Preparation of Potassium Ferrate for the Degradation of Tetracycline

Chapter 25

Preparation of Potassium Ferrate for the Degradation of Tetracycline *

Downloaded by STANFORD UNIV GREEN LIBR on March 13, 2013 | http://pubs.acs.org Publication Date: July 25, 2008 | doi: 10.1021/bk-2008-0985.ch025

Shih-fen Yang and Ruey-an Doon

Department of Biomedical Engineering and Environmental Sciences, National Tsing Hua University, Hsinchu, 30013, Taiwan

Tetracycline antibiotics are widely used in veterinary medicine and growth-promoting antibiotics for treatment and/or prevention of infectious disease because of their broadspectrum activity and cost benefit. Tetracycline is rather persistent and a largefraction,which can be up to 75 %, of a single dose can be excreted in non-metabolized form in manures. In addition, potassium ferrate (Fe(VI)) is a powerful oxidant over a wide pH range and can be used as an environmentally friendly chemical in treated and natural waters. Therefore, the ability of ferrate (VI) to oxidize tetracycline in aqueous solution was examined in this study. The stability of Fe(VI) was monitored by the 2,2'-azinobis(3ethylbenzothiazoline-6-sulfonate) (ABTS) method. Results showed that the decomposition of Fe(VI) is highly dependent on pHs. A minimum decomposition rate constant of 1 x 10 s was observed at pH 9.2. A stoichiometry of 1.4:1 was observed for the reaction of tetracycline with Fe(VI) at pH 9.2. In addition, tetracycline was rapidly transformed in the presence of low concentrations of Fe(VI) in the pH rang 8.310.0. The degradation efficiency of tetracycline is affected by both pH and initial Fe(VI) concentration. The degradation of tetracycline by Fe(VI) was also confirmed by ESI-MS and total organic carbon (TOC) analyses. -4

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© 2008 American Chemical Society

In Ferrates; Sharma, V.; ACS Symposium Series; American Chemical Society: Washington, DC, 2008.

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Introduction The widespread occurrence of pharmaceutical and personal care products (PPCPs) has recently received a great concern. PPCPs are released into aquatic environments as a result of industrial, agricultural, and sewage runoff (1-3). PPCPs such as hormones, antimicrobials, andfragranceshave been detected in sewage or wastewater treatment effluents at concentrations in the range of monogram per liter to microgram per liter (1). Recent studies have indicated that several treatment technologies such as coagulation, sedimentation, adsorption, photo catalytic degradation and chemical oxidation can result in the decrease in concentrations and risks of PPCPs in the environments (4-6). In addition, advanced oxidation processes (AOPs) using ozone, hydrogen peroxide and the combination with UV light have shown to effectively degrade PPCP contaminants in waters and wastewaters (5, 7). Tetracyclines (TCs) such as tetracycline, chlortetracycline, oxytetracycline and doxycycline are broad-spectrum antibiotics which show activity against gram-positive and gram-negative bacteria including some anaerobes (8, 9). These compounds are widely used in veterinary medicine and growth-promoting antibiotics for treatment and/or prevention of infectious disease. In general, TCs are added in animal foodstuffs and present in animal waste products. These antibiotics are rather persistent and a largefraction,which can be up to 75 %, of a single dose can be excreted in non-metabolized form in manures, and then entered into the aquatic environments. Biodegradation of PPCPs such as tetracycline under aerobic conditions is limited. Several studies, however, have demonstrated that tetracycline can be removed with respect to the chemical oxidation. Reyes et al. (10) reported a rapid tetracycline degradation in the presence of 0.5 g/L of T i 0 and UV light. About 50% of the initial tetracycline concentration was photodegraded after 10, 20 and 120 min when tetracycline was irradiated with a UV lamp, a solarium device and a UV-A lamp, respectively. In addition, significant mineralization was also obtained when the UV lamp and solarium were used for photocatalysis. Dodd et al. (11) investigated the oxidation of 14 antibacterial molecules by aqueous ozone and found that tetracycline reacted rapidly with ozone over a wide range of pHs at 20 °C. The degradation of tetracycline by ozone followed second-order kinetics and a second-order rate constant of 1.9 x 10 Ivl'V at pH 7.0 was observed (11) These results clearly depict that oxidation processes are suitable technologies for controlling tetracycline in aquatic environments. Ferrate (Fe(VI)) is a powerful oxidant over a wide pH range and can be used as an environmentallyfriendlychemical in treated and natural waters (12, 13). The redox potentials of ferrate(VI) are 2.20 V and 0.72 V under acidic and alkaline conditions (Eq.(l) and (Eq.(2)), respectively (12, 14). During the oxidation processes, ferrate(VI) is reduced to Fe(III) ions or ferrihydrite (Fe(OH) ), resulting in the simultaneous coagulation in a single unit process. In addition, ferrate(VI) is an efficient coagulant for removing toxic contaminants after oxidation. Therefore, ferrate can serve as a dual-function chemical reagent in water and wastewater treatment. 2

