Removal of sulfadiazine, sulfamethizole, sulfamethoxazole, and sulfathiazole from aqueous solution by ozonation

Chemosphere 79 (2010) 814–820

Contents lists available at ScienceDirect

Chemosphere journal homepage: www.elsevier.com/locate/chemosphere

Removal of sulfadiazine, sulfamethizole, sulfamethoxazole, and sulfathiazole from aqueous solution by ozonation Temesgen Garoma *, Shyam K. Umamaheshwar, Alison Mumper Department of Civil, Construction, and Environmental Engineering, 5500 Campanile Drive, San Diego State University, San Diego, CA 92182, United States

a r t i c l e

i n f o

Article history: Received 10 December 2009 Received in revised form 24 February 2010 Accepted 26 February 2010 Available online 19 March 2010 Keywords: Sulfonamides Ozonation Oxidation Reaction rates

a b s t r a c t The removal of sulfadiazine, sulfamethizole, sulfamethoxazole, and sulfathiazole from aqueous solution by ozonation was studied. The study was conducted experimentally in a semi-batch reactor under different experimental conditions, i.e., varying influent ozone gas concentration, bicarbonate ion concentration, and pH. The results of the study indicated that ozonation could be used to effectively remove the sulfonamides from water. The sulfonamides exhibited moderate reactivity towards aqueous ozone, kO3 > 2  104 M1 s1 at pH of 2 and 22 °C. The mol of ozone absorbed by the solution per mol of sulfonamides removed varied in the range of 5.5–12.0 with lower ranges representing ozone absorption by the solution at the beginning of the ozonation process whereas higher ratios correspond to >99.9% removal of the target sulfonamides. The removal rate of the sulfonamides improved with bicarbonate ion concentration up to 8 mM but further increase in bicarbonate ion decreased removal efficiency. It was also observed that increasing the pH from 2.0 to 10.0 resulted in enhanced removal of the sulfonamides. Ó 2010 Elsevier Ltd. All rights reserved.

1. Introduction Scientific studies conducted in the US and worldwide have reported widespread occurrence of pharmaceuticals in treated drinking water (Ternes et al., 2002; Benotti et al., 2009), surface water (Focazio et al., 2008), groundwater (Holm et al., 1995), and wastewater treatment plant effluent and sludge (Lindberg et al., 2005). In particular, the detection of pharmaceuticals in drinking water sources (surface water and groundwater) and treated drinking water is causing great concern in the scientific communities and public health agencies. Because pharmaceuticals are designed to have certain physiological and biological effects, and uncontrolled ingestion may pose adverse health effects on humans and animals (Daughton and Ternes, 1999). Pharmaceuticals reach the environment through several pathways. Pharmaceuticals ingested by humans and animals, for therapeutic effects, are partially absorbed by the body, and the rest are excreted in feces, urine or manure (Loffler and Ternes, 2003). The excreted pharmaceutical can be in the original form or as a metabolite (Heberer, 2002). For example, sulfamethizole, one of the focus chemicals for this study, can be excreted by the body at about 80% of the original (Scholar and Pratt, 2000). Besides human and animal excretion, pharmaceuticals are released into the environment through effluent discharge from pharmaceutical manufacturing plants (Larsson et al., 2007), leachate from landfills * Corresponding author. Tel.: +1 619 594 0957; fax: +1 619 594 8078. E-mail address: [email protected] (T. Garoma). 0045-6535/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.chemosphere.2010.02.060

containing medical waste (Holm et al., 1995), disposal of unused or expired pharmaceuticals (Ternes et al., 2002), leachate from animal manure used as fertilizer (Hirsch et al., 1999), treated wastewater used for irrigation (Hirsch et al., 1999), and recycled water used for groundwater recharge (Drewes et al., 2003). Current data available in the literature indicate that conventional treatment methods used in water treatment (coagulation, flocculation, sedimentation, sand filtration, and disinfection with chlorine) and wastewater treatment (primary settling, activated sludge or trickling filter, and secondary settling) are not effective for removal of all pharmaceuticals present in raw water and wastewater (Kim et al., 2007; Vieno et al., 2007). This is because pharmaceuticals differ greatly in structure and in their physical and chemical properties which affect their rate of removal during treatment. In an activated sludge treatment process, the rate of removal of pharmaceuticals present in wastewater varied between 34% for clofibric acid to 78% for naproxen (Stumpf et al., 1999). In a pilotscale drinking water treatment processes which used coagulation, sedimentation, rapid sand filtration, ozonation, granular activated carbon filtration, and UV disinfection, the removal rate of pharmaceuticals present in the raw water varied in the range 16% for ciprofloxacin to 99% for carbamazepine (Vieno et al., 2007). On the basis of data currently available in the literature, adsorption to activated carbon, oxidation by ozonation, and separation by membrane processes using reverse osmosis (RO), microfiltration (MF), or nanofiltration (NF) seem to be the most promising methods for the elimination of some pharmaceuticals (Ternes et al., 2002; Huber et al., 2003; Westerhoff et al., 2005). However, the

