasound with or without CB. Coupling activators (i.e. CB with solar light or CB with ultrasound) favored DCF degradation in a synergistic way, with the level of synergy being 45-50%. Keywords: carbocatalysis; kinetics; solar light; synergy; ultrasound; water matrix Introduction
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Pharmaceuticals and personal care products constitute a large group of synthetic organic chemicals inextricably linked to everyday life that have been categorized as contaminants of global concern [1]. These concerns derive from the fact that trace levels of the above-mentioned compounds have been quantified in surface and wastewaters, pointing out the inefficiency of classical physicochemical and biological treatment methods used in conventional wastewater treatment plants [2]. Diclofenac (DCF) is a nonsteroidal anti-inflammatory drug contained in medications under a variety of trade names that has been detected in water bodies [3]. In particular, European Union included DCF in the list of
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priority substances (2013/39/EU) and set the acceptable environmental limits to 100 ng/L for inland and 10 ng/L for coastal waters [4]. Adverse effects to aqueous ecosystems have already been reported due to prolonged exposure to DCF [5], thus creating major concerns for the impact on human health.
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To address this problem, the development of effective methods, such as the advanced oxidation processes (AOPs), for DCF and other pharmaceuticals removal has gained significant interest. AOPs are a family of similar but not identical technologies relying on the formation of various oxidative
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species and, especially, hydroxyl radicals (•OH), which are characterized by an oxidation potential of 2.8 V vs. NHE at pH ~ 6 [6].
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Recently, the activity of sulfate radicals (𝑆𝑂4•− ), which are characterized by an equally high
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oxidation potential (2.6 V vs NHE at pH~6) has been studied in detail; sulfate radicals often lead to
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a more efficient decomposition mechanism [7,8], while some of their advantageous characteristics include longer half-life [9] and reactivity over a wide pH range [10].
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Sulfate radicals can be produced after activation of compounds belonging to the group of persulfates. Generally, persulfates are of comparable cost to hydrogen peroxide in the order of 1 $/kg but they are easier to handle due to their crystalline form, thus limiting problems related to storage
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and transportation [11]. Common activation methods include heat [12] and ultraviolet (UV) radiation [13], that can provide energy for the fission of O-O bond in the persulfate structure and the formation
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of two sulfate radicals according to Equation 1 [14]. In addition, persulfates can be activated via redox reactions with transition metals in both heterogeneous and homogeneous systems leading to the formation of one sulfate radical (Equation 2) [15]. In both cases, 𝑆𝑂4•− may further react resulting in the formation of additional oxidizing species such as •OH (Equation 3). (1)
S2O82- +Mn+ SO42- +SO-•4 +M(n+1)+
(2)
SO-•4 +OH- SO42- + • OH
(3)
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hv/heat S2O82- 2SO-•4
Αctivation by transition metals reduces significantly the energy requirements of the process and consequently the whole cost, making this activation method more attractive for practical applications [14]. Moreover, special emphasis has been given towards the development of heterogeneous,
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transition metal-based activation systems, thus avoiding secondary pollution due to the occurrence of metal ions in aqueous solutions [16]. In particular, copper- and cobalt-based catalysts have been used for heterogeneous activation resulting in successful degradation of many persistent micro-pollutants [17,18]. However, the reported activation mechanisms include irreversible changes in the oxidation state of the metal, thus challenging the true heterogeneous catalytic nature of the process and limiting the privileges of collection and reuse of the catalyst [19]. In order to avoid such phenomena, metal-free carbonaceous materials have been proposed as
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alternative persulfate activators. Activated carbon [20], carbon nanotube [21], reduced graphene oxide [22] and graphene [23], multi-walled carbon nanotubes [24] and biochars [25,26] are some characteristic examples. Considering carbonaceous-based activation systems, several studies propose
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a non-radical oxidation pathway for organic degradation, proceeding through the formation of reactive complexes between activators and the persulfate and subsequent electron abstraction by these complexes [24,27,28]. This mechanism can be described according to the following equation (Εqn 2
S2O8 [OP] 2SO2-4 +CC
(4)
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CC OP CC(e )
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4), where CC is the carbocatalyst and OP the organic pollutant.
