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persulfate: Contributions of hydroxyl and sulfate radicals
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
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Removal of 2-MIB and geosmin using UV/ persulfate: Contributions of hydroxyl and sulfate radicals Pengchao Xie a, Jun Ma a,*, Wei Liu a, Jing Zou a, Siyang Yue a, Xuchun Li b, Mark R. Wiesner c, Jingyun Fang d a
State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin 150090, China b School of the Environmental Science and Engineering, Zhejiang Gongshang University, Hangzhou 310012, China c Civil and Environmental Engineering, Duke University, Durham 27708-0287, USA d School of Environmental Science and Engineering, Sun Yat-sen University, Guangzhou 510275, China
article info
abstract
Article history:
2-methylisoborneol (2-MIB) and geosmin are two odor-causing compounds that are diffi-
Received 7 July 2014
cult to remove and the cause of many consumer complaints. In this study, we assessed the
Received in revised form
degradation of 2-MIB and geosmin using a UV/persulfate process for the first time. The
27 September 2014
results showed that both 2-MIB and geosmin could be degraded effectively using this
Accepted 17 November 2014
process. The process was modeled based on steady-state assumption with respect to the
Available online 26 November 2014
odor-causing compounds and either hydroxyl or sulfate radicals. The second order rate
constants for 2-MIB and geosmin reacting with the sulfate radical (SO4 ) were estimated to 8
1
1
8
1
and (7.6 ± 0.6) 10 M
1
Keywords:
be (4.2 ± 0.6) 10 M
UV/persulfate
contributions of the hydroxyl radical (OH) to 2-MIB and geosmin degradation were 3.5
Hydroxyl radical
times and 2.0 times higher, respectively, than the contribution from SO4 in Milli-Q water
Sulfate radical
with 2 mM phosphate buffer at pH 7.0. The pseudo-first-order rate constants (kso ) of both 2-
2-MIB
MIB and geosmin increased with increasing dosages of persulfate. Although pH did not
Geosmin
affect the degradation of 2-MIB and geosmin directly, different scavenging effects of
s
s
respectively at a pH of 7.0. The
hydrogen phosphate and dihydrogen phosphate resulted in higher values of kso for both 2MIB and geosmin in acidic condition. Bicarbonate and natural organic matter (NOM) inhibited the degradation of both 2-MIB and geosmin dramatically through consuming OH
and SO4 and were likely to be the main radical scavengers in natural waters when using UV/persulfate process to control 2-MIB and geosmin. © 2014 Elsevier Ltd. All rights reserved.
1.
Introduction
Cyanobacteria or blue-green algae can produce 2methylisoborneol (2-MIB) and geosmin which result in many
complaints of taste and odor in drinking water from consumer (Watson et al., 2008; Srinivasan and Sorial, 2011). The concentration of the two odor-causing compounds can even reach more than 1000 ng/L during algae bloom (Mizuno et al.,
* Corresponding author. Tel.: þ86 451 86282292; fax: þ86 451 86283010. E-mail address: [email protected] (J. Ma). http://dx.doi.org/10.1016/j.watres.2014.11.029 0043-1354/© 2014 Elsevier Ltd. All rights reserved.
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w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
Table 1 e Principle reactions in UV/persulfate process. No.
Reactions
UV
r
H
2 3 4
þ S2 O2 8 þ H /…!SO4 þ HSO4 þ OH SO4 þ OH /SO2 4 SO4 þ H2 O/HSO2 4 þ OH
5
SO4 þ SO4 /S2 O2 8
6 7 8 9 10 In the 11 12 In the 13 14 In the 15 16 In the 17 18 In the 19 20 21 22 23 24 In the 25 26 27 28 29 30 31 In the 32 33 In the 34 35
presence of
presence of
presence of
presence of
presence of
presence of
presence of
presence of
SO4
f ¼ εcb=A
þ
Reaction rate ¼ fI0 f ð1 10A Þ=V P A ¼ bðεc þ εi ci Þ
S2 O2 8 !2SO4
1
2 SO4 þ S2 O2 8 4S2 O8 þ SO4 OH þ S2 O2 8 4S2 O8 þ OH OH þ HSO 4 /SO4 þ H2 O OH þ SO4 /HSO 5 2 þ HSO 4 ⇔SO4 þ H 2-MIB SO4 þ 2 MIB/products OH þ 2 MIB/products geosmin SO4 þ geosmin/products OH þ geosmin/products BA BA þ SO4 /products BA þ OH /products NOM NOM þ SO 4 /products NOM þ OH /products 2 HCO3 /CO3 2 HCO 3 þ SO4 /SO4 þ HCO3 2 2 CO3 þ SO4 /SO4 þ CO3 HCO 3 þ OH /H2 O þ CO3 2 CO3 þ OH /OH þ CO3 H2 CO*3 ⇔Hþ þ HCO 3 2 þ HCO 3 ⇔H þ CO3 2 HPO4 /H2 PO4 SO4 þ H2 PO 4 /products 2 SO4 þ HPO2 4 /SO4 þ HPO4 OH þ H2 PO4 /HPO4 þ H2 O OH þ HPO2 4 /HPO4 þ OH þ H3 PO4 ⇔H þ H2 PO4 þ H2 PO 4 ⇔H þ HPO4 2 3 þ HPO4 ⇔H þ PO4 MeOH SO4 þ MeOH/products OH þ MeOH/products TBA SO4 þ TBA/products OH þ TBA/products
References (Crittenden et al., 1999)
i
k2 7:2 103 M1 s1 k3 ¼ 6:5 107 M1 s1 k4 ¼ 8:3 M1 s1 k5 ¼ 8:1 108 M1 s1 ðpH ¼ 5:8Þ 5:0 108 M1 s1 ðpH ¼ 5:0Þ 4:8 108 M1 s1 ðpH ¼ 4:8Þ
(Kolthoff and Miller, 1951) (Neta et al., 1988) (Yu et al., 2004)
k6 ¼ 5:5 105 M1 s1 k7 ¼ 1:4 107 M1 s1 k8 ¼ 6:9 105 M1 s1 k9 ¼ 1 1010 M1 s1 pKa1 ¼ 1:92
(Yu et al., 2004) (Buxton et al., 1990) (Buxton et al., 1988) (Klaning et al., 1991) (Yu et al., 2004)
k11 ¼ ð4:2±0:6Þ 108 M1 s1 k12 ¼ ð4:3±0:2Þ 109 M1 s1
(a) (a)
k13 ¼ ð7:6±0:6Þ 108 M1 s1 k14 ¼ ð5:7±0:2Þ 109 M1 s1
(a) (a)
k15 ¼ 1:2 109 M1 s1 k16 ¼ 5:9 109 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
1 k17 ¼ 2:35 107 M1 C s 1 k18 ¼ 3 108 M1 s C
(a) (Hoigne, 1998)
k19 ¼ 1:6 106 M1 s1 k20 ¼ 6:1 106 M1 s1 k21 ¼ 8:5 106 M1 s1 k22 ¼ 4 108 M1 s1 pKa2 ¼ 6:3 pKa3 ¼ 10:3
(Zuo et al., 1999) (Zuo et al., 1999) (Buxton et al., 1988) (Buxton et al., 1988) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996)
k25 < 7:2 104 M1 s1 k26 ¼ 1:2 106 M1 s1 k27 ¼ 2 104 M1 s1 k28 ¼ 1:5 105 M1 s1 pKa4 ¼ 2:1 pKa5 ¼ 7:2 pKa6 ¼ 12:3
(Neta et al., 1988) (Neta et al., 1988) (Buxton et al., 1988) (Buxton et al., 1988) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996)
k32 ¼ 1:1 107 M1 s1 k33 ¼ 9:7 108 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
k34 ¼ 4 105 M1 s1 k35 ¼ 6:0 108 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
(Neta et al., 1988)
Note: 1) F is quantum yield, which is 1.4 in eq. (1) and (2) ε is persulfate's extinction coefficient, which is 21.58 M1 cm1 determined using a spectrometer; 3) ‘a’ means the data was determined in this work; 4) εgeosmin ¼ 540 M1 cm1 (Jo et al., 2011), ε2-MIB ¼ 776 M1 cm1 (Jo et al., 2011), εBA ¼ 760 M1 cm1 (Guan et al., 2011), εNOM ¼ 0.066 mg1 cm1 L which was determined in this work; 5) k21 was chosen 7.2 104 M1 s1 for model.
2011). The guidelines of either 2-MIB or geosmin for drinking water in China, South Korea and Japan are set at 10 ng/L (MOH and SAC, 2006; Koester, 2011). Due to their extremely low odor threshold concentrations (several ng/L), conventional water treatment processes using coagulation, sedimentation, sand filtration and chlorination are ineffective in removing these compounds (Srinivasan and Sorial, 2011; Antonopoulou et al., 2014). As a result, additional eliminating processes such as oxidation, adsorption or membrane filtration are generally required for such taste and odor problems (Srinivasan and
Sorial, 2011; Agus et al., 2011). Sedlak and co-researchers suggest that reverse osmosis membrane, ozonation followed by biological activated carbon and advanced oxidation technologies (AOTs) such as UV combined with hydrogen peroxide (UV/H2O2) could remove odors from water effectively (Agus et al., 2011). AOTs which typically involve strategies for generating hydroxyl radicals (OH), due to high redox potential (1.9e2.7 V), are efficient in degrading such pollutants (Buxton et al., 1988). Examples of such AOTs including UV/H2O2, vacuum UV, O3/H2O2, electrochemical oxidation and ultrasonic
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
irradiation, all of which have been shown to degrade 2-MIB and geosmin effectively (Jo et al., 2011; Kutschera et al., 2009; Mizuno et al., 2011; Li et al., 2010; Song and O'Shea, 2007; Antonopoulou et al., 2014). There has been recent interest in the generation of the sulfate radical (SO4 ) in AOTs due to its high redox potential of 2.5e3.1 V which is comparable to OH (Neta et al., 1988). However, SO4 is more selective than OH while still reacting rapidly with many organic substrates (Neta et al., 1988). SO4 can be generated from the activation of peroxydisulfate (S2O2 8 ) or peroxymonosulfate (HS2O5 ) by using UV, heat, base, or transition metals (Lau et al., 2007; Guan et al., 2011; Waldemer et al., 2007; Furman et al., 2010; Zou et al., 2013; Zhang et al., 2013). Numerous studies have demonstrated that SO4 can degrade endocrine disrupting compounds, chlorinated compounds, drugs, perfuluorinated compounds and algal toxins (Lau et al., 2007; Guan et al., 2011; Waldemer et al., 2007; Guan et al., 2013; Gao et al., 2012; Hori et al., 2005; Antoniou et al., 2010). Persulfate irradiated by UV was proposed to generate SO4 through reaction 1 in Table 1 with a quantum yield of 1.4 mol Es1 (254 nm) (Mark et al., 1990). In this process, SO4 are produced with activation of persulfate without the production of OH (Mark et al., 1990). However, OH can also be generated when SO4 reacts with water at a rate constant of 1 1 8.3 M s (reaction 4) (Yu et al., 2004), or with OH at a rate constant of 6.5 107 M1 s1 under alkaline conditions (re action 3) (Neta et al., 1988). Thus, SO4 generation may be an attractive AOTs strategy for removing 2-MIB and geosmin. However, to date the degradation efficiency and rates by SO4 have not been reported for 2-MIB and geosmin. The objectives of this study were (1) to investigate the effects of water chemistry and kinetics of the removal of 2-MIB and geosmin by UV (254 nm)/persulfate; (2) to determine the reaction rates between 2-MIB (and geosmin) and SO4 ; (3) and to model the removal of 2-MIB and geosmin in this process with water quality that is likely to be encountered in water treatment.
2.
Experimental section
2.1.