T. Garoma et al. / Chemosphere 79 (2010) 814–820

capital and operational and maintenance costs associated with some of these processes (e.g., RO, MF, and NF) make them less attractive. Sulfonamides are synthetic antimicrobials, developed in the early 1900s. They work by providing a bacteriostatic effect on bacteria, first delay the reproduction of bacteria cells, and then prevent the cells from growing by inhibiting the production of folic acid, which is required for cell growth (Craig and Stitzel, 1994). There are eight common sulfonamides currently used: sulfacetamide, sulfadiazine, sulfadoxine, sulfamethizole, sulfamethoxazole, sulfanilamide, sulfasalazine, sulfisoxazole (Cazes, 2003). The current study focuses on the removal of sulfadiazine (SDZ), sulfamethizole (SFZ), sulfamethoxazole (SMZ), and sulfathiazole (STZ) from aqueous solution using ozonation. Some of the physical and chemical properties for the target sulfonamides are presented in Table 1. The major reasons for selection of these sulfonamides are the following: (a) SMZ is the most widely prescribed antibiotics in the US and other developed countries, and hence frequently detected in the environment (Nicolle, 2002), (b) some of the sulfonamides can be excreted by the body at high rates, as high as 30% for SDZ and 80% for SFZ of the administered dose (Scholar and Pratt, 2000), (c) some of the sulfonamides were detected at very high concentration in the environment, as high as 330 lg L1 for SFZ and 1160 lg L1 for SDZ in groundwater downgradient of a landfill that accepts both household and pharmaceuticals manufacturing waste (Holm et al., 1995), and (d) all the sulfonamides were detected in the environment, including drinking water (Benotti et al., 2009), surface water (Focazio et al., 2008), groundwater (Batt et al., 2006), and wastewater treatment plant effluent (Gobel et al., 2007). The detection of the chemicals in treated drinking water and wastewater treatment plant effluent indicates that they are not effectively removed during conventional water and wastewater treatment. Ozonation has shown to be effective for the oxidation of emerging contaminants, including some pharmaceutics (Ternes et al., 2002; Huber et al., 2003; Westerhoff et al., 2005), phthalates (Gromadzka and Swietlik, 2007; Khan and Jung, 2008), bisphenol A (Deborde et al., 2008; Garoma and Matsumoto, 2009), and other organic contaminants. During ozonation, the degradation of the target chemicals can be initiated by direct reaction with aqueous ozone and hydroxyl radical (OH) which is generated as a result of decomposition of aqueous ozone. In pure water, the decomposition of aqueous ozone is initiated by its reaction with hydroxide ion (OH), and this reaction leads to the production of radicals that propagate the decomposition process by chains of radical–radical or radical–solute reactions and produce OH. The decomposition of aqueous ozone is known to occur very slowly at lower pH levels, and therefore, it can be inferred that direct reaction between aqueous ozone and the target compounds is the main mechanism through which they are removed during ozonation at lower pH levels. Depending on the chemical structure of the target compounds, direct oxidation rate constant for aqueous ozone range from less than 1 to about 109 M1 s1 (Hoigne and Bader, 1983). In contrast, the oxidation of most organic pollutants by OH is extremely rapid, with rate constants of 106–1010 M1 s1 (Haag and Yao, 1992). Besides pH, alkalinity is another water quality parameter that could affect the effectiveness of ozonation as a treatment technique for organics. Bicarbonate and carbonate ions, major constituents of natural alkalinity of water, are known scavengers of OH. In addition, carbonate and bicarbonate ions could also promote ozone decomposition. The reaction between OH and carbonate and bicarbonate ions produce carbonate radical (CO 3 ) which in turn reacts ) produced as a result of aqueous with hydroperoxide ion (HO 2 ozone decomposition and results in the generation of OH through series of radical–radical reactions (Behar et al., 1970; Staehelin and Hoigne, 1982; Buhler et al., 1984; Buxton et al., 1988).