6

1

3


406 2

+

3+

Fe0 " + 8H + 3e" -> Fe + 4H 0 4

2

2

Fe0 + 4H 0+ 3e -> Fe(OH) + 50H"


4

2

3

E° = 2.20 V

(1)

E° = 0.72 V

(2)

Potassium ferrate(VI) has been used to oxidize many priority pollutants including reduced sulfur- and nitrogen-containing compounds, chlorophenols, and some toxic inorganics such as arsenic, cyanides and ammonia (11-16). More recently, ferrate(VI) has been used to decompose endocrine disrupting chemicals (17, 18), sulfonamide antimicrobials (19, 20) and microorganisms (14). However, the use of ferrate for the degradation of tetracycline has received less attention. When the potassium ferrate is dissolved in water, oxygen is evolved and ferric oxide is precipitated, making this oxidant unstable in water. The decomposition rate of ferrate is highly dependent on initial ferrate concentration, pH value and temperature (12, 14). Usually the change in aqueous ferrate concentration is determined by a simple and convenient method which monitors the absorbance at 510 nm wavelength (12, 14, 21). However, the inherent limitation of this simple UV-Vis spectroscopic method is that the molar absorbance coefficient is strongly pH dependent. The molar absorption coefficient ( £ f e , 5io nm) of aqueous ferrate(VI) at pH 9.1, the most stable condition of the aqueous ferrate(VI), is 1150 M^cm" , and the molar absorbance coefficient decreases to 520 M^cm" at pH 6.2. In addition, the molar absorption coefficient is too low to measure the Fe(VI) concentration when it is lower than 100 JLIM. errat

1

1

Recently, another method for monitoring the concentration of aqueous ferrate(VI) has been developed. This method is based on the reaction of ferrate(VI) with 2,2'-azinobis(3-ethylbenzothiazoline-6-sulfonate) (ABTS) (22). The colorless ABTS will react with ferrate(VI) to form a green radical ABTS* , which has a maximum absorbance at 415 nm. The change in molar absorbance for the ABTS method at 415 nm is reported to be 3.40 * 10 M ' W per mol L" of added ferrate(VI) (22). This method is highly sensitive and simple for low concentration of aqueous ferrate(VI), which has been used to determine the apparent second-order rate constant for the reaction of Fe(VI) with phenolic endocrine disruptors in aqueous solution (17). However, the use of ABTS method for the determination of the reaction of tetracycline with Fe(VI) has received less attention. The objective of this study was to determine the potential of Fe(VI) for the degradation of tetracycline at various pHs ranging from 7.0-10.0. The potassium ferrate was prepared using wet method by oxidizing Fe(III) salt under strong alkaline conditions. The stability of the prepared potassium ferrate at various pH values was also examined using ABTS method. The stoichiometric relationship between ferrate(VI) and tetracycline was determined. In addition, the oxidative degradation rates of tetracycline by Fe(VI) as a function of pH were determined. Electrospray ionization-mass spectrum (ESI-MS) was also used to identify the degradation of tetracycline by Fe(VI). +

4

1


1

407 (A)

HC

CH


3

3

Figure 1 The chemical structures of (a) tetracycline and (b) ABTS.