Chemicals
Chemical solutions were prepared with reagent-grade chemicals and ultra-pure water (18.2 MU cm) produced using a MilliQ biocel system. Potassium peroxydisulfate (PDS), benzoic acid (BA), sodium phosphate monobasic monohydrate, sodium phosphate dibasic, potassium hydroxide and N,Ndiethyl-p-phenylenediamine (DPD) were purchased from SigmaeAldrich, USA. 2-MIB and geosmin were supplied by Wako, Japan. Suwannee River natural organic matter (NOM) obtained from International Humic Substances Society was dissolved into ultra-pure water and filtered through a glassfiber membrane of 0.45 mm pore size (Whatman) to make a stock NOM solution. H2O2 solution (35% v/v) bought from Alfa Aesar, USA was standardized by colorimetric method using DPD (Bader et al., 1988). Phosphoric acid (HPLC grade) was obtained from Dima-Tech Inc., and methanol (MeOH) and nhexane of HPLC grade were obtained from Thermo Fisher
225
Scientific Inc. Tert-butyl-alcohol (TBA) was purchased from Tianjin Chemical Reagent Co., Ltd., China. Other reagents were obtained from Sinopharm Chemical Reagent Co., Ltd., China. All of the chemicals were used as received without further purification. Two real water samples (A and B) for simulating the degradation of 2-MIB and geosmin in realistic water matrices were collected from two drinking water productions. The source water of sample A is from a reservoir, while the other one is from a river. Both the two samples which went through coagulation, sedimentation and filtration processes were further filtered through a 0.45 mm glass fiber membrane (Whatman) prior to using. The water quality parameters of the samples are shown in Table S2. 2-MIB or geosmin was spiked to the water samples to prepare odor-containing water.
2.2.
Experimental procedures
A 600 mL sealed cylindrical borosilicate glass reactor equipped with a 6 W low-pressure Hg UV lamp (254 nm, GPH 135T5 L/4, Heraeus Noblelight) was used to conduct photochemical experiments (Fig. S1). The photon flux (I0) entering the solution from the UV source, effective path length and the corresponding average fluency rate were calculated to be 0.757 mE s1, 3.02 cm and 3.82 103 mE s1 cm2 (1.79 mW cm2), respectively, using the methods described in our previous research (Li et al., 2012). Temperature was controlled at 20 ± 1 C using a thermostat (THD-2015, Tianheng, Ningbo, China). Phosphate buffer (2 mM) was used to adjust pH values from 4.0 to 8.0. Samples were withdrawn at predetermined time intervals and quenched using excess sodium sulfite. All experiments were replicated independently at least two times, and the error bars represent the standard deviation among replicates.
2.3.
Analytical methods
Gas chromatography (GC) (model 6890, Agilent, Santa Clara, USA) was employed to measure the concentrations of 2-MIB and geosmin which were extracted by n-hexane. The chromatograph was coupled with a flame ionization detector (FID) and a HP5 column (30 m 0.25 mm, ID 0.32 mm). Other conditions and the method detection limit can be found in SI Text S1 and Table S1. The concentration of total organic carbon (TOC) was measured using a TOC analyzer (Multi N/C 3100, Analytik Jena, Berlin, Germany). Solution pH was measured by a pH meter (UB-7, Denver Instrument, USA). The extinction coefficients of some chemicals and absorbance of experimental solutions were determined by using a Varian Cary 300 spectrometer, USA. The concentration of persulfate was determined by iodometric method (Nagi et al., 1959). The concentrations of carbonate and bicarbonate were calculated from the pKa of carbonic acid (H2CO3) and pH assuming equilibrium. Details of the calculations are shown in SI Text S2. Bromide and chloride were determined by an ion chromatograph (ICS-3000, Dionex, USA). BA was analyzed by HPLC using a binary HPLC pump (model 1525), a dual l absorbance detector (model 2475) which was set at 227 nm, an autosampler (model 717) and a reverse-phase C18 column (5 mm 4.6 mm 150 mm) (all
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Fig. 1 e Degradation of 2-MIB (a) and geosmin (b) in UV/PDS process. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, [PDS]0 was 200 mM for PDS alone, while 10 mM for UV/PDS, I0/V ¼ 1.26 mE s¡1 L¡1.
from Waters, Milford, USA). An eluant water (10.0 mM phosphoric acid) and methanol (50:50, v/v %) at a flow rate of 1.0 mL/min was used in the system. Reaction rate constants of 2-MIB and geosmin with SO4 and OH were measured using a relative rate technique (Aschmann et al., 2011), in which BA whose reaction rate constant (k15 ¼ 1.2 109 M1 s1, k16 ¼ 5.9 109 M1 s1) is reliably known was chosen as a reference compound (see Text S3 for details) (Neta et al., 1988; Buxton et al., 1988). 1 mM TBA was added in the solution to eliminate OH in the UV/PDS process when determining the reaction rate constant of SO4 reacting with 2-MIB or geosmin. Experiments were also conducted in UV/H2O2 process to determine the reaction rate constants of OH reacting with 2-MIB or geosmin.
primarily on SO4 formed from activation of persulfate (reaction 1) and OH (reactions 3 and 4). The concentration of persulfate was assumed to keep stable as it changed insignificantly when the reaction time was less than 900 s (Fig. S2). Then based on the steady-state assumption, the kinetics of target compounds degradation in UV/PDS can be modeled as follows:
d½C ¼ k1;C ½SO4 ss þ k2;C ½OH ss ½C ¼ kso;C ½C dt
where the concentrations of OH and SO4 are taken to be at steady-state with concentrations [OH]ss and [SO4 ]ss, respectively, ‘C’ is target compound which in this case may be either 2-MIB or geosmin, k1,C and k2,C are reaction rate constants of target compound with SO4 and OH, respectively. It is assumed that SO4 is formed from photolysis of persulfate (reaction 1) (Crittenden et al., 1999); OH is formed from the oxidation of water and hydroxyl ions (Neta et al., 1988; Yu et al., 2004). Assuming steady-state, the net formation rates
2.4.
Modeling equations
The model was established based on the hypothesis that the degradation of target compounds in UV/PDS process depends
(1)
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of SO4 and OH are zero and the concentrations of SO4 and OH can be expressed as eqs. (2) and (3), respectively (see Text S4 in SI for details):
½SO4 ss ¼
½OH ss ¼
fI0 f ð1 10A Þ=V P þ k1;C ½C þ ki ½Si k3 ½OH þ k4 ½H2 O þ k6 S2 O2 8
(2)
i
k3 ½OH þ k4 ½H2 O P k2;C ½C þ k7 S2 O2 þ ki' ½Si 8
i
fI0 f ð1 10A Þ=V P k3 ½OH þ k4 ½H2 O þ k6 S2 O2 þ k1;C ½C þ ki ½Si 8
(3)
i
SO4
2-MIB and geosmin were difficult to be degraded by only PDS oxidation and UV irradiation (254 nm), it could be concluded that SO4 and/or OH produced in this system, via reactions 1 and 4 were primarily responsible for the observed degradation. From the incerting small figures in Fig. 1a and b, the pseudo-first-order rate constants for 2-MIB and geosmin degradation by using UV/PDS process were calculated to be 3.16 103 s1 and 4.80 103 s1, respectively.
1
where Si are scavengers of or OH , I0 (E s ) is the photons flux, G is the fraction of light absorbed by PDS, f (mol E s1) is the quantum yield of SO4 , A (cm1) is the absorbance of solution, and V (L) is the volume of solution. The values of kso;C for target compound degradation and the contributions of SO4 and OH under a range of conditions are then calculated by simultaneously solving eqs. (1)e(3). As phosphate radical is a kind of weak oxidant (Crittenden et al., 1999), its contribution to the degradation to 2-MIB and geosmin was ignored.
3.
Results and discussion
3.1. PDS
Removal efficiency of 2-MIB and geosmin using UV/
Fig. 1 shows the degradation of 238 nM (40 mg/L) 2-MIB and 219 nM (40 mg/L) geosmin in the UV/PDS process at pH 7.0 with a 2 mM phosphate buffer. Negligible degradation of 2-MIB and geosmin was observed using PDS alone at a concentration of 200 mM over a period of 1800 s, indicating little to no oxidation of 2-MIB and geosmin by PDS directly. Less than 3% of 2-MIB and 6% of geosmin were degraded, respectively, within 900 s under the UV irradiation alone (I0/V ¼ 1.26 mE s1 L1), which were in accordance with other reports and consistent with their weak UV absorption at 254 nm (Jo et al., 2011; Kutschera et al., 2009). In contrast, UV/PDS process showed high levels of degradation efficiency for both 2-MIB (86.0%) and geosmin (94.5%) at only 10 mM PDS dosages with a 600 s contact time. As
3.2. Reaction rate constants of SO4 reacting with 2-MIB or geosmin Fig. 2 shows a reaction rate constant ratio between 2-MIB (or geosmin) and BA with SO4 k11/k15 ¼ 0.35 ± 0.05 (or k13/ k15 ¼ 0.63 ± 0.05) at pH 7.0 achieved by using relative rate technique. With a reaction rate constant of BA with SO4 k15 ¼ 1.2 109 M1 s1 (Neta et al., 1988), the reaction rate constants k11 ¼ (4.2 ± 0.6) 108 M1 s1 and k13 ¼ (7.6 ± 0.6) 108 M1 s1 for 2-MIB and geosmin, respectively, could be achieved. The ionic strength affects the reac tion rate constants of SO4 with lots of organic pollutants due to the negative charges at high pHs (Rickman and Mezyk, 2010). However, the reaction rate constant of SO4 reacting with either 2-MIB or geosmin was expected to be little influenced by ionic strength due to their alcoholic structure that the hydroxyl group is bonded to the saturated carbon atom and isn't dissociated at pH 7. Using the same method in a UV/ H2O2 process with a rate constant of BA reacting with OH k16 ¼ 5.9 109 M1 s1 (Buxton et al., 1988), we could calculate that the rate constants of 2-MIB and geosmin reacting with OH were (4.3 ± 0.2) 109 M1 s1 and (5.7 ± 0.2) 109 M1 s1, respectively, based on the rate constant ratio shown in Fig. S3. The values were slightly lower than those reported in von Gunten's research (5.09 109 M1 s1 for 2-MIB and 7.8 109 M1 s1 for geosmin) (Peter and von Gunten, 2007).
3.3.
Roles of sulfate radical and hydroxyl radical
MeOH and TBA are usually used to distinguish SO4 and OH based on reactions 32e35. MeOH is considered as an effective quencher for both SO4 (k32 ¼ 1.1 107 M1 s1) and OH (k33 ¼ 9.7 108 M1 s1) (Buxton et al., 1988; Neta et al., 1988).
Fig. 2 e Determination of the reaction rate constants of SO4 reacting with 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 5 mM, [Geosmin]0 ¼ 5 mM, [BA]0 ¼ 5 mM, [PDS]0 ¼ 100 mM, [TBA]0 ¼ 1 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1.
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Fig. 3 e Impacts of co-existing MeOH and TBA on the pseudo-first-order constants of 2-MIB and geosmin. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1. TBA is an effective quencher for OH (k35 ¼ 6 108 M1 s1) (Buxton et al., 1988), but not for SO4 (k34 ¼ 4 105 M1 s1) (Neta et al., 1988). So the difference of degradation rate of target compounds could be used to distinguish the
contribution of SO4 and OH. Co-existence of either MeOH or TBA inhibited the degradation of 2-MIB (or geosmin), while the inhibiting ability of MeOH is stronger than that of TBA (Fig. 3 and Fig. S4). It proved that both SO4 and OH are existing in the experimental condition and played important role in the degradation of 2-MIB and geosmin. Fig. 3 shows the pseudo-first-order rate constants of 2-MIB and geosmin degradation, and the modeling results of the individual and overall contribution of SO4 and OH at different dosages of co-existing MeOH or TBA at pH 7.0 with 2 mM phosphate buffer. The pseudo-first-order rate constants of both 2-MIB and geosmin in the UV/PDS process decreased with increasing the dosages of MeOH or TBA. The modeling results were in good accordance with the experimental data. The modeling results also show that MeOH and TBA mainly suppressed the contribution of OH at low dosages, due to rapid reactions (k33 ¼ 9.7 108 M1 s1, k35 ¼ 6 108 M1 s1). In addition, the presence of MeOH can also inhibit the formation of OH through reaction 4 in Table 1 by consuming SO4 . Initially, the pseudo-first-order rate constants of 2-MIB and geosmin degradation decreased rapidly. With further increasing the dosages of MeOH or TBA, the decreasing speed of the pseudo-first-order rate constants of 2-MIB and geosmin degradation reduced dramatically. The pseudo-first-order rate constant of 2-MIB decreased by 96% (or 94.3%) with increasing the concentration of MeOH (or TBA) from 0 to 1000 mM. While it decreased by 82.7% and 73.3% with adding 100 mM MeOH and TBA, respectively. The pseudo-first-order rate constant of geosmin followed a similar trend to that of 2-MIB. It reduced by 82.7% (or 72.1%) in the presence of 100 mM MeOH (or TBA) and 86.8% (or 76.8%) in the presence of 1000 mM MeOH (or TBA). The contribution of SO4 became the vital reason for the degradation of 2-MIB or geosmin at high dosages of MeOH or TBA. The decreasing speed of the pseudo-first-order rate constants of 2-MIB and geosmin degradation reduced dramatically after the initial fast decreasing stage was due to that the reaction rate con stant of SO4 reacting with MeOH or TBA was much slower than that of OH. When the dosage of MeOH or TBA was zero, the contribution of OH was about 3.5 times and 2.0 times higher than that of SO4 to 2-MIB and geosmin degradation, respectively. The results further suggested that plenty of SO4
Fig. 4 e Impacts of pH on the pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
would react with water to produce OH in the experimental system (reaction 4).