815

The main objective of this study is to evaluate the degradation of four sulfonamides, sulfadiazine (SDZ), sulfamethizole (SFZ), sulfamethoxazole (SMZ), and sulfathiazole (STZ), during ozonation under different experimental conditions by varying influent ozone gas concentration, bicarbonate ion concentration, and pH. The study will also investigate the reaction kinetics of the sulfonamides during ozonation and determine the mol of ozone absorbed by the solution per mol of sulfonamide removed. 2. Material and methods 2.1. Experimental approach and conditions The experimental setup for the study (Fig. 1) consists of an oxygen tank, an ozone gas generator, a reaction vessel, and devices for influent and effluent ozone gas and aqueous ozone measurement. The reaction vessel has an internal diameter of 13 cm and a height of 30 cm, and it was filled with 2.5 L of solution during each experiment. The reaction vessel was equipped with openings for ozone gas inlet and outlet, aqueous ozone probe, and sampling collection. Two glass diffusers were used to sparge ozone gas into the solution at a constant flow rate of 1.2 L min1. The reactor was operated in a semi-batch mode and the contents were stirred continuously using a magnetic stirrer. The experiments were conducted at room temperature, 22 ± 1 °C. During a typical experimental run, solution of target sulfonamide was prepared in deionized water and then the pH of the solution was adjusted to desired value using dilute solutions (1.0 N) of sulfuric acid (H2SO4) and/or sodium hydroxide (NaOH). The concentration of bicarbonate ion was adjusted by using 1.0 M solution of sodium bicarbonate (NaHCO3). Finally, ozone gas was introduced into the solution and 2.0 mL of aliquot samples were withdrawn periodically and analyzed for residual concentration of sulfonamides. The residual ozone was quenched with sodium thiosulfate (Na2S2O3) at 0.5 mM to stop further reaction after sample withdrawal. The degradation of the sulfonamides during ozonation was individually investigated under different experimental conditions (Table 2) by varying influent ozone gas concentration, bicarbonate ion concentration, and pH in the ranges of 1.0–3.2 mg L1, 2.0– 20.0 mM, and 2.0–10.0, respectively. All solutions initially contained about 1000 lg L1 of the target chemical. The concentration of the sulfonamides used was high compared to those found in natural water, but it was chosen to facilitate the investigation of the kinetics of the chemicals reaction during the ozonation process. In the current study, the following operational parameters were used as base values: influent ozone gas concentration of 2.3 mg L1, bicarbonate ion concentration of 2.0 mM, and pH value of 7.0. For all experimental conditions, the influent ozone gas flow rate was kept constant at 1.2 L min1. In addition, a control experiment (exp. run 10) was conducted to check whether the chemicals volatilize from the aqueous phase. In the control experiment, oxygen gas was continuously bubbled into the solution at the same flow rate as for the ozonation experiments. The results indicated that the selected sulfonamides (SMZ and STZ) did not volatilize or strip from the aqueous phase. The base experiment (exp. run 1) was conducted in triplicates and average values are reported in Table 2. 2.2. Analytical methods The influent and effluent ozone gas concentrations were measured using M454 Ozone Analyzer (Teledyne Technologies, Inc., San Diego, CA) calibrated by the Potassium Iodide method (APHA, 1998). The aqueous ozone concentration was determined using

816

Table 1 Chemical and physical properties of selected sulfonamides. SMZ

SDZ

SFZ

STZ

References

CAS No. Molecular formula Molecular weight (g mol1) Water solubility (mg L1) Log Kow (25 °C) (–) Melting point (°C) pKa1, pKa2

723-46-6 C10H11N3O3S 253.3

68-35-9 C10H10N4O2S 250.3

144-82-1 C9H10N4O2S2 270.3

72-14-0 C9H9N3O2S2 255.3

(CRC, 2007) (CRC, 2007) (CRC, 2007)

281a

130c

250a

480b

0.88 171 1.7, 5.6

–0.13 252 2.0, 6.4

0.54 210 1.9, 5.3

0.35 175 2.0, 7.1

(Merck Index, 2001; Wolters and Steffens, 2005; CRC, 2007) (Leo et al., 1971) (CRC, 2007) (Szczepaniak and Szymanski, 2000; Qiang and Adams, 2004)

Chemical structure

O

(Merck Index, 2001; CRC, 2007;)

O

N SMZ

NH

O

S O

NH

N

SFZ

S

NH2

Me

O

NH

SDZ

N a

c

At 25 °C. At 20 °C. At 37 °C.