Materials and Methods Chemicals Tetracyclines ( C H N 0 , 98 %), N-(2-hydroxyethyl)-piperazine-N'-(2ethanosulfonic acid) (HEPES) (99.5%), and 2,2'-azinobis(3-ethylbenzothiazoline-6-sulfonate) (ABTS) were purchased form Sigma-Aldrich Co. (Milwaukee, WI). Figure 1 shows the chemical structures of tetracycline and ABTS. Ascorbic acid (97 %) was purchased from Mallinckrodt Co. (Phillipsburg, NJ). Solutions were prepared with distilled water that was passed through a Milli-Q water purification system (> 18 MQ cm). All chemicals were of analytical grade and were used as received without further treatment 22

24

2

8

Preparation of potassium ferrate Solid phase of potassium ferrate (K Fe0 ) with a purity up to 90% was synthesized by modifying the method of reaction between OC1" and Fe(NO) in strong basic solutions (12, 14). Sodium hypochlorite (NaOCl") wasfirstreacted 2

4

3


408 with hydrochloric acid (HC1) to produce chlorine gas, and then the chlorine gas was flowed into the vacuumed vessel containing Fe(NO) and KOH to produce potassium ferrate. The resulting purple slurry was filtered and then the precipitate was washed with 1 M aqueous KOH solution. The precipitate was then flushed with n-hexane (four times x 25 mL), n-pentane (four times x 25 mL), and methanol (four times x 25 mL) in sequence (22). The resulting potassium ferrate solid was obtained and stored in the glove box prior to use. The concentration of potassium ferrate was determined using UV-Vis spectroscopy at 510 nm. Aqueous Fe(VI) solution was prepared by dissolving solid samples of K Fe0 in a 5 mM phosphate/1 mM borate buffer at pH 9.1 ± 0.1, which is well-known to be most stable for Fe(VI) solution (12, 14). A molar absorption coefficient of aqueous Fe(VI) eferratoionm of 1150 M^cm" at pH 9.0 was used for the calculation of Fe0 " concentration at 510 nm (21).


3

2

4

1

2

4

Stability of ferrate(VI) at various pH values To understand the stability of potassium ferrate in aqueous solution at various pH values, the prepared solid potassium ferrate was dissolved in various pH buffer solutions. The aqueous ferrate solution was buffered with a 5 mM phosphate buffer in the pH range of 6-8, while 5 mM phosphate/1 mM borate buffer was used in the pH range of 9-11. The stability of ferrate was monitored using ABTS method. An appropriate volume of stock aqueous ferrate(VI) was added into solution containing ABTS at various pHs ranging from 7.0 to 11.0. The final concentration of ferrate(VI) and ABTS were 200 and 100 uM, respectively. After the complete reaction of ABTS with ferrate(VI), a green color was formed (22). The solution was then transferred to the cuvette and the absorbance of green radical (ABTS* ) was monitored at 415 nm by using Hitachi U-3010 UV-Vis spectrophotometer. +

Degradation of tetracycline Reactions between tetracycline and ferrate were conducted using a series of 10 ml serum bottles. The ferrate and tetracycline concentrations were on the order of 10 - 80 uM and 200 uM, respectively. 5 mM phosphate buffer was used to maintain the pH at 7-8, while 5 mM phosphate/1 mM borate buffer was used to control the pH range of 8.5-10.0. All the experiments were performed at room temperature (23 - 25 °C). Samples were periodically withdrawn from the reaction vessels, and immediately treated with 0.5 mL of ascorbic acid (25 mM) or Na S0 (6 mM) to quench the reaction. The 0.45 Jim PTFE filters were used to remove the precipitate, and then the absorption of supernatants were monitored at 359 nm to determine the residual concentration of tetracycline using Hitachi U-3010 UV-Vis spectrophotometer. 2

3


409

Results and Discussion


Sensitivity of ABTS method Ferrate(VI) is a powerful oxidant with the high reduction potential and is unstable in water. The aqueous decomposition rate of ferrate is strongly dependent on the pH and temperature. In this study, the effect of pH on the ferrate stability at room temperature was evaluated by the ABTS method. Lee et al. (22) have depicted that ABTS method is a highly sensitive method for the determination of low concentration of aqueous ferrate over a wide pH range. In this study, the ABTS method was also used to determined decomposition of Fe(VI) and degradation of tetracycline at various pH values. Figure 2 shows the calibration curves of Fe(VI) concentration ranging from 0 - 200 at pH 7.0 and 9.2. A good linearity (r = 0.979, n = 10) of the calibration curve of Fe(VI) at pH 7.0 with the slope of 3.76 x 10 M" cm" was obtained. Similar result was also observed for the calibration curve of Fe(VI) at pH 9.2. The absorbance increased linearly upon increasing Fe(VI) concentration from 0 to 154 uivl and then gradually leveled off to saturation when further increasing the Fe(VI) concentration to 220 u.M. A correlation coefficient of 0.995 with the slope of 3.68 x 10 M" cm" was obtained at pH 9.2. Lee et al. (22) depicted that the reaction of Fe(VI) with ABTS has a stoichiometry of 1:1 when ABTS concentration was in excess (73 \xM). The absorbance coefficient was found to be (3.4 ± 0.05) x 10 M" cm" per mole L" of added Fe(VI) ranging between 0.03 and 35 pM. In this study, we also found that the absorbance of Fe(VI) determined by excess ABTS (100 |iM) at various pH values (7.0 - 9.2) were in the same range. It is noted that the linearity of Fe(VI) used in this study can be up to 150 u.M in the presence of excess ABTS, which shows a wide dynamic range than that determined by Lee et al. (22). In addition, the absorbance (~ 3.7 x 10 M" cm' ) is much higher than that used at 510 nm (1150 M" cm" ), showing that ABTS method is highly sensitive and stable for Fe(VI) determination. 2