3.4.
Effects of reaction pH
Degradation rate of 2-MIB and geosmin decreased with increasing pH from 4.0 to 8.0 (Fig. S5). Fig. 4 shows the experimental pseudo-first-order rate constants of 2-MIB and geosmin decreased from 5.59 103 s1 and 8.73 103 s1 to 1.89 103 s1 and 3.35 103 s1, respectively, as pH increased from 4.0 to 8.0. In the range of pH 5.0 to 8.0, the pseudo-first-order rate constants of both 2-MIB and geosmin decreased most quickly. The modeling results shown in Fig. 4 were in accordance with the overall trend of experimental observations, though the experimental values of 2-MIB were lower than those predicted from the model. From the results, it is reasonable to conclude that both SO4 and OH play important roles in the experimental system, but that the contribution of OH was much higher than that of SO4 . Though persulfate dissociation can be enhanced by acid through reaction 2, it can only be important in strong acidic conditions and could be eliminated in the pH range of interest here (Kolthoff and Miller, 1951). The quantum efficiency of photodissociation of persulfate ion was identical at different pHs (Yu et al., 2004). So the formation of sulfate radicals would not be affected by pH, which was in accordance with the degradation of persulfate shown in Fig. S6. Fig. 2a and S7a show that the reaction rate constants be tween SO4 and 2-MIB were stable at pH 4.0, 7.0 and 8.0 ((4.2 ± 0.6) 108 M1 s1 z (4.2 ± 0.6) 108 M1 s1 z (4.5 ± 0.6) 108 M1 s1). The reaction rate constants between SO4 and geosmin calculated from Fig. 2b and S7b show the same trend ((7.5 ± 0.3) 108 M1 s1 z (7.6 ± 0.6) 108 M1 s1 z (7.3 ± 0.7) 108 M1 s1). The results suggested that the reaction rate constants of SO4 for both 2-MIB and geosmin were not affected by pH. 2 The distributions of H2PO 4 and HPO4 are affected by pH (reactions 29e31). There were 2 mM phosphate buffer whose concentration was over 8000 times more than the initial
229
concentration of either 2-MIB (238 nM) or geosmin (219 nM) in the solution, which were scavengers in the UV/PDS process based on reactions 25e28. The reaction rate constant between 6 1 1 s ) is higher than that SO4 and HPO2 4 (k26 ¼ 1.2 10 M between SO4 and H2PO4 (k25 < 7.2 104 M1 s1) (Neta et al., 1988). In addition, the reaction rate constant of OH reacting 5 1 1 s ) is also higher than that with HPO2 4 (k28 ¼ 1.5 10 M reacting with H2PO4 (k27 ¼ 2 104 M1 s1) (Buxton et al., 1988). So higher concentration of HPO2 4 can reduce the theoretical steady-state concentrations of SO4 and OH based on eqs. (2) and (3). The distribution of H2PO4 and HPO2 4 shown in Fig. S8 was also in accordance with the variation of the pseudo-firstorder rate constants of 2-MIB and geosmin. The percent of increased dramatically from pH 5 to pH 8, which HPO2 4 increased the consumption of SO4 and OH and decreased the steady-state concentration of both SO4 and OH. So the pseudo-first-order rate constants decreased obviously from pH 5 to pH 8, while kept stable when the pH was less than 5 where nearly none of HPO2 4 existed. The discussions suggested that pH would not affect the degradation of 2-MIB and geosmin directly, but only through 2 affecting the distribution of H2PO 4 and HPO4 in the experimental process. It should be noted that the contribution of diphosphate radical to the degradation of 2-MIB or geosmin was eliminated when applying the steady-state model.
3.5.
Effects of initial persulfate dosages
The degradation rates of 2-MIB and geosmin increased with increasing persulfate dosage from 0 mM to 50 mM (Fig. S9). Fig. 5 shows that the experimental and calculated data of overall pseudo-first-order rate constants of both 2-MIB and geosmin were in accordance well. The experimental values increased with increasing the dosage of persulfate. The formation rate of SO4 increased with increasing the dosage of persulfate (re action 1), then the ideal steady-state concentrations of SO4 and OH were expected to increase based on eqs. (2) and (3). The modeling results of the individual contributions of SO4 and OH to 2-MIB and geosmin degradation were also shown
Fig. 5 e Impacts of PDS dosages on the pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
230
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
Fig. 6 e Impacts of alkalinity on the pseudo-first-order constants of 2-MIB and geosmin. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
in Fig. 5. OH was the main reactive species, whose contribution to 2-MIB or geosmin degradation was higher than that of SO4 at 10e50 mM persulfate dosages all the time.
3.6.
Effects of alkalinity
Different dosages of bicarbonate (0e2 mM) were added in the reaction solutions at pH 7.0 with 2 mM phosphate buffer to study the effects of alkalinity on the degradation of 2-MIB and geosmin (Fig. S10), and to discuss the individual and the sum of contribution of SO4 and OH to either 2-MIB or geosmin degradation (Fig. 6). At the chosen pH, carbonate species include bicarbonate (HCO 3 ) and H2CO3* (about 16.7%) whose rate constants of reacting with OH and SO4 are unknown. 2 While carbonate (CO3 ) could be eliminated as its percentage 2 of all the carbonate species (i.e., H2CO*3, HCO 3 , CO3 ) was very low (about 0.04%). We here treated it as Fang's method that the rate constants of H2CO3* reacting with OH and SO4 were assumed to be similar with those of HCO 3 reacting with OH and SO4 , respectively (Fang et al., 2014), which still needs to 2 be further proved. HCO 3 and CO3 react with OH or SO4 to produce carbonate radical (reactions 19e22 in Table 1) which is a weak oxidant (Buxton et al., 1988; Zuo et al., 1999). But the
oxidation reactions are insignificant and can be ignored for treating most organic pollutants (Crittenden et al., 1999). So the contribution of carbonate radicals in the experimental system was not considered here. Fig. 6 shows the overall pseudo-first-order rate constants of 2-MIB and geosmin degradation decreased from (3.00 ± 0.17) 103 s1 and 3 1 s (4.58 ± 0.25) 103 s1 without HCO 3 to (0.30 ± 0.03) 10 and (0.57 ± 0.05) 103 s1 in the presence of 2 mM HCO 3, respectively. The results show that bicarbonate could reduce the degradation efficiency of both 2-MIB and geosmin effec tively, which was because it could scavenge both SO4 and OH 6 1 1 at relative high reaction rates (1.6 10 M s for SO4 , 8.5 106 M1 s1 for OH) (Buxton et al., 1988; Zuo et al., 1999). Both the steady-state concentrations of OH and SO4 could be reduced by bicarbonate, but the inhibition effect on OH was much stronger than that on SO4 . It is not only because that the reaction rate of OH is higher than that of SO4 , but also due to the fact that the steady-state concentration of OH is based on the steady state concentration of SO4 (eq. (3)). The concentration of 2-MIB and geosmin could be detected as low as a few ng/L in natural water (Pirbazari et al., 1993). While the bicarbonate concentration varied from several dozen to several hundred mg/L (Crittenden et al., 2005), which was
Fig. 7 e Impacts of co-existing NOM on pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
231
much higher than the concentration of 2-MIB and geosmin. So alkalinity should be an important scavenging factor for 2-MIB and geosmin degradation in water production when using UV/ PDS process.
3.7.
Effects of NOM
Different dosages of NOM (0e3 mg/L) were added in the samples to elucidate their effects on the degradation of 2-MIB and geosmin (Fig. S12), and the individual and the total con tributions of SO4 and OH to either 2-MIB or geosmin degradation at pH 7.0 with 2 mM phosphate buffer (Fig. 7). Fig. 7 shows that the experimental pseudo-first-order rate constants of both 2-MIB and geosmin were decreased with increasing dosages of co-existing NOM, which were in accordance with the modeling overall pseudo-first-order rate constants well. The pseudo-first-order rate constant of OH for both 2-MIB and geosmin decreased sharply at low dosages of NOM, with a much faster decreasing speed than that of SO4 . The results can be explained by a steady-state concentration of OH that decreased much faster than that of SO4 with coexistence of scavengers through eqs. (2) and (3). When the dosage of NOM exceeded 1.0 mg/L, the contribution of OH could be nearly eliminated. The extinction coefficient of the NOM used here was 0.066 mg1 cm1 L (by TOC) at 254 nm, so it could act as an inner filter to reduce the rate of the activation of persulfate to produce SO4 based on reaction 1 in Table 1. The filtering effect of NOM was also incorporated into the modeling. The reaction 1 rate constant of NOM reacting with OH is 3 108 M1 C s (Hoigne, 1998), and the reaction rate constant of NOM reacting 1 (see Text S5 and Fig. S11 with SO4 is about 2.35 107 M1 C s in SI for details). So NOM competed the reactive radicals with target compound (2-MIB or geosmin) based on reactions 17 and 18 to reduce the steady-state concentration of radicals. The reaction rate constant between NOM and OH is over 10 times higher than that between NOM and SO4 , and the steady-state concentration of OH was affected by that of SO4 via eq. (3). So the decreasing speed of pseudo-first-rate con stant for OH is much quicker than that for SO4 . It is reasonable to conclude that the fast decreasing pseudo-first-order rate constants of 2-MIB and geosmin with low dosages of NOM was mainly due to the sharply decreased contribution of OH (Fig. 7). The pseudo-first-order rate constants of 2-MIB and geosmin decreased much more slowly with further increasing the dosage of NOM, which was because that only SO4 that was competed by NOM with a slower rate constant played roles at high dosage of NOM. It should be noted that NOM could be activated by UV light to produce reactive species (such as singlet oxygen) (Latch and McNeill, 2006; Zhang and Hsu-Kim, 2010), whose contribution to 2-MIB (or geosmin) degradation was eliminated in the model.
3.8.
Degradation efficiency in real water
Two different real water samples (A and B) were used to elucidate the degradation efficiencies of 2-MIB and geosmin (Fig. S13) and the application of the kinetic model in real water (Fig. 8). Fig. 8 shows the experimental values were somewhat lower than the calculated values from the kinetic model. The
Fig. 8 e Application of the kinetic model in two different real waters. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 100 mM, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1. experimental pseudo-first-order rate constants of 2-MIB were about 78.4% and 76.9% of the modeling values in Samples A and B, respectively, and those of geosmin were 81.0% and 77.5% of the modeling values, respectively. Fig. S14 shows the impacts of chloride (0e2 mM) on the degradation of 2-MIB and geosmin was negligible. So the effects of halides were negligible to the degradation kinetics of 2-MIB and geosmin as the concentrations in both the two real water samples were less than 0.6 mM (Table S2). The small discrepancy between experimental data and calculated values from the model was likely due to the difference of the reaction rate constant of NOM in the real waters reacting with radicals (SO4 or OH) from that of Suwannee River NOM used in the model. Variability in the properties of NOM is likely to yield variations in their reaction rate constants with the radicals. For example, the range of reaction rate constants between organic matter 1 at room temperature and OH was 1.21e9.37 108 M1 C s (Mckay et al., 2011), which leads to the difference of the scavenging effect of NOM on OH. The variation should be also existed for SO4 . From the modeling results considering NOM and bicarbonate, the reaction rate constant of SO4 reacting with NOM from both the two real water samples was esti1 which was a little higher mated be about 3.7 107 M1 C s than the reaction rate constant of Suwannee River NOM 1 (2.35 107 M1 C s ). Besides, the effects of other water components, such as the coexisting micropollutants, may also contribute to the discrepancy.
