Oxidation of 17β-Estradiol in Water by Slow-Release Permanganate Candles

Abstract

Lakes and streams associated with discharge water from concentrated animal feeding operations (CAFOs) usually contain trace concentrations of endocrine-disrupting compounds (EDCs) that can cause abnormalities in aquatic organisms. Our objectives were to quantify the efficacy of slow-release permanganate (MnO₄⁻) candles to treat 17β-estradiol (E2) in batch experiments and a flow-through system. Kinetic experiments verified that the E2-MnO₄⁻ reaction (at pH ∼6) was first order with respect to MnO₄⁻ and the initial E2 concentration (second-order rate: 59.0 ± 1.07 M⁻¹ s⁻¹). Reactions were pH dependent with the highest rates observed at pH 11, which was above the pKa for E2. We manufactured 0.5 cm (diameter) × 0.64 cm (length) slow-release MnO₄⁻ candles with a 4.6:1 ratio using 23 g KMnO₄ and 5 g paraffin wax to treat E2-spiked environmental matrices collected from a dairy feedlot in a rural area of Nakhon Ratchasima, Thailand. Use of MnO₄⁻ candles completely degraded E2 within 40–60 min, which was faster than the reaction in laboratory water. To test the concept of using candles in a series of discharges from CAFOs, we used a flow-through system where E2-spiked discharge water was pumped over a miniature candle in three discharge events. Results showed that the candle was able to remove ∼92% of E2 by continuously releasing MnO₄⁻. These results provide proof-of-concept that MnO₄⁻ candles have potential for development and could provide a practical approach to treat discharge water before its release into the ecosystem.

Introduction

Endocrine-disrupting chemicals (EDCs) in terrestrial and aquatic environments have gained increasing attention in the past decade (Pickering and Sumpter, 2003; Zhang et al., 2006; Pereira et al., 2011; Lapworth et al., 2012; Xiao-Yan et al., 2015). Recent studies suggested that the dissemination of EDCs into waterways can affect the endocrine function of aquatic organisms by interfering with the metabolic process of natural hormones (Pan et al., 2009). Concentrated animal feeding operations (CAFOs) have been recognized as an important source of manure-borne chemicals such as natural estrogen (Hutchins et al., 2007; Zheng et al., 2013). Dairy waste stored in drainage ditches can literally become contaminated with these compounds before flowing to the receiving watersheds.

Estrogenic chemicals, such as estrone (E1), 17β-estradiol (E2), and estriol (E3), are free estrogens that are found in subsurface water systems at very low concentrations due to their hydrophobicity (Baronti et al., 2000; Kolpin et al., 2002). Among these well-known natural estrogens, E2 is responsible for the development of female characteristics and it showed the highest estrogenic activity based on a yeast estrogenicity screening assay, followed by E1 and E3 (Metcalfe et al., 2001; Rutishauser et al., 2009; Puma et al., 2010). From the various research surveys conducted, E2 concentrations identified in water bodies have ranged from tens to thousands of nanograms per liter (Biswas et al., 2013). Watkinson et al. (2007) found that these compounds as low as nanograms per liter can cause abnormalities of aquatic organisms.

Although E2 has a high octanol–water partitioning coefficient (log K_ow = 3.94; Ying and Kookana, 2005), E2 is not entirely adsorbed by subsurface soil, but rather its level fluctuates with time and may eventually be percolated to groundwater (Prater et al., 2015). For example, E2 concentrations ranging from 6 to 66 ng/L have been reported in groundwater (Peterson et al., 2001). Because E2 in environmental water can contribute to significant deterioration of the ecosystem, the USEPA listed E1, E2, and E3 on the draft Contaminant Candidate List 4 and the EU Water Framework Directive has recently established Environmental Quality Standards for E2 at 0.4 ng/L (Jarošová et al., 2015). As such, a practical approach to remediating E2-contaminated discharge water from the CAFOs is warranted to prevent human and aquatic life from long-term exposure.

Chemical oxidative processes are capable of degrading EDCs and include treatments such as ozone (Esplugas et al., 2007), chlorine (Chamberlain and Adams, 2006), and UV photolysis (Rosenfeldt and Linden, 2004). Permanganate (MnO₄⁻) is an oxidant commonly used in waterworks to control dissolved manganese and iron, taste, odor, color, biological growth, as a disinfectant, and for in situ chemical oxidation. Past research has shown that MnO₄⁻ preferentially attacks compounds with carbon–carbon double bonds, aldehyde groups, or hydroxyl groups, and several phenolic EDCs such as bisphenol F and estrone (Oberle and Schroder, 2000; Jiang et al., 2012; Yang et al., 2013).