NH2

Me O

b

S

O N

S O

STZ

NH 2

NH S

S O

NH2

T. Garoma et al. / Chemosphere 79 (2010) 814–820

Property

817

T. Garoma et al. / Chemosphere 79 (2010) 814–820

Alfa Aesar. Other chemicals and reagents used in this study were purchased either from Sigma–Aldrich, Fisher Scientific, or J.T. Baker chemical companies. Ozone gas was generated from zero grade air (Praxair, Inc.) using LG-7 ozone generator (Ozone Engineering, Inc., El Sobrante, CA), and the amount of ozone produced in the oxygen gas was controlled by changing power input to the generator. 3. Results and discussion 3.1. Kinetics of selected sulfonamides with aqueous ozone The reaction of a target sulfonamide with OH and aqueous ozone during ozonation can be expressed by the well established second-order kinetics as follows:

d½M ¼ k OH ½M½ OH  kO3 ½M½O3  dt Fig. 1. Experimental setup.

Q45H Dissolved Ozone Analyzer (Analytical Technology, Inc., Collegeville, PA) calibrated by the Indigo method. Sulfonamides were analyzed using an Agilent 1100 Mass Detector (MSD) (Palo Alto, CA) equipped with an Atmospheric Pressure Chemical Ionization (APCI) source and operated in a positive-ion mode. Nitrogen was used as a drying and nebulizing gas. ChemStation (V B.03.01) was utilized for both data acquisition and processing. The MSD was attached to an Agilent 1100 High Performance Liquid Chromatography (HPLC) (Palo Alto, CA). Depending on the level of oxidation of the sulfonamides, 1–100 lL aliquot samples were injected onto an Agilent C-18 column (Eclipse Plus C18, 2.1  150 mm, 3.5 lm) and the column temperature was maintained at 25 °C. The separation of analytes was achieved using a flow rate of 0.5 mL min1 and a gradient method with two mobile phases: mobile phase A was 95% H2O, 5% acetonitrile, and 0.1% formic acid and mobile phase B was 95% acetonitrile, 5% H2O, and 0.1% formic acid. During a typical run cycle, the column was equilibrated with 10% mobile phase for 2 min, linearly increased to 50% in 8 min, and again the percentage of B was linearly increased to 95% during the next 5 min. Phase B was held at 95% for the 5 min, then the percentage of phase B was linearly decreased to 50% during the next 5 min, again linearly decreased to 10% in 8 min, and then the initial condition was re-established in 2 min. The detection limit for the method was 0.1 lg L1.

ð1Þ

where k OH and kO3 are the reaction rate constant of the target sulfonamide with OH and aqueous ozone, respectively and [M] is the molar concentration of the chemical. At lower pH levels, the degradation of the sulfonamides primarily occurs by direct reaction with aqueous ozone and Eq. (1) can be simplified into Eqs. (2) and (3).

d½M ¼ kO3 ½M½O3  dt ln

½M ¼ kO3 ½Mo

Z

ð2Þ

½O3 dt

ð3Þ

R The plot of lnð½M=½Mo Þ vs: ½O3 dt (Fig. 2) yields a straight line with a slope of kO3 . The results show that the target sulfonamides have moderate reactivity towards aqueous ozone (kO3  2 to 3  104 M1 s1 at pH of 2 and 22 °C). Huber et al. (2003) reported ozone rate constant of 2.5  106 M1 s1 for sulfamethoxazole at pH of 7 and 22 °C. They also estimated that the reactivity of ozone with other sulfonamides is 105 M1 s1 at pH of 7 and 20 °C. The rate constants for the target sulfonamides obtained in this study are less than the ones estimated by Huber et al. (2003). This is because the reactivity of sulfonamides towards ozone is strongly related to their pKa values; deprotonated species have greater reactivity compared to protonated species (Hoigne and Bader, 1983). In Huber et al. (2003), the study was conducted at pH value of 7 where deprotonated species are predominant while in the current study, conducted at pH of 2, protonated species predominate (refer to Table 1 for pKa values of the sulfonamides).

2.3. Materials 3.2. Effect of influent ozone gas All chemicals and reagents used in the research were analytical grade. SFZ and STZ were purchased from MP Biomedicals, Inc., SMZ was purchased from Sigma–Aldrich, and SDZ was obtained from

The effect of influent ozone gas concentration on the removal of the target sulfonamides from dilute aqueous solution was

Table 2 Experimental conditions.

Base condition Effect of ozone Effect of bicarbonate ion

Effect of pH Control a b

Triplicate run. Average values.