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Carbon black (CB), mainly consisting of fine particles of carbon, has recently been proposed as a low-cost alternative to expensive carbon materials, such as graphene, in the field of electrochemistry
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showing great performance as an electric conductive agent [29]. In particular, the current worldwide CB production exceeds 8 Mt/y, with 90% used in rubber applications. This has a positive impact on
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its commercial price, being several orders of magnitude lower than other carbonaceous materials. A complicated blend of carbon chemistry and surface energy makes it ideal for a variety of applications
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explaining its ever-increasing popularity and huge utility. In the present work, the efficiency of CB for SPS activation and the subsequent degradation of DCF is tested. To the best of our knowledge,
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only o handful of studies report DCF removal from water matrices with heterogeneous persulfate activation systems, such as metal-doped, manganese oxide octahedral molecular sieve catalysts [30], perovskite oxides [31] and graphene oxide-TiO2 nanosheets [32], whereas this is the first work using CB instead of high-priced carbonaceous materials for SPS activation. The effects of CB, SPS and
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DCF concentrations on DCF degradation kinetics are investigated. Emphasis is given on the effect of water matrix, while possible synergistic effects arising from the simultaneous use of solar or ultrasound irradiation are also discussed.
2. Experimental 2.1. Chemicals and water matrices
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Carbon black (Vulcan XC72R) was supplied by Cabot. Diclofenac sodium salt (C14H10Cl12NNaO2, CAS:15307-79-6) and sodium persulfate (Na2S2O8, CAS: 7775-27-1) were purchased from Sigma Aldrich. Potassium iodide (KI, CAS:7681-11-0) for SPS consumption measurements and for radical quenching experiments was purchased from Carlo Erba reagents. Hydrogen peroxide (H2O2, CAS: 7722-84-1), sodium chloride (NaCl, CAS:7647-14-5), sodium bicarbonate (NaHCO3, CAS: 144-558), humic acid (HA, CAS:1415-93-6), methanol (CH3OH, CAS: 67-56-1), ethanol (CH3CH2OH, CAS: 64-17-5) sodium hydroxide (NaOH, CAS: 1310-73-2), sulfuric acid (H2SO4, CAS: 7664-93-9)
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and acetonitrile (CH3CN, CAS: 75-05-8, for HPLC analysis) were also purchased from Sigma Aldrich.
Ultrapure water (pH=6, EASYpureRF-Barnstead/Thermolyne, USA) was employed in most of the
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experiments presented in this study. Moreover, bottled water (BW) (pH=7.8, 0.4 mS/cm conductivity, 182 mg/L HCO3- , 2 mg/L Cl- , 14 mg/L SO24 , 8.6 mg/L NO3- and 76 mg/L of various metal ions), surface water (SW) taken from a river in the region of Athens, Greece (pH=7.5, 491 μS/cm
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conductivity, 2.7 mg/L total organic carbon, 274 mg/L SO24 , 5 mg/L Cl- ) and secondary treated
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wastewater (WW) (pH=8, 21 mg/L chemical oxygen demand) were also used in order to study the
2.2 Catalyst characterization
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effect of water matrix.
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CB was characterized with respect to its specific surface area according to the Brunaueur-EmmettTeller (BET) method with the use of a Micromeritics (Gemini III 2375), whereas Brucker D8 Advance (Cu Kα) was used for X-Ray diffraction pattern of the sample. The zeta potential
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measurements of CB were performed via laser doppler micro-electrophoresis employing a Malvern Zetasizer (Nano-ZS). Briefly, small amounts of CB powder were ultrasonically dispersed in UPW.