Encapsulating oxidants for sustained release is a relatively new technology proposed for subsurface remediation (Ross et al., 2005; Christenson et al., 2012; Rauscher et al., 2012; Lee and Gupta, 2014; Chokejaroenrat et al., 2015). Usually, paraffin is used as the releasing material as it provides a slow, but sustained release over a long period (Christenson et al., 2012; Kambhu et al., 2012; Rauscher et al., 2012; Chokejaroenrat et al., 2015). This slow-release oxidant, in so-called candles, could negate the need for oxidant replenishment and greatly simplify the application process. In theory, these candles could also be placed in discharge conduits that handle wastewater from the CAFOs. Because these candles only release permanganate when submerged, the daily flush of water from the milking parlor and barns would dictate when the MnO₄⁻ is released from the candle and potentially offer a low-maintenance treatment.

Although the impact of estrogenic chemicals has been well documented, to our best knowledge, information is lacking on using slow-release MnO₄⁻ candles to degrade 17β-estradiol from dairy discharge water. Therefore, the objectives of this study were as follows: (1) to determine the kinetics of E2 removal by MnO₄⁻ in aqueous solution and effect of initial pH on the E2-MnO₄⁻ reaction; (2) to determine if slow-release MnO₄⁻ candles could be used to degrade E2 in dairy farm wastewater; and (3) to evaluate the possibility of using slow-release MnO₄⁻ candles in the flow-through system. This was accomplished by quantifying the efficacy of permanganate in a batch-scale experiment and then testing the effectiveness of candles in removing E2 from dairy farm wastewater in a vessel and in a flow-through system.

Materials and Methods

Standards and reagents

Chemicals used in experiments were purchased from a variety of vendors and used as received. These chemicals included the following: potassium permanganate (Ajax Finechem, New South Wales, Australia), ascorbic acid (Ajax Finechem), methanol (Fisher, Loughborough, UK), acetonitrile (Mallinckrodt, St. Louis, NJ), estrone (E1; Sigma-Aldrich, St. Louis, MO), and 17β-estradiol (E2; Sigma-Aldrich) (Table 1). Permanganate stock solutions were prepared in distilled water, while E1 and E2 were prepared in methanol.

Table 1.

Physiochemical Characteristics of Natural Estrogens (Li et al., 2008 )

Natural estrogen	Chemical structure	Molecular formula	MW (g/mol)	Water solubility (mg/L) at 25°C
E2, Estradiol		C₁₈H₂₄O₂	272.39	3.9

Temporal changes in E2 were quantified at 200 nm using high-performance liquid chromatography (HPLC) equipped with a UV detector (Waters, Milford, MA). Peak separations were performed by injecting 20 μL of sample into a Kinetex, 5 μm C18 250 × 4.6 mm, coupled with a guard column (Phenomenex, Torrance, CA). The mobile phase was an isocratic mixture of acetonitrile and H₂O (45:55) at a flow rate of 1 mL/min.

To quantify the natural estrogen (i.e., E1 and E2) concentrations in the dairy discharge water, we used solid-phase extraction cartridges to extract and concentrate the E2 (Pedrouzo et al., 2009). Samples were first filtered with a No. 4 filter paper (Whatman, Buckinghamshire, UK). The Oasis HLB cartridges (Waters) were first washed with 5 mL methanol and 5 mL distilled water. Then, the 200 mL of samples were applied to the cartridges at a flow rate of 12 mL/min and eluted with 5 mL methanol (with 5% acetonitrile) into glass test tubes. Extracted solutions were concentrated using N₂ gas at 50°C until dry. Final extracts were then dissolved with 2 mL acetonitrile and water (45:55). The solutions were then transferred to vials for HPLC analysis as discussed above. Limit of detection (LOD) analysis was also performed for E2 in both laboratory water and dairy farm wastewater by using analytical method applied from Wang et al. (2006).