4.
Conclusions
Both 2-MIB and geosmin in water can be removed effectively by using UV/persulfate process. The second order rate con stants for 2-MIB and geosmin reacting with SO4 were esti8 M1 s1 and mated to be (4.2 ± 0.6) 10 8 1 1 (7.6 ± 0.6) 10 M s respectively at a pH of 7.0. A model based on steady-state assumption suggested that both SO4 and OH contributed to the degradation of 2-MIB and geosmin. The contributions of the OH to 2-MIB and geosmin
232
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
degradation were higher than the contribution from SO4 in ultra-pure water. While the contributions of OH to 2-MIB and geosmin degradation were inhibited much more significantly than those from SO4 by the co-existence of scavengers including bicarbonate and NOM. Increasing dosages of persulfate increased the degradation of 2-MIB and geosmin. pH did not affect the degradation of 2-MIB and geosmin directly. Faster degradations of both 2-MIB and geosmin in acidic condition when using phosphate buffer were achieved due to the different scavenging effects of hydrogen phosphate and dihydrogen phosphate. NOM and bicarbonate are likely to be the main radical scavengers in natural waters when using UV/ persulfate process to control 2-MIB and geosmin as they can consume both SO4 and OH effectively. The process was also effective in removing 2-MIB and geosmin in real water, and the degradation could also be well described by the model developed here.
Acknowledgments This study was supported by the Natural Science Foundation of China (grants 51178134, 51108117, 21307057 and 51108111). It was also supported by China's Fund for National Creative Research Groups (grant 51121062), and the State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology (grant ES201006). The authors greatly thank Dr Yinghong Guan for the discussion and helpful comments on the model proposed here.
Appendix A. Supplementary data Supplementary data related to this article can be found at http://dx.doi.org/10.1016/j.watres.2014.11.029.
references
Agus, E., Lim, M.H., Zhang, L., Sedlak, D.L., 2011. Odorous compounds in municipal wastewater effluent and potable water reuse systems. Environ. Sci. Technol. 45 (21), 9347e9355. Antoniou, M.G., de la Cruz, A.A., Dionysiou, D.D., 2010. Intermediates and reaction pathways from the degradation of Microcystin-LR with sulfate radicals. Environ. Sci. Technol. 44 (19), 7238e7244. Antonopoulou, M., Evgenidou, E., Lambropoulou, D., Konstantinou, I., 2014. A review on advanced oxidation processes for the removal of taste and odor compounds from aqueous media. Water Res. 53 (0), 215e234. Aschmann, S.M., Arey, J., Atkinson, R., 2011. Kinetics and products of the reaction of OH radicals with 3-Methoxy-3methyl-1-tubanol. Environ. Sci. Technol. 45 (16), 6896e6901. Bader, H., Sturzenegger, V., Hoigne, J., 1988. Photometric method for the determination of low concentrations of hydrogen peroxide by the peroxidase catalyzed oxidation of N,Ndiethyl-p-phenylenediamine (DPD). Water Res. 22 (9), 1109e1115. Buxton, G.V., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical-review of rate constants for reactions of hydrated
electrons, hydrogen-atoms and hydroxyl radicals (OH/O) in aqueous solution. J. Phys. Chem. Ref. Data 17 (2), 513e886. Buxton, G.V., Salmon, G.A., Wood, N., 1990. A pulse radiolysis study of the chemistry of oxysulphur radicals in aqueous solution. In: Physico-chemical Behaviour of Atmospheric Pollutants. Springer, Netherlands, pp. 245e250. Crittenden, J.C., Hu, S., Hand, D.W., Green, S.A., 1999. A kinetic model for H2O2/UV process in a completely mixed batch reactor. Water Res. 33 (10), 2315e2328. Crittenden, J.C., Trussell, R.R., Hand, D.W., Howe, K.J., Tchobanoglous, G., 2005. Water Treatment: Principles and Design, second ed. John Wiley & Sons, Inc, Hoboken, New Jersey. Fang, J., Fu, Y., Shang, C., 2014. The roles of reactive species in micropollutant degradation in the UV/free chlorine system. Environ. Sci. Technol. 48 (3), 1859e1868. Furman, O.S., Teel, A.L., Watts, R.J., 2010. Mechanism of base activation of persulfate. Environ. Sci. Technol. 44 (16), 6423e6428. Gao, Y.-Q., Gao, N.-Y., Deng, Y., Yang, Y.-Q., Ma, Y., 2012. Ultraviolet (UV) light-activated persulfate oxidation of sulfamethazine in water. Chem. Eng. J. 195196 (1), 248e253. Guan, Y., Ma, J., Li, X., Fang, J., Chen, L., 2011. Influence of pH on the formation of sulfate and hydroxyl radicals in the UV/ peroxymonosulfate system. Environ. Sci. Technol. 45 (21), 9308e9314. Guan, Y., Ma, J., Ren, Y., Liu, Y., Xiao, J., Lin, L., Zhang, C., 2013. Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals. Water Res. 47 (14), 5431e5438. Hoigne, J., 1998. Chemistry of aqueous ozone, and transformation of pollutants by ozone and advanced oxidation process. In: The Handbook of Environmental Chemistry Quality and Treatment of Drinking Water. Springer, Berlin. Hori, H., Yamamoto, A., Hayakawa, E., Taniyasu, S., Yamashita, N., Kutsuna, S., 2005. Efficient decomposition of environmentally persistent perfluorocarboxylic acids by use of persulfate as a photochemical oxidant. Environ. Sci. Technol. 39 (7), 2383e2388. Jo, C.H., Dietrich, A.M., Tanko, J.M., 2011. Simultaneous degradation of disinfection byproducts and earthy-musty odorants by the UV/H2O2 advanced oxidation process. Water Res. 45 (8), 2507e2516. Klaning, U.K., Sehested, K., Appelman, E.H., 1991. Laser flash photolysis and pulse radiolysis of aqueous solutions of the fluoroxysulfate ion, SO4F. Inorg. Chem. 30 (18), 3582e3584. Koester, K., 2011. Measuring and Modeling Geosmin Removal from Horsetooth Reservoir Water by Powdered Activated Carbon for Selected Contact Times. Colorado State University, Fort Collins, Colorado. Kolthoff, I.M., Miller, I.K., 1951. The chemistry of persulfate. I. The kinetics and mechanism of the decomposition of the persulfate ion in aqueous medium. J. Am. Chem. Soc. 73 (7), 3055e3059. Kutschera, K., Bornick, H., Worch, E., 2009. Photoinitiated oxidation of geosmin and 2-methylisoborneol by irradiation with 254 nm and 185 nm UV light. Water Res. 43 (8), 2224e2232. Latch, D.E., McNeill, K., 2006. Microheterogeneity of singlet oxygen distributions in irradiated humic acid solutions. Science 311 (5768), 1743e1747. Lau, T.K., Chu, W., Graham, N.J.D., 2007. The aqueous degradation of butylated hydroxyanisole by UV/S2O2 8 : study of reaction mechanisms via dimerization and mineralization. Environ. Sci. Technol. 41 (2), 613e619. Li, M., Xue, Q., Zhang, Z., Feng, C., Chen, N., Lei, X., Shen, Z., Sugiura, N., 2010. Removal of geosmin (trans-1,10-dimethyl-
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
trans-9-decalol) from aqueous solution using an indirect electrochemical method. Electrochim. Acta 55 (23), 6979e6982. Li, X., Ma, J., Liu, G., Fang, J., Yue, S., Guan, Y., Chen, L., Liu, X., 2012. Efficient reductive dechlorination of monochloroacetic acid by sulfite/UV process. Environ. Sci. Technol. 46 (13), 7342e7349. Mark, G., Schuchmann, M.N., Schuchmann, H.-P., von Sonntag, C., 1990. The photolysis of potassium peroxodisulphate in aqueous solution in the presence of tertbutanol: a simple actinometer for 254 nm radiation. J. Photochem. Photobiol. A Chem. 55 (2), 157e168. Mckay, G., Dong, M.M., Kleinman, J.L., Mezyk, S.P., RosarioOrtiz, F.L., 2011. Temperature dependence of the reaction between the hydroxyl radical and organic matter. Environ. Sci. Technol. 45 (16), 6932e6937. Ministry of Health of The People's Republic of China (MOH), Standardization Administration of The People's Republic of China (SAC), 2006. Standards for Drinking Water Quality (GB 5749-2006). Beijing. Mizuno, T., Ohara, S., Nishimura, F., Tsuno, H., 2011. O3/H2O2 process for both removal of odorous algal-derived compounds and control of bromate ion formation. Ozone Sci. Eng. 33 (2), 121e135. Nagi, W., El Asmar, M.F., El Sadr, M.M., 1959. Iodometric method for determination of persulfate. Anal. Chem. 31 (11), 1870e1871. Neta, P., Huie, R.E., Ross, A.B., 1988. Rate constants for reactions of inorganic radicals in aqueous solution. J. Phys. Chem. Ref. Data 17 (3), 1027e1284. Peter, A., von Gunten, U., 2007. Oxidation kinetics of selected taste and odor compounds during ozonation of drinking water. Environ. Sci. Technol. 41 (2), 6896e6901. Pirbazari, M., Ravindran, V., Badriyha, B.N., Craig, S., McGuire, M.J., 1993. GAC adsorber design protocol for the removal of off-flavors. Water Res. 27 (7), 1153e1166. Rickman, K.A., Mezyk, S.P., 2010. Kinetics and mechanisms of sulfate radical oxidation of b-lactam antibiotics in water. Chemosphere 81 (3), 359e365.
233
Song, W., O'Shea, K.E., 2007. Ultrasonically induced degradation of 2-methylisoborneol and geosmin. Water Res. 41 (12), 2672e2678. Srinivasan, R., Sorial, G.A., 2011. Treatment of taste and odor causing compounds 2-methyl isoborneol and geosmin in drinking water: a critical review. J. Environ. Sci. 23 (1), 1e13. Stumm, W., Morgan, J.J., 1996. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, third ed. John Wiely & SONS, INC. Waldemer, R.H., Tratnyek, P.G., Johnson, R.L., Nurmi, J.T., 2007. Oxidation of chlorinated ethenes by heat-activated persulfate: kinetics and products. Environ. Sci. Technol. 41 (3), 1010e1015. Watson, S.B., Ridal, J., Boyer, G.L., 2008. Taste and odour and cyanobacterial toxins: impairment, prediction, and management in the Great Lakes. Can. J. Fish. Aquat. Sci. 65 (8), 1779e1796. Yu, X.-Y., Bao, Z.-C., Barker, J.R., 2004. Free radical reactions involving Cl, Cl 2 , and SO4 in the 248 nm photolysis of aqueous solutions containing S2O2 8 and Cl . J. Phys. Chem. A 18 (2), 295e308. Zhang, T., Zhu, H., Croue, J.-P., 2013. Magnetically separable CuFe2O4 spinel in water: efficiency, stability, and mechanism. Environ. Sci. Technol. 47 (6), 2784e2791. Zhang, T., Hsu-Kim, H., 2010. Photolytic degradation of methylmercury enhanced by binding to natural organic ligands. Nature Geosci. 3 (7), 473e476. Zou, J., Ma, J., Chen, L., Li, X., Guan, Y., Xie, P., Pan, C., 2013. Rapid acceleration of ferrous iron/peroxymonosulfate oxidation of organic pollutants by promoting Fe(III)/Fe(II) cycle with hydroxylamine. Environ. Sci. Technol. 47 (20), 11685e11691. Zuo, Z., Cai, Z., Katsumura, Y., Chitose, N., Muroya, Y., 1999. Reinvestigation of the acid-base equilibrium of the (bi) carbonate radical and pH dependence of its reactivity with inorganic reactants. Radiat. Phys. Chem. 55 (1), 15e23.