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The pH of these suspensions was adjusted using H2SO4 or NaOH solutions. The velocity of the CB particles of each suspension was measured implementing a patented laser interferometric technique (phase analysis light scattering), enabling the calculation of electrophoretic mobility of the particles
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and, thus, the estimation of the zeta potential values. Fourier transform infrared (FTIR) spectroscopy was performed using a Perkin Elmer Spectrum RX FTIR system. The measurement range was 4000– 450 cm-1. Vulcan (about 0.5%) and KBr were sieved and pressurized to produce a homogeneous disk.
2.3 Experimental procedure
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Degradation experiments were carried out in a pyrex vessel of 250 mL capacity open to the atmosphere. In a typical run, 120 mL of an aqueous solution with the desired concentration of DCF (1 mg/L in most cases) was loaded in the vessel followed by the addition of CB and SPS. The solution was under continuous magnetic stirring. Samples of 1.2 mL were withdrawn from the vessel at desired time periods followed by the addition of 0.3 mL of methanol in order to quench the reaction. Samples were then filtered (0.22 μm PVDF filters) to remove CB particles and DCF concentration
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was measured with the use of high performance liquid chromatography (HPLC).
2.4 Analytical methods
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An HPLC system (Waters Alliance 2695) consisting of a photodiode array detector (Waters 2996), a gradient pump (Waters 2695) and a Kinetex column (C18 100A, 150 mm × 3 mm, 2.6 μm particle size) maintained at 45oC was employed. The mobile phase consisted of 68:32 water:acetonitrile,
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while the injection volume was 40 μL.
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Persulfate was measured spectrophotometrically based on its reaction with potassium iodide [33]. A solar simulator (Oriel, model LCS-100) equipped with a 100 W Xenon (ozone free) lamp and a
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Branson 450 horn-type digital sonifier operating at a frequency of 20 kHz and an actual power density
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of 36 W/L were employed to study the combined activation of persulfate by CB/solar light or
3. Results
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3.1 Characterization of CB
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CB/ultrasound. Details regarding the solar simulator and the sonifier can be found elsewhere [25].
The XRD pattern of CB is presented in Figure 1A. It consists of reflections at 24o and 43.5o attributed to graphitic regions (JCPDS 2-456) [34]. Measurements of zeta potential (Figure 1B) show
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that the CB surface is negatively charged at pH>2 with the negative charge increasing at more alkaline conditions. The specific surface area of the sample was measured equal to 216 m2/g. The FTIR spectrum of CB is shown in Figure 1C. The peak at 3450 cm–1 is assigned to stretching vibration of
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the –OH groups of CB. As can be seen, the surface of CB exhibits a significant number of –OH groups. The absence of a peak at about 1100 cm–1 denotes that there are no C-O-C bonds. The peaks at about 1640 and 1380 cm–1 correspond to C=C and C-OH groups.
3.2 Screening experiments
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Figure 2A shows normalized DCF concentration-time profiles for experiments where CB and SPS were employed either alone or together. CB (25 mg/L) alone led to ca 20% decrease in DCF concentration after 30 min presumably due to adsorption phenomena. A comparable concentration decrease was observed when SPS (100 mg/L) was used alone due its mild oxidizing ability. On the contrary, the combined application of CB and SPS led to complete degradation of 1 mg/L DCF after 15 min, indicating the successful activation of SPS by CB. In addition, Figure 2B shows the conversion of 100 mg/L SPS in the presence of 25 mg/L CB and 1 mg/L DCF. It is observed that ca
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30% of SPS was consumed within the first 5 min of reaction, while its concentration remained almost constant thereafter.
The efficiency of SPS for DCF degradation was compared to that of H2O2. As seen in Figure 3,
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SPS at 100 mg/L was more effective than H2O2 at 740-2960 mg/L for the degradation of 4 mg/L DCF
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in the presence of 25 mg/L CB; this implies that CB is not a suitable agent to activate H2O2.