4

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1

4

4

1

1

1

1

1

1

1

1

1

1

Stability of Fe(VI) at various pH values The stability of Fe(VI) at various pH values ranging between 7.0 and 11.0 was further examined using ABTS method. Figure 3 shows the concentration profiles of 200 u.M Fe(VI) solutions at various pH values. Results showed that the stability of ferrate(VI) solution is highly dependent on pH values. The ferrate(VI) solution is more stable under alkaline conditions. Only 22 % of the original ferrate(VI) solution remained at pH 7.0 after 540 sec (9 min), while 49 and 93 % of those remained at pH 8.0 and 9.2, respectively. Further increase in



410

Figure 2. The calibration curves of Fe(VI) solution determined by ABTS method at pH 7.0 and 9. 2.

CD

>

03

50

100 150 200 250 300 350 400 450 500 550

Time (sec) Figure 3. The decomposition of ferrate(VI) in aqueous solution at various pH values. The initial ferrate concentration was 200 pM.



411 pH value slightly decreased the stability of ferrate solution, and 63-75 % of the original ferrate remained in solutions when pH was higher than 10. Several sduies have reported that Fe(VI) solution was stable in the pH range 9.0-10.0 (14, 15, 24, 25), which is consistent with the result obtained in this study. The reason for the slight increase in decomposition ratio of Fe(VI) at pH > 10 is not clear. Lee and Gai (25) depicted that the lowest initial reeduction rate of ferrate by water occurred at pH 9.4-9.7, below which decomposition occurred by a reaction that was secondorder with respect to ferrate concentration and approximately first-order above pH 10. In addition, Graham et al. (15) showed that the stability of the ferrate is highly pH dependent and there appeared to be a maximum stability at approximately pH 10. They proposed that the descrease in ferrate stability at pH > 10 may be attributed to the association with a different reduction pathway leading to the formation of anionic ion species instead of Fe(OH) ). The decomposition of ferrate in aqueous solution can be explained by a first-order reaction kinetics with respect to ferrate(VI) concentration (23). Table 1 shows the first-order rate constants (ki) for ferrate decomposition at various pH values. The k\ for ferrate decomposition decreased significantly from 3.9 x 10" s" at pH 7.0 to 1 x 10" s" at pH 9.2. Further increase in pH value slightly increased the decomposition rate of ferrate, and the k for ferrate decomposition was in the rang 5.0-9.7 x 10" s" at pH 10-11. Li et al. (23) reported that the decomposition rate constant of ferrate ranged between (1 - 37) x 10" s" in the pH range of 7.1 - 11.9, and the rate constant has a minimum value between pH 9.2 and 9.4. In this study, a minimal rate constant for ferrate decomposition was found to be 9.2, which shows that the potassium ferrate prepared in this study is most stable under alkaline conditions. It is also noted that the decomposition rate of ferrate(VI) in solution can be neglected during the experiment because the reaction of tetracycline with ferrate can be complete within 3 min. 3(S

3

1

4

1

{

4

1

4

1

Table 1. Thefirst-orderrate constant (ki) for ferrate decomposition at various pH values. 7.0 8.0 9.2 10.0 11.0

3.9 1.8 1.0 5.0 9.7

3

x 10" x 10" x 10' x lO x 10"

3

4

-4

4

0.967 0.988 0.989 0.969 0.988

The ferrate solution was purple initially and gradually changed to a yellowish color when decomposition occurred. Under acidic conditions, ferrate has a high oxidation potential which may lead to a rapid redox reaction with water, and thus results in the instability of ferrate(VI) solution at low pH.