Exp. run

Target chemical

(O3)inf (mg L1)

Bicarbonate (mM)

pH (–)

1a 2 3 4 5 6 7 8 9 10

SDZ, SFZ, SMZ, STZ SDZ, SFZ, SMZ, STZ

1.0b 2.3 3.2 2.3 2.3 2.3 2.3 2.3 2.3 –

2.0b 2.0 2.0 4.0 8.0 16.0 20.0 2.0 2.0 2.0

7.0b 7.0 7.0 7.0 7.0 7.0 7.0 2.0 10.0 7.0

SMZ, STZ

SMZ SMZ, STZ

818

T. Garoma et al. / Chemosphere 79 (2010) 814–820

Ln([SFZ]/[SFZ]o)

0.E+00 0.0

2.E-05

4.E-05

6.E-05

8.E-05

1.E-04

-0.5 y = -2.84E+04x R² = 9.84E-01

-1.0

b Ln([SFZ]/[SFZ]o)

a

-1.5 -2.0 -2.5 -3.0

0.E+00 0.0

2.E-05

4.E-05

-0.5

1.E-04

1.E-04

-1.0 -1.5 -2.0 -2.5 -3.0

2.E-05

4.E-05

-0.5

[O3]dt

6.E-05

8.E-05

1.E-04

d Ln([SFZ]/[SFZ]o)

0.E+00 0.0

Ln([SFZ]/[SFZ]o)

8.E-05

y = -2.17E+04x R² = 9.78E-01

[O3]dt

c

6.E-05

y = -1.90E+04x R² = 9.96E-01

-1.0

-1.5

-2.0

0.E+00 0.0

2.E-05

4.E-05

6.E-05

8.E-05

1.E-04

-0.5 y = -2.15E+04x R² = 9.86E-01

-1.0

-1.5

-2.0

[O3]dt

[O3]dt

R

½M Fig. 2. lnð½M Þ vs. ½O3 dt: (a) sulfamethoxazole, (b) sulfamethizole, (c) sulfathiazole, and (d) sulfadiazine. o

1.000

0

30

60

90

120

150

180

0.100

b

[C]/[C]o

[C]/[C]o

a

0.010

1.000

0

30

60

90

120

150

180

0.100

0.010

150

180

120

150

180

(O3)inf=3.0 mg/L

0.010

Time (Sec) 0

120

(O3)inf=2.2 mg/L

(O3)inf=3.6 mg/L

d

1.000

[C]/[C]o

[C]/[C]o

1.000

90

(O3)inf=1.0 mg/L

(O3)inf=2.3 mg/L

c

60

0.100

(O3)inf=1.3 mg/L

0.001

30

0.100

Time (Sec) 0

30

(O3)inf=1.0 mg/L

(O3)inf=1.1 mg/L

(O3)inf=2.3 mg/L

(O3)inf=2.1 mg/L

(O3)inf=3.3 mg/L

Time (Sec)

0.010

60

90

(O3)inf=3.1 mg/L

Time (Sec)

Fig. 3. Effect of influent ozone gas concentration on the removal: (a) sulfamethoxazole, (b) sulfamethizole, (c) sulfathiazole, and (d) sulfadiazine.

investigated using influent ozone gas concentrations of 1.0, 2.3, and 3.2 mg L1 (exp. runs 1–3). Other operational parameters, including initial target sulfonamides concentration, bicarbonate ion concentration, pH, and ozone gas flow rate, were kept constant at 1000 lg L1, 2.0 mM, 7.0, and 1.2 L min1, respectively. The experimental results are presented in Fig. 3, on a semi-log scale.

The experimental sets at 2.3 mg L1 of influent ozone gas (exp. run 1) were conducted in triplicates. The error bars shown in the figure correspond to ±1 standard deviation from the average values obtained from the triplicate experiments. The experimental results revealed that as the influent ozone gas concentration increased, the removal of the sulfonamides increased. By the end of 120 s of ozon-

819

T. Garoma et al. / Chemosphere 79 (2010) 814–820

ation, greater than 90%, 95%, and 99% removals of SFZ, SDZ and STZ, and SMZ, respectively, were achieved for an influent ozone gas concentration of about 3.2 mg L1. On the hand 65–80% removals were obtained for the same duration at 2.3 mg L1 of influent ozone gas. This is expected because an increase in the influent ozone gas concentration results in an increase in aqueous ozone concentration which either directly reacts with the sulfonamides or decomposes to produce OH which in turn reacts with the sulfonamides.