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3.3 Effect of DCF, SPS and CB concentrations on DCF degradation
The impact of increasing DCF concentration in the range 0.5-4 mg/L on its degradation is
rate=-
𝑑[SMX] dt
=k app [SMX]
(5)
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reaction rate can be expressed as follows [35]:
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presented in Figure 4A. Considering that DCF degradation follows pseudo-first-order kinetics, the
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Apparent rate constants (kapp) were calculated from the linearized form of Equation (5) and are shown in brackets. As expected, increasing DCF concentration results in longer time periods for its
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complete degradation. This is clearly reflected to the apparent rate constants, which decrease from 0.52 to 0.015 min-1 increasing DCF concentration from 0.5 to 4 mg/L.
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Figure 4B shows the effect of varying CB concentration between 0-75 mg/L. It is observed that increasing CB concentration from 25 to 50 mg/L halved the time needed for the complete degradation of 1 mg/L DCF, as more active sites were available for SPS activation; kapp increased from 0.22 min1
at 25 mg/L CB to 0.667 min-1 at 50 mg/L. However, further increase of CB concentration to 75
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mg/L had practically no impact on kapp, implying that the reaction rate reached a plateau. The existence of such a plateau is common and is associated with the number of the catalyst’s active centers available relative to SPS concentration [25]. Experiments were also performed varying SPS concentration (0-200 mg/L) in the presence of 25 mg/L CB. As seen in Figure 4C, an increase in SPS concentration from 25 to 50 mg/L led to an increase in DCF removal from 30% to 90% after 5 min. Further increase of SPS to 100 and 200 mg/L had an adverse effect on degradation rate with the apparent rate constant decreasing from 0.341 min-
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at 50 mg/L to 0.201 min-1 at 100 mg/L. This may be attributed to the fact that excess amount of SPS
scavenges the highly reactive species forming species of lower reactivity [26,36].
3.4 The water matrix effect All the experiments so far were carried out in UPW, thus ignoring possible effects associated with inorganic and organic species inherently present in real aqueous matrices. In order to investigate such
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interactions, experiments were performed in BW, SW and WW and the results are shown in Figure 5. As clearly seen, DCF degradation rate decreased as the complexity of water matrix increased. Specifically, kapp in WW was approximately 7 times lower than in UPW, showing that degradation
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was significantly hindered in secondary treated wastewater.
In an attempt to understand the impact of main water constituents on degradation, experiments were performed in UPW spiked with various concentrations of bicarbonate, chloride or humic acid
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and the results are shown in Figure 6. Interestingly, the addition of 50 mg/L 𝐻𝐶𝑂3− enhanced DCF degradation and kapp increased from 0.201 to 0.308 min-1 (Figure 6A). This positive effect became
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more pronounced for the higher concentration of 𝐻𝐶𝑂3− tested (500 mg/L) as 1 mg/L DCF was
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completely degraded after only 5 min. In most cases, the presence of 𝐻𝐶𝑂3− in similar oxidation
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systems has a negative effect on degradation, since 𝐻𝐶𝑂3− may scavenge the highly active 𝑆𝑂4•− and/or •OH producing the less reactive carbonate radicals (𝐻𝐶𝑂3•) [37]. However, beneficial effects
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have also been reported, which are attributed mainly to the enhanced formation of •OH induced by 𝐻𝐶𝑂3− [7,38,39]. Additionally, 𝐻𝐶𝑂3• is a one-electron oxidant that acts by both electron transfer and hydrogen abstraction mechanisms to produce radicals from the oxidized targets which may partly
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offset the scavenging of hydroxyl radicals [40,41]. On the other hand, the addition of Cl- (Figure 6B) retarded slightly DCF degradation with the
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effect being more evident for the higher concentration tested (500 mg/L). In accordance with 𝐻𝐶𝑂3− , [42], SO•-4 may react with chloride to form other less reactive radicals such as chlorine and dichloride radicals ( Cl• , Cl•-2 ) [43], thus lowering degradation rates. However, it should be noted that positive
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outcomes due to the addition of Cl- cannot be excluded [21,44]. Humic acid was used to evaluate the effect of natural organic matter typically found in real
matrices. As seen in Figure 6C, the addition of 10-20 mg/L HA resulted in a decrease in kapp from 0.201 to 0.062-0.036 min-1. This phenomenon may be ascribed to the competitive (i) adsorption of HA on the active sites of CB [45], and (ii) consumption of reactive species by target and non-target organics.