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Removal of 2-MIB and geosmin using UV/ persulfate: Contributions of hydroxyl and sulfate radicals Pengchao Xie a, Jun Ma a,*, Wei Liu a, Jing Zou a, Siyang Yue a, Xuchun Li b, Mark R. Wiesner c, Jingyun Fang d a
State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin 150090, China b School of the Environmental Science and Engineering, Zhejiang Gongshang University, Hangzhou 310012, China c Civil and Environmental Engineering, Duke University, Durham 27708-0287, USA d School of Environmental Science and Engineering, Sun Yat-sen University, Guangzhou 510275, China
article info
abstract
Article history:
2-methylisoborneol (2-MIB) and geosmin are two odor-causing compounds that are diffi-
Received 7 July 2014
cult to remove and the cause of many consumer complaints. In this study, we assessed the
Received in revised form
degradation of 2-MIB and geosmin using a UV/persulfate process for the first time. The
27 September 2014
results showed that both 2-MIB and geosmin could be degraded effectively using this
Accepted 17 November 2014
process. The process was modeled based on steady-state assumption with respect to the
Available online 26 November 2014
odor-causing compounds and either hydroxyl or sulfate radicals. The second order rate
constants for 2-MIB and geosmin reacting with the sulfate radical (SO4 ) were estimated to 8
1
1
8
1
and (7.6 ± 0.6) 10 M
1
Keywords:
be (4.2 ± 0.6) 10 M
UV/persulfate
contributions of the hydroxyl radical (OH) to 2-MIB and geosmin degradation were 3.5
Hydroxyl radical
times and 2.0 times higher, respectively, than the contribution from SO4 in Milli-Q water
Sulfate radical
with 2 mM phosphate buffer at pH 7.0. The pseudo-first-order rate constants (kso ) of both 2-
2-MIB
MIB and geosmin increased with increasing dosages of persulfate. Although pH did not
Geosmin
affect the degradation of 2-MIB and geosmin directly, different scavenging effects of
s
s
respectively at a pH of 7.0. The
hydrogen phosphate and dihydrogen phosphate resulted in higher values of kso for both 2MIB and geosmin in acidic condition. Bicarbonate and natural organic matter (NOM) inhibited the degradation of both 2-MIB and geosmin dramatically through consuming OH
and SO4 and were likely to be the main radical scavengers in natural waters when using UV/persulfate process to control 2-MIB and geosmin. © 2014 Elsevier Ltd. All rights reserved.
1.
Introduction
Cyanobacteria or blue-green algae can produce 2methylisoborneol (2-MIB) and geosmin which result in many
complaints of taste and odor in drinking water from consumer (Watson et al., 2008; Srinivasan and Sorial, 2011). The concentration of the two odor-causing compounds can even reach more than 1000 ng/L during algae bloom (Mizuno et al.,
* Corresponding author. Tel.: þ86 451 86282292; fax: þ86 451 86283010. E-mail address: [email protected] (J. Ma). http://dx.doi.org/10.1016/j.watres.2014.11.029 0043-1354/© 2014 Elsevier Ltd. All rights reserved.
224
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
Table 1 e Principle reactions in UV/persulfate process. No.
Reactions
UV
r
H
2 3 4
þ S2 O2 8 þ H /…!SO4 þ HSO4 þ OH SO4 þ OH /SO2 4 SO4 þ H2 O/HSO2 4 þ OH
5
SO4 þ SO4 /S2 O2 8
6 7 8 9 10 In the 11 12 In the 13 14 In the 15 16 In the 17 18 In the 19 20 21 22 23 24 In the 25 26 27 28 29 30 31 In the 32 33 In the 34 35
presence of
presence of
presence of
presence of
presence of
presence of
presence of
presence of
SO4
f ¼ εcb=A
þ
Reaction rate ¼ fI0 f ð1 10A Þ=V P A ¼ bðεc þ εi ci Þ
S2 O2 8 !2SO4
1
2 SO4 þ S2 O2 8 4S2 O8 þ SO4 OH þ S2 O2 8 4S2 O8 þ OH OH þ HSO 4 /SO4 þ H2 O OH þ SO4 /HSO 5 2 þ HSO 4 ⇔SO4 þ H 2-MIB SO4 þ 2 MIB/products OH þ 2 MIB/products geosmin SO4 þ geosmin/products OH þ geosmin/products BA BA þ SO4 /products BA þ OH /products NOM NOM þ SO 4 /products NOM þ OH /products 2 HCO3 /CO3 2 HCO 3 þ SO4 /SO4 þ HCO3 2 2 CO3 þ SO4 /SO4 þ CO3 HCO 3 þ OH /H2 O þ CO3 2 CO3 þ OH /OH þ CO3 H2 CO*3 ⇔Hþ þ HCO 3 2 þ HCO 3 ⇔H þ CO3 2 HPO4 /H2 PO4 SO4 þ H2 PO 4 /products 2 SO4 þ HPO2 4 /SO4 þ HPO4 OH þ H2 PO4 /HPO4 þ H2 O OH þ HPO2 4 /HPO4 þ OH þ H3 PO4 ⇔H þ H2 PO4 þ H2 PO 4 ⇔H þ HPO4 2 3 þ HPO4 ⇔H þ PO4 MeOH SO4 þ MeOH/products OH þ MeOH/products TBA SO4 þ TBA/products OH þ TBA/products
References (Crittenden et al., 1999)
i
k2 7:2 103 M1 s1 k3 ¼ 6:5 107 M1 s1 k4 ¼ 8:3 M1 s1 k5 ¼ 8:1 108 M1 s1 ðpH ¼ 5:8Þ 5:0 108 M1 s1 ðpH ¼ 5:0Þ 4:8 108 M1 s1 ðpH ¼ 4:8Þ
(Kolthoff and Miller, 1951) (Neta et al., 1988) (Yu et al., 2004)
k6 ¼ 5:5 105 M1 s1 k7 ¼ 1:4 107 M1 s1 k8 ¼ 6:9 105 M1 s1 k9 ¼ 1 1010 M1 s1 pKa1 ¼ 1:92
(Yu et al., 2004) (Buxton et al., 1990) (Buxton et al., 1988) (Klaning et al., 1991) (Yu et al., 2004)
k11 ¼ ð4:2±0:6Þ 108 M1 s1 k12 ¼ ð4:3±0:2Þ 109 M1 s1
(a) (a)
k13 ¼ ð7:6±0:6Þ 108 M1 s1 k14 ¼ ð5:7±0:2Þ 109 M1 s1
(a) (a)
k15 ¼ 1:2 109 M1 s1 k16 ¼ 5:9 109 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
1 k17 ¼ 2:35 107 M1 C s 1 k18 ¼ 3 108 M1 s C
(a) (Hoigne, 1998)
k19 ¼ 1:6 106 M1 s1 k20 ¼ 6:1 106 M1 s1 k21 ¼ 8:5 106 M1 s1 k22 ¼ 4 108 M1 s1 pKa2 ¼ 6:3 pKa3 ¼ 10:3
(Zuo et al., 1999) (Zuo et al., 1999) (Buxton et al., 1988) (Buxton et al., 1988) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996)
k25 < 7:2 104 M1 s1 k26 ¼ 1:2 106 M1 s1 k27 ¼ 2 104 M1 s1 k28 ¼ 1:5 105 M1 s1 pKa4 ¼ 2:1 pKa5 ¼ 7:2 pKa6 ¼ 12:3
(Neta et al., 1988) (Neta et al., 1988) (Buxton et al., 1988) (Buxton et al., 1988) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996) (Stumm & Morgan, 1996)
k32 ¼ 1:1 107 M1 s1 k33 ¼ 9:7 108 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
k34 ¼ 4 105 M1 s1 k35 ¼ 6:0 108 M1 s1
(Neta et al., 1988) (Buxton et al., 1988)
(Neta et al., 1988)
Note: 1) F is quantum yield, which is 1.4 in eq. (1) and (2) ε is persulfate's extinction coefficient, which is 21.58 M1 cm1 determined using a spectrometer; 3) ‘a’ means the data was determined in this work; 4) εgeosmin ¼ 540 M1 cm1 (Jo et al., 2011), ε2-MIB ¼ 776 M1 cm1 (Jo et al., 2011), εBA ¼ 760 M1 cm1 (Guan et al., 2011), εNOM ¼ 0.066 mg1 cm1 L which was determined in this work; 5) k21 was chosen 7.2 104 M1 s1 for model.
2011). The guidelines of either 2-MIB or geosmin for drinking water in China, South Korea and Japan are set at 10 ng/L (MOH and SAC, 2006; Koester, 2011). Due to their extremely low odor threshold concentrations (several ng/L), conventional water treatment processes using coagulation, sedimentation, sand filtration and chlorination are ineffective in removing these compounds (Srinivasan and Sorial, 2011; Antonopoulou et al., 2014). As a result, additional eliminating processes such as oxidation, adsorption or membrane filtration are generally required for such taste and odor problems (Srinivasan and
Sorial, 2011; Agus et al., 2011). Sedlak and co-researchers suggest that reverse osmosis membrane, ozonation followed by biological activated carbon and advanced oxidation technologies (AOTs) such as UV combined with hydrogen peroxide (UV/H2O2) could remove odors from water effectively (Agus et al., 2011). AOTs which typically involve strategies for generating hydroxyl radicals (OH), due to high redox potential (1.9e2.7 V), are efficient in degrading such pollutants (Buxton et al., 1988). Examples of such AOTs including UV/H2O2, vacuum UV, O3/H2O2, electrochemical oxidation and ultrasonic
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
irradiation, all of which have been shown to degrade 2-MIB and geosmin effectively (Jo et al., 2011; Kutschera et al., 2009; Mizuno et al., 2011; Li et al., 2010; Song and O'Shea, 2007; Antonopoulou et al., 2014). There has been recent interest in the generation of the sulfate radical (SO4 ) in AOTs due to its high redox potential of 2.5e3.1 V which is comparable to OH (Neta et al., 1988). However, SO4 is more selective than OH while still reacting rapidly with many organic substrates (Neta et al., 1988). SO4 can be generated from the activation of peroxydisulfate (S2O2 8 ) or peroxymonosulfate (HS2O5 ) by using UV, heat, base, or transition metals (Lau et al., 2007; Guan et al., 2011; Waldemer et al., 2007; Furman et al., 2010; Zou et al., 2013; Zhang et al., 2013). Numerous studies have demonstrated that SO4 can degrade endocrine disrupting compounds, chlorinated compounds, drugs, perfuluorinated compounds and algal toxins (Lau et al., 2007; Guan et al., 2011; Waldemer et al., 2007; Guan et al., 2013; Gao et al., 2012; Hori et al., 2005; Antoniou et al., 2010). Persulfate irradiated by UV was proposed to generate SO4 through reaction 1 in Table 1 with a quantum yield of 1.4 mol Es1 (254 nm) (Mark et al., 1990). In this process, SO4 are produced with activation of persulfate without the production of OH (Mark et al., 1990). However, OH can also be generated when SO4 reacts with water at a rate constant of 1 1 8.3 M s (reaction 4) (Yu et al., 2004), or with OH at a rate constant of 6.5 107 M1 s1 under alkaline conditions (re action 3) (Neta et al., 1988). Thus, SO4 generation may be an attractive AOTs strategy for removing 2-MIB and geosmin. However, to date the degradation efficiency and rates by SO4 have not been reported for 2-MIB and geosmin. The objectives of this study were (1) to investigate the effects of water chemistry and kinetics of the removal of 2-MIB and geosmin by UV (254 nm)/persulfate; (2) to determine the reaction rates between 2-MIB (and geosmin) and SO4 ; (3) and to model the removal of 2-MIB and geosmin in this process with water quality that is likely to be encountered in water treatment.
2.
Experimental section
2.1.