412 +

H F e 0 -» H + HFe0 " 2

4

4

+

HFe0 " -> H + F e 0 4

2_ 4

pK = 3.5 (12, 26)

(3)

pK = 7.23 (12, 13)

(4)

a

a

4 K Fe0 + 10 H 0 -> 4 Fe(OH) + 8 KOH + 30 2

4

2

3

(5)

2

2

On the contrary, the dominant ferrate species is Fe0 ' when the pH is higher than 8.3 (5), which is more stable and persists longer time in solution. It is believed that ferrate could undergo a different reaction pathway at pH > 10, leading to the formation of anionic iron species (Fe(OH)* and Fe(OH) ") instead of the formation of ferrihydrite (Fe(OH) ) (15, 27, 28). Therefore, the pH value of 9.2 was selected to prepare the ferrate(VI) solution for further experiments. 4


3

4

6

3

Stoichiometry The stoichiometric experiments were conducted by fixing the initial concentration of aqueous ferrate(VI) at 70 u.M and varied the initial concentration of aqueous tetracycline rangingfrom0-120 uM in the solution at pH 9.2. The residual concentration of ferrate(VI) after reaction was measured by ABTS method. Figure 4 illustrates the residual ferrate concentration as a

Figure 4. The stoichiometry of ferrate to tetracycline during the reaction of tetracycline with Fe(VI). The initial concentrations of Fe(VI) was 70 pM and the Fe(VI) concentration in solution was determined by ABTS method.


413


function of tetracycline concentration. The ferrate concentration decreased progressively with the increase in tetracycline concentration. When the tetracycline concentration was higher than 50 |iM, the residual ferrate concentration in solution was below to the detection limit. The slope of linear relationship between ferrate and tetracycline in the concentration range of lower than 50 u,M was found to be 1.37. In addition, a yellow color after the reaction of tetracycline with ferrate was clearly observed, suggesting that the final product of Fe(VI) was Fe(III) ions. Therefore, a stoichiometry of 1.4 (Fe(VI):tetracycline) was proposed in this study for the degradation of tetracycline by Fe(VI). 2

1.4 Fe0 ' + tetracycline -> 1.4 Fe(III) + 0 + product(s) 4

2

(6)

Degradation of tetracycline The mixing of aqueous tetracycline and ferrate(VI) leads to a series of solution color changes. The solution starts out purple, rapidly turns yellowish brown, and then slowly faded to colorless. Therefore, the UV-Vis spectra of Fe(VI) and tetracycline solutions were determined first. The UV-Vis spectra of tetracycline showed two peaks at 278 and 359 nm, while a strong absorption peak at about 510 nm was observed for ferrate. Since 278 nm has a strong interference after the reaction of tetracycline with ferrate, the wavelength of 359 nm was used to monitor the change in tetracycline concentration for further experiments. The effect of solution pH on the degradation of tetracycline was first studied. Figure 5 shows the concentration profile of tetracycline as a function of pH value. The degradation efficiency of tetracycline increased upon increasing pH value (Figure 6). The degradation efficiency of tetracycline at pH 7.5 was only 35%, while the oxidative removal efficiencies of tetracycline increased to 53-64% when the pH were higher than 8.3. The low degradation efficiency at pH 7.5 may probably be due to that ferrate(VI) is unstable at pH < 8. This result clearly depicts that pH plays a crucial role in controlling the degradation of tetracycline. The degradation experiments were further undertaken in terms of initial ferrate(VI) concentration where an excess tetracycline (200 uJVI) was used. The effect of ferrate concentration on the degradation of 200 (iM tetracycline at pH 10 was examined. Figure 7 shows the effect of initial Fe(VI) concentration on the degradation of tetracycline. The degradation efficiency increased upon increasing initial ferrate(VI) concentrations ranging from 10 to 75 uM, showing that ferrate(VI) can be used as an environmentallyfriendlyreagent to efficiently oxidative remove tetracycline in aquatic environments. Figure 8 shows the electrospray ionization-mass spectrometry (ESI-MS) spectra of tetracycline and the products after the reaction of tetracycline with


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Preparation of Potassium Ferrate for the Degradation of Tetracycline

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