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In this light, an additional set of experiments was performed adding an excess (i.e. 2 g/L) of methanol, ethanol or potassium iodide as possible quenchers of •OH and 𝑆𝑂4•− radicals. Methanol and ethanol react with •OH and 𝑆𝑂4•− radicals with comparable rate constants of 1.2-2.8×109 M−1s−1 and 1.6-7.7×107 M−1s−1, respectively [46,47]. On the other hand, potassium iodide is known to be a surface-bound free radicals scavenger [48]. As seen in Figure 6D, ethanol has a seriously detrimental effect on DCF degradation with the rate constant decreasing from 0.201 to 0.039 min-1, while the effect of methanol is far less important with the rate constant becoming 0.143 min-1. Addition of KI
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practically stopped DCF degradation confirming the existence of surface bound radicals. Therefore, it is concluded that reactive radicals (𝑆𝑂4•− , •OH) seem to play a significant role in the present
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degradation system
3.5 Reusability of CB
Reusability tests were carried out to assess the activity of CB upon repeated use, as follows: the
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reaction vessel was added with 120 mL of ultrapure water containing 1 mg/L DCF, 50 mg/L CB and
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50 mg/L SPS. DCF concentration was monitored at regular time intervals for 20 min (first cycle).
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Fresh DCF solution was then added to bring the initial concentration to 1 mg/L and the liquid volume to 120 mL and this procedure was repeated for five cycles. Figure 7 shows the 20-min DCF removal
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for each cycle. It is observed that fresh CB (first cycle) decomposed completely 1 mg/L DCF in 20 min, while DCF removal remained practically unchanged after 3 cycles. Even after 5 cycles DCF
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removal remained at high levels (~80%) showing that CB retains its ability to activate persulfate upon repeated use.
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Summarizing, DCF degradation (see Figure 2A) takes place only in the presence of CB, proving its efficiency towards activating SPS. Moreover, CB maintains its activity (see Figure 7) upon repeated use showing the heterogeneous nature of the present system. As stated before, most studies
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dealing with carbon-based, activated persulfate systems propose a non-radical pathway, involving interactions between the reactant, carbon catalysts and persulfate, as follows: 2
S2O8 [DCF] 2SO2-4 +CB
(6)
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CB DCF CB(e )
In this work though, the inhibitory effect associated with the use of common scavengers (see Figure 6D) shows that the main mechanism involves the formation of radicals and can be described as follows: CB 2S2O82 S2O8 SO42 + SO 4
(7)
2 SO 4 H 2 O SO4 OH H
(8)
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This pathway has significant similarities with heterogeneous activation systems with metals [49]. No matter whether the mechanism is radical or non-radical, CB serves as electron mediator transferring electrons from and/or to SPS.