Chemicals
Chemical solutions were prepared with reagent-grade chemicals and ultra-pure water (18.2 MU cm) produced using a MilliQ biocel system. Potassium peroxydisulfate (PDS), benzoic acid (BA), sodium phosphate monobasic monohydrate, sodium phosphate dibasic, potassium hydroxide and N,Ndiethyl-p-phenylenediamine (DPD) were purchased from SigmaeAldrich, USA. 2-MIB and geosmin were supplied by Wako, Japan. Suwannee River natural organic matter (NOM) obtained from International Humic Substances Society was dissolved into ultra-pure water and filtered through a glassfiber membrane of 0.45 mm pore size (Whatman) to make a stock NOM solution. H2O2 solution (35% v/v) bought from Alfa Aesar, USA was standardized by colorimetric method using DPD (Bader et al., 1988). Phosphoric acid (HPLC grade) was obtained from Dima-Tech Inc., and methanol (MeOH) and nhexane of HPLC grade were obtained from Thermo Fisher
225
Scientific Inc. Tert-butyl-alcohol (TBA) was purchased from Tianjin Chemical Reagent Co., Ltd., China. Other reagents were obtained from Sinopharm Chemical Reagent Co., Ltd., China. All of the chemicals were used as received without further purification. Two real water samples (A and B) for simulating the degradation of 2-MIB and geosmin in realistic water matrices were collected from two drinking water productions. The source water of sample A is from a reservoir, while the other one is from a river. Both the two samples which went through coagulation, sedimentation and filtration processes were further filtered through a 0.45 mm glass fiber membrane (Whatman) prior to using. The water quality parameters of the samples are shown in Table S2. 2-MIB or geosmin was spiked to the water samples to prepare odor-containing water.
2.2.
Experimental procedures
A 600 mL sealed cylindrical borosilicate glass reactor equipped with a 6 W low-pressure Hg UV lamp (254 nm, GPH 135T5 L/4, Heraeus Noblelight) was used to conduct photochemical experiments (Fig. S1). The photon flux (I0) entering the solution from the UV source, effective path length and the corresponding average fluency rate were calculated to be 0.757 mE s1, 3.02 cm and 3.82 103 mE s1 cm2 (1.79 mW cm2), respectively, using the methods described in our previous research (Li et al., 2012). Temperature was controlled at 20 ± 1 C using a thermostat (THD-2015, Tianheng, Ningbo, China). Phosphate buffer (2 mM) was used to adjust pH values from 4.0 to 8.0. Samples were withdrawn at predetermined time intervals and quenched using excess sodium sulfite. All experiments were replicated independently at least two times, and the error bars represent the standard deviation among replicates.
2.3.
Analytical methods
Gas chromatography (GC) (model 6890, Agilent, Santa Clara, USA) was employed to measure the concentrations of 2-MIB and geosmin which were extracted by n-hexane. The chromatograph was coupled with a flame ionization detector (FID) and a HP5 column (30 m 0.25 mm, ID 0.32 mm). Other conditions and the method detection limit can be found in SI Text S1 and Table S1. The concentration of total organic carbon (TOC) was measured using a TOC analyzer (Multi N/C 3100, Analytik Jena, Berlin, Germany). Solution pH was measured by a pH meter (UB-7, Denver Instrument, USA). The extinction coefficients of some chemicals and absorbance of experimental solutions were determined by using a Varian Cary 300 spectrometer, USA. The concentration of persulfate was determined by iodometric method (Nagi et al., 1959). The concentrations of carbonate and bicarbonate were calculated from the pKa of carbonic acid (H2CO3) and pH assuming equilibrium. Details of the calculations are shown in SI Text S2. Bromide and chloride were determined by an ion chromatograph (ICS-3000, Dionex, USA). BA was analyzed by HPLC using a binary HPLC pump (model 1525), a dual l absorbance detector (model 2475) which was set at 227 nm, an autosampler (model 717) and a reverse-phase C18 column (5 mm 4.6 mm 150 mm) (all
226
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
Fig. 1 e Degradation of 2-MIB (a) and geosmin (b) in UV/PDS process. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, [PDS]0 was 200 mM for PDS alone, while 10 mM for UV/PDS, I0/V ¼ 1.26 mE s¡1 L¡1.
from Waters, Milford, USA). An eluant water (10.0 mM phosphoric acid) and methanol (50:50, v/v %) at a flow rate of 1.0 mL/min was used in the system. Reaction rate constants of 2-MIB and geosmin with SO4 and OH were measured using a relative rate technique (Aschmann et al., 2011), in which BA whose reaction rate constant (k15 ¼ 1.2 109 M1 s1, k16 ¼ 5.9 109 M1 s1) is reliably known was chosen as a reference compound (see Text S3 for details) (Neta et al., 1988; Buxton et al., 1988). 1 mM TBA was added in the solution to eliminate OH in the UV/PDS process when determining the reaction rate constant of SO4 reacting with 2-MIB or geosmin. Experiments were also conducted in UV/H2O2 process to determine the reaction rate constants of OH reacting with 2-MIB or geosmin.
primarily on SO4 formed from activation of persulfate (reaction 1) and OH (reactions 3 and 4). The concentration of persulfate was assumed to keep stable as it changed insignificantly when the reaction time was less than 900 s (Fig. S2). Then based on the steady-state assumption, the kinetics of target compounds degradation in UV/PDS can be modeled as follows:
d½C ¼ k1;C ½SO4 ss þ k2;C ½OH ss ½C ¼ kso;C ½C dt
where the concentrations of OH and SO4 are taken to be at steady-state with concentrations [OH]ss and [SO4 ]ss, respectively, ‘C’ is target compound which in this case may be either 2-MIB or geosmin, k1,C and k2,C are reaction rate constants of target compound with SO4 and OH, respectively. It is assumed that SO4 is formed from photolysis of persulfate (reaction 1) (Crittenden et al., 1999); OH is formed from the oxidation of water and hydroxyl ions (Neta et al., 1988; Yu et al., 2004). Assuming steady-state, the net formation rates
2.4.
Modeling equations
The model was established based on the hypothesis that the degradation of target compounds in UV/PDS process depends
(1)
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of SO4 and OH are zero and the concentrations of SO4 and OH can be expressed as eqs. (2) and (3), respectively (see Text S4 in SI for details):
½SO4 ss ¼
½OH ss ¼
fI0 f ð1 10A Þ=V P þ k1;C ½C þ ki ½Si k3 ½OH þ k4 ½H2 O þ k6 S2 O2 8
(2)
i
k3 ½OH þ k4 ½H2 O P k2;C ½C þ k7 S2 O2 þ ki' ½Si 8
i
fI0 f ð1 10A Þ=V P k3 ½OH þ k4 ½H2 O þ k6 S2 O2 þ k1;C ½C þ ki ½Si 8
(3)
i
SO4
2-MIB and geosmin were difficult to be degraded by only PDS oxidation and UV irradiation (254 nm), it could be concluded that SO4 and/or OH produced in this system, via reactions 1 and 4 were primarily responsible for the observed degradation. From the incerting small figures in Fig. 1a and b, the pseudo-first-order rate constants for 2-MIB and geosmin degradation by using UV/PDS process were calculated to be 3.16 103 s1 and 4.80 103 s1, respectively.
1
where Si are scavengers of or OH , I0 (E s ) is the photons flux, G is the fraction of light absorbed by PDS, f (mol E s1) is the quantum yield of SO4 , A (cm1) is the absorbance of solution, and V (L) is the volume of solution. The values of kso;C for target compound degradation and the contributions of SO4 and OH under a range of conditions are then calculated by simultaneously solving eqs. (1)e(3). As phosphate radical is a kind of weak oxidant (Crittenden et al., 1999), its contribution to the degradation to 2-MIB and geosmin was ignored.
3.
Results and discussion
3.1. PDS
Removal efficiency of 2-MIB and geosmin using UV/
Fig. 1 shows the degradation of 238 nM (40 mg/L) 2-MIB and 219 nM (40 mg/L) geosmin in the UV/PDS process at pH 7.0 with a 2 mM phosphate buffer. Negligible degradation of 2-MIB and geosmin was observed using PDS alone at a concentration of 200 mM over a period of 1800 s, indicating little to no oxidation of 2-MIB and geosmin by PDS directly. Less than 3% of 2-MIB and 6% of geosmin were degraded, respectively, within 900 s under the UV irradiation alone (I0/V ¼ 1.26 mE s1 L1), which were in accordance with other reports and consistent with their weak UV absorption at 254 nm (Jo et al., 2011; Kutschera et al., 2009). In contrast, UV/PDS process showed high levels of degradation efficiency for both 2-MIB (86.0%) and geosmin (94.5%) at only 10 mM PDS dosages with a 600 s contact time. As
3.2. Reaction rate constants of SO4 reacting with 2-MIB or geosmin Fig. 2 shows a reaction rate constant ratio between 2-MIB (or geosmin) and BA with SO4 k11/k15 ¼ 0.35 ± 0.05 (or k13/ k15 ¼ 0.63 ± 0.05) at pH 7.0 achieved by using relative rate technique. With a reaction rate constant of BA with SO4 k15 ¼ 1.2 109 M1 s1 (Neta et al., 1988), the reaction rate constants k11 ¼ (4.2 ± 0.6) 108 M1 s1 and k13 ¼ (7.6 ± 0.6) 108 M1 s1 for 2-MIB and geosmin, respectively, could be achieved. The ionic strength affects the reac tion rate constants of SO4 with lots of organic pollutants due to the negative charges at high pHs (Rickman and Mezyk, 2010). However, the reaction rate constant of SO4 reacting with either 2-MIB or geosmin was expected to be little influenced by ionic strength due to their alcoholic structure that the hydroxyl group is bonded to the saturated carbon atom and isn't dissociated at pH 7. Using the same method in a UV/ H2O2 process with a rate constant of BA reacting with OH k16 ¼ 5.9 109 M1 s1 (Buxton et al., 1988), we could calculate that the rate constants of 2-MIB and geosmin reacting with OH were (4.3 ± 0.2) 109 M1 s1 and (5.7 ± 0.2) 109 M1 s1, respectively, based on the rate constant ratio shown in Fig. S3. The values were slightly lower than those reported in von Gunten's research (5.09 109 M1 s1 for 2-MIB and 7.8 109 M1 s1 for geosmin) (Peter and von Gunten, 2007).
3.3.
Roles of sulfate radical and hydroxyl radical
MeOH and TBA are usually used to distinguish SO4 and OH based on reactions 32e35. MeOH is considered as an effective quencher for both SO4 (k32 ¼ 1.1 107 M1 s1) and OH (k33 ¼ 9.7 108 M1 s1) (Buxton et al., 1988; Neta et al., 1988).
Fig. 2 e Determination of the reaction rate constants of SO4 reacting with 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 5 mM, [Geosmin]0 ¼ 5 mM, [BA]0 ¼ 5 mM, [PDS]0 ¼ 100 mM, [TBA]0 ¼ 1 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1.
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w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
Fig. 3 e Impacts of co-existing MeOH and TBA on the pseudo-first-order constants of 2-MIB and geosmin. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1. TBA is an effective quencher for OH (k35 ¼ 6 108 M1 s1) (Buxton et al., 1988), but not for SO4 (k34 ¼ 4 105 M1 s1) (Neta et al., 1988). So the difference of degradation rate of target compounds could be used to distinguish the
contribution of SO4 and OH. Co-existence of either MeOH or TBA inhibited the degradation of 2-MIB (or geosmin), while the inhibiting ability of MeOH is stronger than that of TBA (Fig. 3 and Fig. S4). It proved that both SO4 and OH are existing in the experimental condition and played important role in the degradation of 2-MIB and geosmin. Fig. 3 shows the pseudo-first-order rate constants of 2-MIB and geosmin degradation, and the modeling results of the individual and overall contribution of SO4 and OH at different dosages of co-existing MeOH or TBA at pH 7.0 with 2 mM phosphate buffer. The pseudo-first-order rate constants of both 2-MIB and geosmin in the UV/PDS process decreased with increasing the dosages of MeOH or TBA. The modeling results were in good accordance with the experimental data. The modeling results also show that MeOH and TBA mainly suppressed the contribution of OH at low dosages, due to rapid reactions (k33 ¼ 9.7 108 M1 s1, k35 ¼ 6 108 M1 s1). In addition, the presence of MeOH can also inhibit the formation of OH through reaction 4 in Table 1 by consuming SO4 . Initially, the pseudo-first-order rate constants of 2-MIB and geosmin degradation decreased rapidly. With further increasing the dosages of MeOH or TBA, the decreasing speed of the pseudo-first-order rate constants of 2-MIB and geosmin degradation reduced dramatically. The pseudo-first-order rate constant of 2-MIB decreased by 96% (or 94.3%) with increasing the concentration of MeOH (or TBA) from 0 to 1000 mM. While it decreased by 82.7% and 73.3% with adding 100 mM MeOH and TBA, respectively. The pseudo-first-order rate constant of geosmin followed a similar trend to that of 2-MIB. It reduced by 82.7% (or 72.1%) in the presence of 100 mM MeOH (or TBA) and 86.8% (or 76.8%) in the presence of 1000 mM MeOH (or TBA). The contribution of SO4 became the vital reason for the degradation of 2-MIB or geosmin at high dosages of MeOH or TBA. The decreasing speed of the pseudo-first-order rate constants of 2-MIB and geosmin degradation reduced dramatically after the initial fast decreasing stage was due to that the reaction rate con stant of SO4 reacting with MeOH or TBA was much slower than that of OH. When the dosage of MeOH or TBA was zero, the contribution of OH was about 3.5 times and 2.0 times higher than that of SO4 to 2-MIB and geosmin degradation, respectively. The results further suggested that plenty of SO4
Fig. 4 e Impacts of pH on the pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
would react with water to produce OH in the experimental system (reaction 4).