0

30

60

90

120

150

180

[C]/[C]o

1.000

0.100

0.010 pH=2 pH=7

3.3. Ozone absorbed

pH=10

0.001

The number of mol of ozone gas absorbed by the solution per mol of a target sulfonamide removed during ozonation was estimated using the following expression:

R R ½O3 absorbed ð ðO3 Þinf dt  ðO3 Þeff dtÞQ g ¼ ½Mremoved ð½Mo  ½MÞV

ð4Þ

where ½O3 absorbed is the number mol of ozone gas absorbed by the solution, ½Mremoved represents the number of mol of the target sulR R fonamide removed, ðO3 Þinf dt and ðO3 Þeff dt are the area under the influent and effluent ozone gas curve plotted as function of time, respectively, Qg is ozone gas flow rate in L s1, ½Mo is initial concentration of sulfonamide in mol L1, and V is volume of reactor in L. On the basis of Eq. (4), the mol of ozone absorbed by the solution per mol of sulfonamides removed varied in the ranges of 5.6–10.0, 5.8–11.5, 7.7–12.0, and 5.5–10.7 for SDZ, SFZ, SMZ, and STZ, respectively. The lower ranges represent the ozone absorption by the solution at the beginning of the ozonation process whereas the higher ranges correspond to ozone absorption for complete removal of the target sulfonamides. This is because early during the ozonation process ozone primarily reacts with the target sulfonamides; however, at extended reaction time ozone could be also consumed by target sulfonamides’ reaction-intermediates resulting in higher molar stoichiometric ratio. 3.4. Effect of bicarbonate ion The effect of bicarbonate ion on the rate of removal of sulfamethoxazole (SMZ) and sulfathiazole (STZ) during ozonation was investigated by varying bicarbonate ion concentration in the range of 2.0–20.0 mM (122–1220 ppm). The experimental results are presented in Fig. 4, on semi-log scale. The removal of SMZ and STZ increased with an increase in bicarbonate concentration up to 8 mM and then decreased. As pointed out in the introduction, bicarbonate and carbonate ions are known promoters of ozone decomposition and scavengers of OH. Carbonate and bicarbonate ions react with OH and produces carbonate radical (CO 3 ) which

1.000

0

30

60

90

0.100

[HCO3-]=2mM

0.010

0.001

120

150

180

Fig. 5. Effect of pH on the removal sulfamethoxazole.

in turn reacts with radicals produced as a result of aqueous ozone decomposition and results in the generation of OH. In the current study, it can be inferred that the bicarbonate ion primarily served as promoter of ozone decomposition up to 8 mM dose and further increase in dose resulted in the scavenging of OH, hence reducing the removal rate of SMZ and STZ. The reactivity of bicarbonate ion (a predominant species at pH = 7) towards OH, 8.5  106 M1 s1 (Buxton et al., 1988), is significantly lower than the reactivity of SMZ and STZ with OH, 5.8  109 and 7.1  109 M1 s1, respectively (Ikehata et al., 2006), however, the concentration of SMZ and STZ are very small compared to that of bicarbonate ion resulting in k OH ½M << k OH ½HCO 3  as the bicarbonate ion concentration increases beyond 8 mM; where [M] is the concentration of SMZ or STZ. 3.5. Effect of pH Fig. 5 presents the effect of pH on the removal of SMZ during ozonation. At pH value of 2.0, the removal of the sulfonamide was the smallest, about 92% by the end of 3 min of ozonation. Percent SMZ removed increased as the pH increased from 2.0 to 10.0. This is expected, because the deprotonated SMZ species, which is known to have higher reactivity towards ozone compared to protonated species (Hoigne and Bader, 1983), becomes predominant as the pH increases. In addition, at lower pH levels the decomposition of aqueous ozone is known to occur very slowly and the removal of the sulfonamide from dilute aqueous solution is primarily due to its reaction with aqueous ozone. It is well known that the decomposition of aqueous ozone increases with an increase in pH. Thus an increase in pH will result in an increase in  OH and the reactivity the SMZ with OH is significantly higher than with aqueous ozone. However, Lin et al. (2009) reported that the degradation of sulfonamides by ozone decreased with increased pH. The authors

b

[C]/[C]o

[C]/[C]o

a

Time (Sec)

0

30

[HCO3-]=2mM

[HCO3-]=8mM

[HCO3-]=4mM

[HCO3-]=16mM

[HCO3-]=8mM

Time (Sec)

90

0.100

[HCO3-]=4mM

[HCO3-]=20mM

60

1.000

0.010

[HCO3-]=16mM

Time (Sec)

Fig. 4. Effect of bicarbonate ion on the removal: (a) sulfamethoxazole and (b) sulfathiazole.