3.6 Effect of pH on DCF degradation In an attempt to evaluate the effect of solution pH on the process, experiments were carried out
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altering the initial inherent value of 5.5 with the addition of either H2SO4 or NaOH. All experiments were performed in UPW. Having in mind that the adsorption capacity of carbonaceous materials is highly pH-dependent [50], both adsorption (i.e. without SPS) and oxidation experiments (i.e. with
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SPS) were performed; Figure 8 shows the extent of DCF removal after 15 and 30 min at various pH values. As clearly seen, DCF removal by adsorption is favored at highly acidic conditions and this can be explained by the electrostatic forces developed between DCF and the surface of CB. The pKa
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of DCF is equal to 4.15 (at 25oC) [51], i.e. DCF is positively charged at pH< pKa, whereas negatively charged at at pH>pKa. Since the surface of CB is negatively charged at pH>2 (Figure 1B), enhanced
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adsorption is expected to occur at pH=3 owing to the electrostatic attraction between CB and DCF,
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while electrostatic repulsion leads to lower adsorption at near-neutral and basic pH. Interestingly, the
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extent of oxidation seems not to be affected by the initial solution pH, i.e. complete DCF degradation occurred after 30 min irrespective of the pH. These findings also imply that the discrepancy in reactivities of the various actual matrices shown in Figure 5 is likely to be due to their composition
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rather than their different pH values.
It should be noticed here that for the experiments carried out at initial pH values of 5.5 and 3, the
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final pH values remained almost the same (i.e. 5 and 2.9, respectively). Only for the experiment at
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pH 9.5, did the final value drop by about two units.
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3.7 Coupling CB with solar irradiation or ultrasound for SPS activation In order to explore the possible existence of synergistic effects on SPS activation associated with
the simultaneous use of CB and (i) simulated solar or (ii) ultrasound (US) radiation, additional experiments were carried out and the results are presented in Figure 9. As seen in Figure 9A, solar radiation in the absence of CB and SPS resulted in about 30% DCF removal after 30 min of reaction, presumably due to its photolytic degradation. The addition of 100 mg/L SPS had practically no impact on degradation, implying that SPS activation by simulated solar radiation was inconsiderable. On the
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contrary, the simultaneous use of solar radiation, SPS and CB resulted in complete removal of DCF in 7 min, i.e. half of the time needed for the combined CB and SPS system in the dark. Figure 9A also shows (dashed line) the theoretical DCF profile that would be derived if the individual processes (i.e. solar radiation and CB/SPS) were simply added up; this indeed is slower than the actual profile associated with the combined process (i.e. CB/SPS/solar), which implies synergy. To quantify the extent of synergy (S), Equation 9 was applied: n
1
(9)
kcombined
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S (%) 100
kcombined ki
where kcombined is the rate constant of the combined process (CB/SPS/solar) and ki is the rate constant
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of each individual process (CB/SPS and solar). The simultaneous activation of SPS by CB and solar radiation resulted in 50% synergy.
A similar procedure was then repeated replacing solar radiation with ultrasound and the results are
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shown in Figure 9B. US alone resulted in ca 70% DCF degradation after 60 min of reaction and this increased to 100% when it was coupled with SPS, thus implying SPS activation by ultrasound. The
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combined process (i.e. CB/SPS/US) was again faster than the sum of the individual processes (i.e.
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US and CB/SPS) with the level of synergy being 45%.
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These results are in line with the work of Kemmou et al. [25], who studied SPS activation by biochar for the degradation of sulfamethoxazole; coupling biochar with solar radiation or ultrasound resulted in 23% and 35% synergy, respectively. Metheniti el al. [43], who examined the degradation
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of propyl paraben with SPS activated by a heterogeneous catalyst (i.e. iron in carbon xerogel) combined with ultrasound, reported synergy between 35% and 50% depending on the water matrix;
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4. Conclusions
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this was attributed to the enhancement of mass transfer and the alteration of the catalyst surface.
Carbon black, a low-cost carbonaceous material, was used as a heterogeneous activator of persulfate for the degradation of diclofenac. The main conclusions can be summarized as follows: 1)
CB exhibits satisfactory SPS activation capacity leading to complete degradation of DCF in
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short periods of time. Metal-free carbocatalysts are relatively new materials in persulfate oxidation systems and their use can pave the way for substituting costly activators. 2)
DCF degradation follows pseudo-first-order kinetics, which are determined by CB, SPS and
DCF concentrations, a behavior typical to the majority of persulfate-assisted oxidation systems. 3)
Complex interactions between CB, SPS and the various inorganic and organic species present
in real aqueous matrices retard DCF degradation, especially in the case of wastewater, thus requiring
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longer treatment times. However, the effects of individual species may be either positive (i.e. bicarbonate) or detrimental (i.e. humic acid), thus making the final, net effect case-specific. 4)
The CB/SPS system can successfully be combined with other activators, such as solar and
ultrasound radiation, thus resulting in synergistic phenomena and boosting degradation efficiency.