3.4.
Effects of reaction pH
Degradation rate of 2-MIB and geosmin decreased with increasing pH from 4.0 to 8.0 (Fig. S5). Fig. 4 shows the experimental pseudo-first-order rate constants of 2-MIB and geosmin decreased from 5.59 103 s1 and 8.73 103 s1 to 1.89 103 s1 and 3.35 103 s1, respectively, as pH increased from 4.0 to 8.0. In the range of pH 5.0 to 8.0, the pseudo-first-order rate constants of both 2-MIB and geosmin decreased most quickly. The modeling results shown in Fig. 4 were in accordance with the overall trend of experimental observations, though the experimental values of 2-MIB were lower than those predicted from the model. From the results, it is reasonable to conclude that both SO4 and OH play important roles in the experimental system, but that the contribution of OH was much higher than that of SO4 . Though persulfate dissociation can be enhanced by acid through reaction 2, it can only be important in strong acidic conditions and could be eliminated in the pH range of interest here (Kolthoff and Miller, 1951). The quantum efficiency of photodissociation of persulfate ion was identical at different pHs (Yu et al., 2004). So the formation of sulfate radicals would not be affected by pH, which was in accordance with the degradation of persulfate shown in Fig. S6. Fig. 2a and S7a show that the reaction rate constants be tween SO4 and 2-MIB were stable at pH 4.0, 7.0 and 8.0 ((4.2 ± 0.6) 108 M1 s1 z (4.2 ± 0.6) 108 M1 s1 z (4.5 ± 0.6) 108 M1 s1). The reaction rate constants between SO4 and geosmin calculated from Fig. 2b and S7b show the same trend ((7.5 ± 0.3) 108 M1 s1 z (7.6 ± 0.6) 108 M1 s1 z (7.3 ± 0.7) 108 M1 s1). The results suggested that the reaction rate constants of SO4 for both 2-MIB and geosmin were not affected by pH. 2 The distributions of H2PO 4 and HPO4 are affected by pH (reactions 29e31). There were 2 mM phosphate buffer whose concentration was over 8000 times more than the initial
229
concentration of either 2-MIB (238 nM) or geosmin (219 nM) in the solution, which were scavengers in the UV/PDS process based on reactions 25e28. The reaction rate constant between 6 1 1 s ) is higher than that SO4 and HPO2 4 (k26 ¼ 1.2 10 M between SO4 and H2PO4 (k25 < 7.2 104 M1 s1) (Neta et al., 1988). In addition, the reaction rate constant of OH reacting 5 1 1 s ) is also higher than that with HPO2 4 (k28 ¼ 1.5 10 M reacting with H2PO4 (k27 ¼ 2 104 M1 s1) (Buxton et al., 1988). So higher concentration of HPO2 4 can reduce the theoretical steady-state concentrations of SO4 and OH based on eqs. (2) and (3). The distribution of H2PO4 and HPO2 4 shown in Fig. S8 was also in accordance with the variation of the pseudo-firstorder rate constants of 2-MIB and geosmin. The percent of increased dramatically from pH 5 to pH 8, which HPO2 4 increased the consumption of SO4 and OH and decreased the steady-state concentration of both SO4 and OH. So the pseudo-first-order rate constants decreased obviously from pH 5 to pH 8, while kept stable when the pH was less than 5 where nearly none of HPO2 4 existed. The discussions suggested that pH would not affect the degradation of 2-MIB and geosmin directly, but only through 2 affecting the distribution of H2PO 4 and HPO4 in the experimental process. It should be noted that the contribution of diphosphate radical to the degradation of 2-MIB or geosmin was eliminated when applying the steady-state model.
3.5.
Effects of initial persulfate dosages
The degradation rates of 2-MIB and geosmin increased with increasing persulfate dosage from 0 mM to 50 mM (Fig. S9). Fig. 5 shows that the experimental and calculated data of overall pseudo-first-order rate constants of both 2-MIB and geosmin were in accordance well. The experimental values increased with increasing the dosage of persulfate. The formation rate of SO4 increased with increasing the dosage of persulfate (re action 1), then the ideal steady-state concentrations of SO4 and OH were expected to increase based on eqs. (2) and (3). The modeling results of the individual contributions of SO4 and OH to 2-MIB and geosmin degradation were also shown
Fig. 5 e Impacts of PDS dosages on the pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
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Fig. 6 e Impacts of alkalinity on the pseudo-first-order constants of 2-MIB and geosmin. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
in Fig. 5. OH was the main reactive species, whose contribution to 2-MIB or geosmin degradation was higher than that of SO4 at 10e50 mM persulfate dosages all the time.
3.6.
Effects of alkalinity
Different dosages of bicarbonate (0e2 mM) were added in the reaction solutions at pH 7.0 with 2 mM phosphate buffer to study the effects of alkalinity on the degradation of 2-MIB and geosmin (Fig. S10), and to discuss the individual and the sum of contribution of SO4 and OH to either 2-MIB or geosmin degradation (Fig. 6). At the chosen pH, carbonate species include bicarbonate (HCO 3 ) and H2CO3* (about 16.7%) whose rate constants of reacting with OH and SO4 are unknown. 2 While carbonate (CO3 ) could be eliminated as its percentage 2 of all the carbonate species (i.e., H2CO*3, HCO 3 , CO3 ) was very low (about 0.04%). We here treated it as Fang's method that the rate constants of H2CO3* reacting with OH and SO4 were assumed to be similar with those of HCO 3 reacting with OH and SO4 , respectively (Fang et al., 2014), which still needs to 2 be further proved. HCO 3 and CO3 react with OH or SO4 to produce carbonate radical (reactions 19e22 in Table 1) which is a weak oxidant (Buxton et al., 1988; Zuo et al., 1999). But the
oxidation reactions are insignificant and can be ignored for treating most organic pollutants (Crittenden et al., 1999). So the contribution of carbonate radicals in the experimental system was not considered here. Fig. 6 shows the overall pseudo-first-order rate constants of 2-MIB and geosmin degradation decreased from (3.00 ± 0.17) 103 s1 and 3 1 s (4.58 ± 0.25) 103 s1 without HCO 3 to (0.30 ± 0.03) 10 and (0.57 ± 0.05) 103 s1 in the presence of 2 mM HCO 3, respectively. The results show that bicarbonate could reduce the degradation efficiency of both 2-MIB and geosmin effec tively, which was because it could scavenge both SO4 and OH 6 1 1 at relative high reaction rates (1.6 10 M s for SO4 , 8.5 106 M1 s1 for OH) (Buxton et al., 1988; Zuo et al., 1999). Both the steady-state concentrations of OH and SO4 could be reduced by bicarbonate, but the inhibition effect on OH was much stronger than that on SO4 . It is not only because that the reaction rate of OH is higher than that of SO4 , but also due to the fact that the steady-state concentration of OH is based on the steady state concentration of SO4 (eq. (3)). The concentration of 2-MIB and geosmin could be detected as low as a few ng/L in natural water (Pirbazari et al., 1993). While the bicarbonate concentration varied from several dozen to several hundred mg/L (Crittenden et al., 2005), which was
Fig. 7 e Impacts of co-existing NOM on pseudo-first-order constants of 2-MIB (a) and geosmin (b). Conditions: [2MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 10 mM, 2 mM phosphate buffer, pH ¼ 7.0, 20 C, I0/V ¼ 1.26 mE s¡1 L¡1.
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much higher than the concentration of 2-MIB and geosmin. So alkalinity should be an important scavenging factor for 2-MIB and geosmin degradation in water production when using UV/ PDS process.
3.7.
Effects of NOM
Different dosages of NOM (0e3 mg/L) were added in the samples to elucidate their effects on the degradation of 2-MIB and geosmin (Fig. S12), and the individual and the total con tributions of SO4 and OH to either 2-MIB or geosmin degradation at pH 7.0 with 2 mM phosphate buffer (Fig. 7). Fig. 7 shows that the experimental pseudo-first-order rate constants of both 2-MIB and geosmin were decreased with increasing dosages of co-existing NOM, which were in accordance with the modeling overall pseudo-first-order rate constants well. The pseudo-first-order rate constant of OH for both 2-MIB and geosmin decreased sharply at low dosages of NOM, with a much faster decreasing speed than that of SO4 . The results can be explained by a steady-state concentration of OH that decreased much faster than that of SO4 with coexistence of scavengers through eqs. (2) and (3). When the dosage of NOM exceeded 1.0 mg/L, the contribution of OH could be nearly eliminated. The extinction coefficient of the NOM used here was 0.066 mg1 cm1 L (by TOC) at 254 nm, so it could act as an inner filter to reduce the rate of the activation of persulfate to produce SO4 based on reaction 1 in Table 1. The filtering effect of NOM was also incorporated into the modeling. The reaction 1 rate constant of NOM reacting with OH is 3 108 M1 C s (Hoigne, 1998), and the reaction rate constant of NOM reacting 1 (see Text S5 and Fig. S11 with SO4 is about 2.35 107 M1 C s in SI for details). So NOM competed the reactive radicals with target compound (2-MIB or geosmin) based on reactions 17 and 18 to reduce the steady-state concentration of radicals. The reaction rate constant between NOM and OH is over 10 times higher than that between NOM and SO4 , and the steady-state concentration of OH was affected by that of SO4 via eq. (3). So the decreasing speed of pseudo-first-rate con stant for OH is much quicker than that for SO4 . It is reasonable to conclude that the fast decreasing pseudo-first-order rate constants of 2-MIB and geosmin with low dosages of NOM was mainly due to the sharply decreased contribution of OH (Fig. 7). The pseudo-first-order rate constants of 2-MIB and geosmin decreased much more slowly with further increasing the dosage of NOM, which was because that only SO4 that was competed by NOM with a slower rate constant played roles at high dosage of NOM. It should be noted that NOM could be activated by UV light to produce reactive species (such as singlet oxygen) (Latch and McNeill, 2006; Zhang and Hsu-Kim, 2010), whose contribution to 2-MIB (or geosmin) degradation was eliminated in the model.
3.8.
Degradation efficiency in real water
Two different real water samples (A and B) were used to elucidate the degradation efficiencies of 2-MIB and geosmin (Fig. S13) and the application of the kinetic model in real water (Fig. 8). Fig. 8 shows the experimental values were somewhat lower than the calculated values from the kinetic model. The
Fig. 8 e Application of the kinetic model in two different real waters. Conditions: [2-MIB]0 ¼ 238 nM, [Geosmin]0 ¼ 219 nM, [PDS]0 ¼ 100 mM, 20 C, I0/ V ¼ 1.26 mE s¡1 L¡1. experimental pseudo-first-order rate constants of 2-MIB were about 78.4% and 76.9% of the modeling values in Samples A and B, respectively, and those of geosmin were 81.0% and 77.5% of the modeling values, respectively. Fig. S14 shows the impacts of chloride (0e2 mM) on the degradation of 2-MIB and geosmin was negligible. So the effects of halides were negligible to the degradation kinetics of 2-MIB and geosmin as the concentrations in both the two real water samples were less than 0.6 mM (Table S2). The small discrepancy between experimental data and calculated values from the model was likely due to the difference of the reaction rate constant of NOM in the real waters reacting with radicals (SO4 or OH) from that of Suwannee River NOM used in the model. Variability in the properties of NOM is likely to yield variations in their reaction rate constants with the radicals. For example, the range of reaction rate constants between organic matter 1 at room temperature and OH was 1.21e9.37 108 M1 C s (Mckay et al., 2011), which leads to the difference of the scavenging effect of NOM on OH. The variation should be also existed for SO4 . From the modeling results considering NOM and bicarbonate, the reaction rate constant of SO4 reacting with NOM from both the two real water samples was esti1 which was a little higher mated be about 3.7 107 M1 C s than the reaction rate constant of Suwannee River NOM 1 (2.35 107 M1 C s ). Besides, the effects of other water components, such as the coexisting micropollutants, may also contribute to the discrepancy.