120

150

180

820

T. Garoma et al. / Chemosphere 79 (2010) 814–820

hypothesized that the decrease in degradation rate was due to a decrease in aqueous ozone concentration resulting from increased ozone decomposition, resulting in the generation of OH. But the authors did not explain the effect of OH on the degradation of the sulfonamides, and they did not appear to use OH scavenger to rule out the contribution of OH. 4. Summary It was observed that ozonation was effective in removing sulfadiazine, sulfamethizole, sulfamethoxazole, and sulfathiazole from aqueous and could lead to complete removal of the chemical from contaminated water. In addition, the target sulfonamides exhibited moderate reactivity towards aqueous ozone, kO3 > 2  104 M1 s1 at pH of 2 and 22 °C. The mol of ozone absorbed by the solution per mol of sulfonamides removed varied in the range of 5.5–12.0 with lower ranges representing ozone absorption by the solution at the beginning of the ozonation process whereas higher ratios correspond to >99.9% removal of the target chemicals. The removal rate of the sulfonamides improved with bicarbonate ion concentration up to 8 mM but further increase in bicarbonate ion decreased removal efficiency. Increasing the pH from 2.0 to 10.0 resulted in enhanced removal of the sulfonamides. The results of this research showed that water containing the target sulfonamides can be treated by ozonation and complete removal could be achieved, if desired. Acknowledgements This research was partially supported by the San Diego State University Division of Research Affairs. References American Public Health Association (APHA), 1998. In: Clescerl, L., Greenberg, A., Eaton, A. (Eds.), Standard Methods for Examination of Water and Wastewater. Washington, DC. Batt, A.L., Snow, D.D., Aga, D.S., 2006. Occurrence of sulfonamide antimicrobials in private water wells in Washington County, Idaho, USA. Chemosphere 64, 1963– 1971. Behar, D., Czapski, G., Duchovny, I., 1970. Carbonate radical in flash photolysis and pulse radiolysis of aqueous carbonate solutions. J. Phys. Chem. 74, 2206–2210. Benotti, M.J., Trenholm, R.A., Vanderford, B.J., Holady, J.C., Stanford, B.D., Snyder, S.A., 2009. Pharmaceuticals and endocrine disrupting compounds in US drinking water. Environ. Sci. Technol. 43, 597–603. Buhler, R.E., Staehelin, J., Hoigne, J., 1984. Ozone decomposition in water studied by pulse-radiolysis 1. HO2/O2 and HO3/O3 as intermediates. J. Phys. Chem. 88, 2560–2564. 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, 513–886. Cazes, J., 2003. The Merck Manual of Diagnosis and Therapy, 17th ed. Craig, C.R., Stitzel, R.E., 1994. Modern Pharmacology, fourth ed. New York, USA. CRC, 2007. CRC Handbook of Chemistry and Physics, 87th ed., David R. Lide (Editorin-Chief). Daughton, C.G., Ternes, T.A., 1999. Pharmaceuticals and personal care products in the environment: agents of subtle change? Environ. Health Perspect. 107 (Suppl. 6), 907–938. Deborde, M., Rabouan, S., Mazellier, P., Duguet, J.P., Legube, B., 2008. Oxidation of bisphenol A by ozone in aqueous solution. Water Res. 42, 4299–4308. Drewes, J.E., Heberer, T., Rauch, T., Reddersen, K., 2003. Fate of pharmaceuticals during ground water recharge. Ground Water Monit. R 23, 64–72. Focazio, M.J., Kolpin, D.W., Barnes, K.K., Furlong, E.T., Meyer, M.T., Zaugg, S.D., Barber, L.B., Thurman, M.E., 2008. A national reconnaissance for pharmaceuticals and other organic wastewater contaminants in the United States – II) Untreated drinking water sources. Sci. Total Environ. 402, 201–216.