Acknowledgements
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AP and DM acknowledge support of this work by the project “INVALOR: Research Infrastructure for Waste Valorization and Sustainable Management” (MIS 5002495) which is implemented under the Action “Reinforcement of the Research and Innovation Infrastructure”, funded by the Operational
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Programme "Competitiveness, Entrepreneurship and Innovation" (NSRF 2014-2020) and co-
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financed by Greece and the European Union (European Regional Development Fund).
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Environ. Chem. 11 (2014) 51-62.
18
(A)
Graphite
Intensity (a.u.)
(002)
10
20
30
40
50
60
ο
70
IP T
(101)
80
U A
N
-20
0
2
70
(C)
4
6
pH
8
10
12
14
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%T
60
ED
M
-40
-60
SC R
(B)
0
PT
Zeta potential (mV)
2θ ( )
A
50
40
4000
3500
3000
2500
2000
1500 -1
wavenumber (cm ) Figure 1.
1000
500
19
Figure 1. (A) X-Ray diffraction pattern of CB; (B) Zeta potential of CB as a function of pH; (C)
A
CC E
PT
ED
M
A
N
U
SC R
IP T
FTIR spectrum of CB.
20
(A)
1.0
0.6 0.4
SPS CB CB/SPS
0.2 0.0
0
5
10
15
20
25
IP T
C/C0
0.8
30
(B)
N
0.8
A
0.4
M
C/C0
0.6
0
5
10
ED
0.2 0.0
U
1.0
SC R
Time (min)
15
20
25
30
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PT
Time (min)
Figure 2.
Figure 2. (A) Removal of DCF by SPS alone, CB alone and SPS/CB; (B) SPS conversion during the
A
SPS/CB experiment. Conditions: 1 mg/L DCF, 100 mg/L SPS, 25 mg/L CB in UPW.
21
1.0 0.8
0.4
IP T
C/C0
0.6
100 mg/L CB/SPS 740 mg/L CB/H2O2
0.2
300 mg/L CB/H2O2
0
10
20
30
40
50
60
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0.0
Time (min)
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Figure 3.
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Figure 3. Comparison between SPS and H2O2 for the degradation of 4 mg/L DCF with 25 mg/L CB
A
CC E
PT
ED
M
A
in UPW.
22
1.0
0.5 mg/L 1.0 mg/L 3.0 mg/L 4.0 mg/L
0.8
(A)
(0.520) (0.201) (0.044) (0.015)
C/C0
0.6 0.4
0.0
0
10
20
30
40
50
IP T
0.2
60
(B)
1.0
0.4
A
0.0 mg/L (ND) 25 mg/L (0.201) 50 mg/L (0.677) 75 mg/L (0.605)
M
C/C0
0.6
N
U
0.8
0
5
10
ED
0.2 0.0
15
20
25
30
PT
Time (min)
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1.0
(C)
0.6 0.4
A
C/C0
0.8
0.0 mg/L (ND) 25 mg/L (0.185) 50 mg/L (0.341) 100 mg/L (0.201) 200 mg/L (0.225)
0.2 0.0
0
5
10
Time (min) Figure 4.
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Time (min)
15
23
Figure 4. Effect of (A) initial DCF concentration on its degradation with 25 mg/L CB and 100 mg/L SPS in UPW; (B) CB concentration on 1 mg/L DCF degradation with 100 mg/L SPS in UPW; (C) SPS concentration on 1 mg/L DCF degradation with 25 mg/L CB in UPW. Numbers in brackets show
A
CC E
PT
ED
M
A
N
U
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apparent rate constants in min-1. ND: not determined.