4.
Conclusions
Both 2-MIB and geosmin in water can be removed effectively by using UV/persulfate process. The second order rate con stants for 2-MIB and geosmin reacting with SO4 were esti8 M1 s1 and mated to be (4.2 ± 0.6) 10 8 1 1 (7.6 ± 0.6) 10 M s respectively at a pH of 7.0. A model based on steady-state assumption suggested that both SO4 and OH contributed to the degradation of 2-MIB and geosmin. The contributions of the OH to 2-MIB and geosmin
232
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degradation were higher than the contribution from SO4 in ultra-pure water. While the contributions of OH to 2-MIB and geosmin degradation were inhibited much more significantly than those from SO4 by the co-existence of scavengers including bicarbonate and NOM. Increasing dosages of persulfate increased the degradation of 2-MIB and geosmin. pH did not affect the degradation of 2-MIB and geosmin directly. Faster degradations of both 2-MIB and geosmin in acidic condition when using phosphate buffer were achieved due to the different scavenging effects of hydrogen phosphate and dihydrogen phosphate. NOM and bicarbonate are likely to be the main radical scavengers in natural waters when using UV/ persulfate process to control 2-MIB and geosmin as they can consume both SO4 and OH effectively. The process was also effective in removing 2-MIB and geosmin in real water, and the degradation could also be well described by the model developed here.
Acknowledgments This study was supported by the Natural Science Foundation of China (grants 51178134, 51108117, 21307057 and 51108111). It was also supported by China's Fund for National Creative Research Groups (grant 51121062), and the State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology (grant ES201006). The authors greatly thank Dr Yinghong Guan for the discussion and helpful comments on the model proposed here.
Appendix A. Supplementary data Supplementary data related to this article can be found at http://dx.doi.org/10.1016/j.watres.2014.11.029.
references
Agus, E., Lim, M.H., Zhang, L., Sedlak, D.L., 2011. Odorous compounds in municipal wastewater effluent and potable water reuse systems. Environ. Sci. Technol. 45 (21), 9347e9355. Antoniou, M.G., de la Cruz, A.A., Dionysiou, D.D., 2010. Intermediates and reaction pathways from the degradation of Microcystin-LR with sulfate radicals. Environ. Sci. Technol. 44 (19), 7238e7244. Antonopoulou, M., Evgenidou, E., Lambropoulou, D., Konstantinou, I., 2014. A review on advanced oxidation processes for the removal of taste and odor compounds from aqueous media. Water Res. 53 (0), 215e234. Aschmann, S.M., Arey, J., Atkinson, R., 2011. Kinetics and products of the reaction of OH radicals with 3-Methoxy-3methyl-1-tubanol. Environ. Sci. Technol. 45 (16), 6896e6901. Bader, H., Sturzenegger, V., Hoigne, J., 1988. Photometric method for the determination of low concentrations of hydrogen peroxide by the peroxidase catalyzed oxidation of N,Ndiethyl-p-phenylenediamine (DPD). Water Res. 22 (9), 1109e1115. Buxton, G.V., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical-review of rate constants for reactions of hydrated
electrons, hydrogen-atoms and hydroxyl radicals (OH/O) in aqueous solution. J. Phys. Chem. Ref. Data 17 (2), 513e886. Buxton, G.V., Salmon, G.A., Wood, N., 1990. A pulse radiolysis study of the chemistry of oxysulphur radicals in aqueous solution. In: Physico-chemical Behaviour of Atmospheric Pollutants. Springer, Netherlands, pp. 245e250. Crittenden, J.C., Hu, S., Hand, D.W., Green, S.A., 1999. A kinetic model for H2O2/UV process in a completely mixed batch reactor. Water Res. 33 (10), 2315e2328. Crittenden, J.C., Trussell, R.R., Hand, D.W., Howe, K.J., Tchobanoglous, G., 2005. Water Treatment: Principles and Design, second ed. John Wiley & Sons, Inc, Hoboken, New Jersey. Fang, J., Fu, Y., Shang, C., 2014. The roles of reactive species in micropollutant degradation in the UV/free chlorine system. Environ. Sci. Technol. 48 (3), 1859e1868. Furman, O.S., Teel, A.L., Watts, R.J., 2010. Mechanism of base activation of persulfate. Environ. Sci. Technol. 44 (16), 6423e6428. Gao, Y.-Q., Gao, N.-Y., Deng, Y., Yang, Y.-Q., Ma, Y., 2012. Ultraviolet (UV) light-activated persulfate oxidation of sulfamethazine in water. Chem. Eng. J. 195196 (1), 248e253. Guan, Y., Ma, J., Li, X., Fang, J., Chen, L., 2011. Influence of pH on the formation of sulfate and hydroxyl radicals in the UV/ peroxymonosulfate system. Environ. Sci. Technol. 45 (21), 9308e9314. Guan, Y., Ma, J., Ren, Y., Liu, Y., Xiao, J., Lin, L., Zhang, C., 2013. Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals. Water Res. 47 (14), 5431e5438. Hoigne, J., 1998. Chemistry of aqueous ozone, and transformation of pollutants by ozone and advanced oxidation process. In: The Handbook of Environmental Chemistry Quality and Treatment of Drinking Water. Springer, Berlin. Hori, H., Yamamoto, A., Hayakawa, E., Taniyasu, S., Yamashita, N., Kutsuna, S., 2005. Efficient decomposition of environmentally persistent perfluorocarboxylic acids by use of persulfate as a photochemical oxidant. Environ. Sci. Technol. 39 (7), 2383e2388. Jo, C.H., Dietrich, A.M., Tanko, J.M., 2011. Simultaneous degradation of disinfection byproducts and earthy-musty odorants by the UV/H2O2 advanced oxidation process. Water Res. 45 (8), 2507e2516. Klaning, U.K., Sehested, K., Appelman, E.H., 1991. Laser flash photolysis and pulse radiolysis of aqueous solutions of the fluoroxysulfate ion, SO4F. Inorg. Chem. 30 (18), 3582e3584. Koester, K., 2011. Measuring and Modeling Geosmin Removal from Horsetooth Reservoir Water by Powdered Activated Carbon for Selected Contact Times. Colorado State University, Fort Collins, Colorado. Kolthoff, I.M., Miller, I.K., 1951. The chemistry of persulfate. I. The kinetics and mechanism of the decomposition of the persulfate ion in aqueous medium. J. Am. Chem. Soc. 73 (7), 3055e3059. Kutschera, K., Bornick, H., Worch, E., 2009. Photoinitiated oxidation of geosmin and 2-methylisoborneol by irradiation with 254 nm and 185 nm UV light. Water Res. 43 (8), 2224e2232. Latch, D.E., McNeill, K., 2006. Microheterogeneity of singlet oxygen distributions in irradiated humic acid solutions. Science 311 (5768), 1743e1747. Lau, T.K., Chu, W., Graham, N.J.D., 2007. The aqueous degradation of butylated hydroxyanisole by UV/S2O2 8 : study of reaction mechanisms via dimerization and mineralization. Environ. Sci. Technol. 41 (2), 613e619. Li, M., Xue, Q., Zhang, Z., Feng, C., Chen, N., Lei, X., Shen, Z., Sugiura, N., 2010. Removal of geosmin (trans-1,10-dimethyl-
w a t e r r e s e a r c h 6 9 ( 2 0 1 5 ) 2 2 3 e2 3 3
trans-9-decalol) from aqueous solution using an indirect electrochemical method. Electrochim. Acta 55 (23), 6979e6982. Li, X., Ma, J., Liu, G., Fang, J., Yue, S., Guan, Y., Chen, L., Liu, X., 2012. Efficient reductive dechlorination of monochloroacetic acid by sulfite/UV process. Environ. Sci. Technol. 46 (13), 7342e7349. Mark, G., Schuchmann, M.N., Schuchmann, H.-P., von Sonntag, C., 1990. The photolysis of potassium peroxodisulphate in aqueous solution in the presence of tertbutanol: a simple actinometer for 254 nm radiation. J. Photochem. Photobiol. A Chem. 55 (2), 157e168. Mckay, G., Dong, M.M., Kleinman, J.L., Mezyk, S.P., RosarioOrtiz, F.L., 2011. Temperature dependence of the reaction between the hydroxyl radical and organic matter. Environ. Sci. Technol. 45 (16), 6932e6937. Ministry of Health of The People's Republic of China (MOH), Standardization Administration of The People's Republic of China (SAC), 2006. Standards for Drinking Water Quality (GB 5749-2006). Beijing. Mizuno, T., Ohara, S., Nishimura, F., Tsuno, H., 2011. O3/H2O2 process for both removal of odorous algal-derived compounds and control of bromate ion formation. Ozone Sci. Eng. 33 (2), 121e135. Nagi, W., El Asmar, M.F., El Sadr, M.M., 1959. Iodometric method for determination of persulfate. Anal. Chem. 31 (11), 1870e1871. Neta, P., Huie, R.E., Ross, A.B., 1988. Rate constants for reactions of inorganic radicals in aqueous solution. J. Phys. Chem. Ref. Data 17 (3), 1027e1284. Peter, A., von Gunten, U., 2007. Oxidation kinetics of selected taste and odor compounds during ozonation of drinking water. Environ. Sci. Technol. 41 (2), 6896e6901. Pirbazari, M., Ravindran, V., Badriyha, B.N., Craig, S., McGuire, M.J., 1993. GAC adsorber design protocol for the removal of off-flavors. Water Res. 27 (7), 1153e1166. Rickman, K.A., Mezyk, S.P., 2010. Kinetics and mechanisms of sulfate radical oxidation of b-lactam antibiotics in water. Chemosphere 81 (3), 359e365.
233
Song, W., O'Shea, K.E., 2007. Ultrasonically induced degradation of 2-methylisoborneol and geosmin. Water Res. 41 (12), 2672e2678. Srinivasan, R., Sorial, G.A., 2011. Treatment of taste and odor causing compounds 2-methyl isoborneol and geosmin in drinking water: a critical review. J. Environ. Sci. 23 (1), 1e13. Stumm, W., Morgan, J.J., 1996. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, third ed. John Wiely & SONS, INC. Waldemer, R.H., Tratnyek, P.G., Johnson, R.L., Nurmi, J.T., 2007. Oxidation of chlorinated ethenes by heat-activated persulfate: kinetics and products. Environ. Sci. Technol. 41 (3), 1010e1015. Watson, S.B., Ridal, J., Boyer, G.L., 2008. Taste and odour and cyanobacterial toxins: impairment, prediction, and management in the Great Lakes. Can. J. Fish. Aquat. Sci. 65 (8), 1779e1796. Yu, X.-Y., Bao, Z.-C., Barker, J.R., 2004. Free radical reactions involving Cl, Cl 2 , and SO4 in the 248 nm photolysis of aqueous solutions containing S2O2 8 and Cl . J. Phys. Chem. A 18 (2), 295e308. Zhang, T., Zhu, H., Croue, J.-P., 2013. Magnetically separable CuFe2O4 spinel in water: efficiency, stability, and mechanism. Environ. Sci. Technol. 47 (6), 2784e2791. Zhang, T., Hsu-Kim, H., 2010. Photolytic degradation of methylmercury enhanced by binding to natural organic ligands. Nature Geosci. 3 (7), 473e476. Zou, J., Ma, J., Chen, L., Li, X., Guan, Y., Xie, P., Pan, C., 2013. Rapid acceleration of ferrous iron/peroxymonosulfate oxidation of organic pollutants by promoting Fe(III)/Fe(II) cycle with hydroxylamine. Environ. Sci. Technol. 47 (20), 11685e11691. Zuo, Z., Cai, Z., Katsumura, Y., Chitose, N., Muroya, Y., 1999. Reinvestigation of the acid-base equilibrium of the (bi) carbonate radical and pH dependence of its reactivity with inorganic reactants. Radiat. Phys. Chem. 55 (1), 15e23.