Garoma, T., Matsumoto, S., 2009. Ozonation of aqueous solution containing bisphenol A: effect of operational parameters. J. Hazard. Mater. 167, 1185– 1191. Gobel, A., McArdell, C.S., Joss, A., Siegrist, H., Giger, W., 2007. Fate of sulfonamides, macrolides, and trimethoprim in different wastewater treatment technologies. Sci. Total Environ. 372, 361–371. Gromadzka, K., Swietlik, J., 2007. Organic micropollutants degradation in ozoneloaded system with perfluorinated solvent. Water Res. 41, 2572–2580. Haag, W.R., Yao, C.C.D., 1992. Rate constants for reaction of hydroxyl radicals with several drinking-water contaminants. Environ. Sci. Technol. 26, 1005–1013. Heberer, T., 2002. Occurrence, fate, and removal of pharmaceutical residues in the aquatic environment: a review of recent research data. Toxicol. Lett. 131, 5–17. Hirsch, R., Ternes, T., Haberer, K., Kratz, K.-L., 1999. Ocurrence of antibiotics in the aquatic environment. Sci. Total Environ. 225, 109–118. Hoigne, J., Bader, H., 1983. Rate constants of reactions of ozone with organic and inorganic-compounds in water. 2. Dissociating organic-compounds. Water Res. 17, 185–194. Holm, J.V., Rugge, K., Bjerg, P.L., Christensen, T., 1995. Occurence and distribution of pharmaceutical organic compounds in the groundwater downgradient of a landfill (Grindsted, Denmark). Environ. Sci. Technol. 29, 1415–1420. Huber, M.M., Canonica, S., Park, G.Y., von Gunten, U., 2003. Oxidation of pharmaceuticals during ozonation and advanced oxidation processes. Environ. Sci. Technol. 37, 1016–1024. Ikehata, K.N., Naeimeh, J., El-Din, Mohamed.G., 2006. Degradation of aqueous pharmaceuticals by ozonation and advanced oxidation processes: a review. Ozone Sci. Eng. 28, 353–414. Khan, M.H., Jung, J.Y., 2008. Ozonation catalyzed by homogeneous and heterogeneous catalysts for degradation of DEHP in aqueous phase. Chemosphere 72, 690–696. Kim, S.D., Cho, J., Kim, I.S., Vanderford, B.J., Snyder, S.A., 2007. Occurrence and removal of pharmaceuticals and endocrine disruptors in South Korean surface, drinking, and waste waters. Water Res. 41, 1013–1021. Larsson, D.G.J., de Pedro, C., Paxeus, N., 2007. Effluent from drug manufactures contains extremely high levels of pharmaceuticals. J. Hazard. Mater. 148, 751– 755. Leo, A., Hansch, C., Elkins, D., 1971. Partition coefficients and their uses. Chem. Rev. 71, 525–616. Lin, A.Y.C., Lin, C.F., Chiou, J.M., Hong, P.K.A., 2009. O3 and O3/H2O2 treatment of sulfonamide and macrolide antibiotics in wastewater. J. Hazard. Mater. 171, 452–458. Lindberg, R.H., Wennberg, P., Johansson, M., Tysklind, M., Andersson, B., 2005. Screening of human antibiotic substances and determination of weekly mass flows in five sewage treatment plants in Sweden. Environ. Sci. Technol. 39, 3421–3429. Loffler, D., Ternes, T.A., 2003. Determination of acidic pharmaceuticals, antibiotics and ivermectin in river sediment using liquid chromatography–tandem mass spectrometry. J. Chromatogr. A 1021, 133–144. Merck Index, 2001. In: O’Neil, M.J., Smith, A., Heckelman, P.E., Obenchain, J.R., Gallipeau, J.-A.R., D’Arecca, M.A. (Eds.), An Encyclopedia of Chemicals Drugs and Biological. Wiley, New York. Nicolle, L.E., 2002. Urinary tract infection: traditional pharmacologic therapies. Am. J. Med. 113, 35s–44s. Qiang, Z.M., Adams, C., 2004. Potentiometric determination of acid dissociation constants (pK(a)) for human and veterinary antibiotics. Water Res. 38, 2874– 2890. Scholar, E.M., Pratt, W.B., 2000. The Antimicrobial Drugs, second ed. Oxford University Press, Oxford, UK. Staehelin, J., Hoigne, J., 1982. Decomposition of ozone in water – rate of initiation by hydroxide ions and hydrogen-peroxide. Environ. Sci. Technol. 16, 676–681. Stumpf, M., Ternes, T.A., Wilken, R.D., Rodrigues, S.V., Baumann, W., 1999. Polar drug residues in sewage and natural waters in the state of Rio de Janeiro, Brazil. Environ. Sci. Technol. 225, 135–141. Szczepaniak, W., Szymanski, A., 2000. Relationship between hydrophobic properties of amphoteric sulfonamides and their retention in micellar reversed phase liquid chromatography. J. Liquid Chromatogr. Relat. Technol. 23, 1217–1231. Ternes, T.A., Meisenheimer, M., McDowell, D.S.F., Brauch, H.-J., Haist-Gulde, B., Pruess, G., Wilme, U., Zulei-Seibert, N., 2002. Removal of pharmaceuticals during drinking water treatment. Environ. Sci. Technol. 36, 3855–3863. Vieno, N.M., Harkki, H., Tuhkanen, T., Kronberg, L., 2007. Occurrence of pharmaceuticals in river water and their elimination in a pilot-scale drinking water treatment plant. Environ. Sci. Technol. 41, 5077–5084. Westerhoff, P., Yoon, Y., Snyder, S., Wert, E., 2005. Fate of endocrine-disruptor, pharmaceutical, and personal care product chemicals during simulated drinking water treatment processes. Environ. Sci. Technol. 39, 6649–6666. Wolters, A., Steffens, M., 2005. Photodegradation of antibiotics on soil surfaces: laboratory studies on sulfadiazine in an ozone-controlled environment. Environ. Sci. Technol. 39, 6071–6078.