24
1.0 UPW (0.201) BW (0.157) SW (0.111) WW (0.030)
0.6
IP T
C/C0
0.8
0.4
0.0
0
10
20
30
40
50
60
U
Time (min)
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0.2
N
Figure 5.
A
Figure 5. Effect of water matrix on 1 mg/L DCF degradation with 25 mg/L CB and 100 mg/L SPS.
A
CC E
PT
ED
M
Numbers in brackets show apparent rate constants in min-1.
25 (A)
1.0
C/C0
0.8 0.0 mg/L (0.201) 50 mg/L (0.308) 100 mg/L (0.417) 500 mg/L (0.959)
0.6 0.4 0.2 0.0
0
5
10
15
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Time (min) (B)
1.0
0.0 mg/L (0.201) 100 mg/L (0.140) 500 mg/L (0.084)
0.6
C/C0
SC R
0.8
0.4
0
10
20
30
40
N
0.0
U
0.2
1.0
ED
C/C0
0 mg/L (0.201) 10 mg/L (0.062) 20 mg/L (0.036)
0.6
M
(C)
0.8
0.4
0
10
PT
0.2 0.0
A
Time (min)
20
30
40
50
60
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Time (min)
1.0
(D)
0.8
C/C0
A
0.6 0.4
0.0 g/L (0.201) 2.0 g/L MeOH (0.143) 2.0 g/L EtOH (0.039) 2.0 g/L KI (0.016)
0.2 0.0
0
10
20
Time (min)
Figure 6.
30
26
Figure 6. Effect of (A) bicarbonate; (B) chloride; (C) humic acid; (D) MeOH, EtOH and KI on 1 mg/L DCF degradation with 25 mg/L CB and 100 mg/L SPS. Numbers in brackets show apparent
A
CC E
PT
ED
M
A
N
U
SC R
IP T
rate constants in min-1.
100
75
IP T
50
25
th
th
4
5
rd
nd
2
3
st
1
0
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DCF 20min-removal (%)
27
Experimental runs
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Figure 7.
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Figure 7. Removal of DCF (1 mg/L) after 20 min of reaction for five consecutive runs with 50 mg/L
A
CC E
PT
ED
M
A
CB and 50 mg/L SPS.
50
50
25
25
0
0 .5 =9
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pH
pH
=5
.0 =3 pH
pH
=9
.5
.5 =5 pH
=3 pH
IP T
75
SC R
75
.5
100
.0
100
DCF 30min-removal (%)
DCF 15min-removal (%)
28
Figure 8.
N
Figure 8. Effect of initial solution pH on 15-min DCF removal (left axis) and 30-min DCF removal
A
(right axis) in the absence (patterned columns) and in the presence (non-patterned columns) of 100
A
CC E
PT
ED
M
mg/L SPS. Conditions: 1 mg/L DCF, 25 mg/L CB in UPW.
29
(A)
1.0
0.6
0.2 0.0
0
10
20
IP T
solar (0.016) solar/SPS (0.021) CB/SPS (0.201) sum CB/SPS/solar (0.433)
0.4
30
SC R
C/C0
0.8
US (0.024) US/SPS (0.067) CB/SPS (0.201) sum CB/SPS/US (0.409)
A
M
0.6 0.4 0.2
0
10
20
PT
0.0
ED
C/C0
0.8
N
(B)
1.0
U
Time (min)
30
40
50
60
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Time (min)
Figure 9.
Figure 9. Synergistic effect of (A) solar radiation; (B) ultrasound radiation on the degradation of 1
A
mg/L DCF with 25 mg/L CB and 100 mg/L SPS in UPW. Numbers in brackets show apparent rate constants in min-1. Dashed lines show the cumulative concentration-time profile of the respective individual processes.