E2-MnO₄⁻ kinetic experiments

To evaluate the pseudo-first-order rate constant, we performed a batch experiment where the E2 initial concentration was 11.01 μM, and MnO₄⁻ concentrations were ranged from 84.03 to 252.1 μM. To determine the initial reaction rates of E2 and MnO₄⁻, we conducted another set of batch experiments with an initial MnO₄⁻ concentration of 84.03 μM, which was assumed in sufficient excess, to be approximately constant during the E2 oxidation. This experiment received different initial E2 concentrations ranging from 1.84 to 12.85 μM. To avoid complications from subsequent reactions and to save in time, we used an initial rate method, applied from Casado et al. (1986). This method would negate some reaction complications such as the autodecomposition of the reagents or the presence of competitive reactions. All kinetic experiments were performed in 250-mL Erlenmeyer flasks and agitated using an orbital shaker at 150 rpm. Each vial was covered with aluminum foil to prevent MnO₄⁻ photodegradation. Batch experiments were run in triplicate. Solution samples were taken periodically and quenched with 0.1 mL of a sample of 0.1 mM ascorbic acid per milliliter (Broseus et al., 2009). The quenching procedure involved removing 1 mL aliquots from the E2-MnO₄⁻ batch reactors, placing each sample in a 1.5-mL centrifuge tube, adding the quenching agent, and centrifuging at 14,000 rpm for 10 min. The supernatant was then transferred for HPLC and stored at 4°C until analysis.

Controlled-release permanganate candles

Slow-release permanganate candles were manufactured for laboratory experiments and produced in batches by mixing potassium permanganate and paraffin in a 4.6:1 ratio using 23 g KMnO₄ and 5 g paraffin wax. The candle-making procedure has been described in Chokejaroenrat et al. (2015). Candles shaped as cylinders were initially 0.5 cm in diameter and 2.56 cm in length. When control candles were needed, the candles were made in the same mold, using the same proportions, where KCl replaced the MnO₄⁻ (Rauscher et al., 2012) and the procedures used to make the MnO₄⁻ candles were followed.

To quantify the MnO₄⁻ release rates, batch experiments were conducted with different candle lengths (i.e., 1.28, 0.64, 0.32 cm). Because granules of KMnO₄ sink when mixed with melted paraffin, all candles were trimmed of both the top and bottom parts to allow for even distribution of MnO₄⁻ throughout the entire candle. Each candle was placed in a separate Erlenmeyer flask filled with 500 mL distilled water, which was maintained at room temperature. Previous research has established that the MnO₄⁻ candles have two phases in their dissolution pattern: (1) large fluxes of MnO₄⁻ are observed when the candles meet water and (2) continuous fluxes of MnO₄⁻ are observed with time (Christenson et al., 2012). Because the mass of MnO₄⁻ released is linear and diffusion controlled, we determined differences in the release of MnO₄⁻ between freshly prepared MnO₄⁻ candles versus washed (i.e., aged for ∼24 h) candles.

Each treatment mixture of candle was replicated thrice. Immediately before sampling, candles were removed and the solutions were gently mixed by swirling. Aliquots were withdrawn at selected times and the MnO₄⁻ concentration was measured at a wavelength of 525 nm with a Spectronic21 spectrophotometer (Milton Roy, Houston, TX). Samples were diluted when necessary to fall within the linear range of the MnO₄⁻ standard curve. The criteria used to select the appropriate length of candles for successive experiments were that the slow-release MnO₄⁻ should provide a continuous supply of MnO₄⁻ during the discharge activity by yielding enough concentration of MnO₄⁻ to treat both target compounds and natural oxidant demands and maintain its cylindrical shape.

Treatment of E2 using permanganate candles

Evaluation of MnO₄⁻ candle efficiency

Permanganate candles were used to treat E2 in distilled water in a 5 L glass carboy. A sample of ∼20 mg-E2/L was treated with a 0.64-cm aged candle. Candles were adjusted so they were completely immersed in the water. Batch experiments were run in triplicate. Candles were removed and the solutions were gently swirled before each sampling. Samples of 2 mL were periodically collected every 0.5–1 h for the first 12 h and every day for the remainder of the experiment. One milliliter was prepared for the E2 analysis using HPLC, while the other was used for the MnO₄⁻ analysis using UV spectrophotometry as discussed earlier. Samples were diluted when necessary to fall within the linear range of the E2 or MnO₄⁻ standard curve.

Effect of initial pH

A single treatment consisting of 11.01 μM of E2 and a 0.64-cm aged candle was used in this study. Different treatments were started at pH 3, 5, 7, 9, and 11. A recent study has shown that phosphate buffer could interfere with the degradation rate of contaminants (Jiang et al., 2010). Therefore, in our study, an acidic pH was obtained by adding 0.1 N HNO₃, while a basic pH was obtained by adding 0.1 N NaOH. Changes in the E2 concentrations were measured at preselected times over a 180-min period. Samples were quenched with ascorbic acid as described earlier. The experiment was performed in triplicate and conducted at room temperature. All flasks were shaken using an orbital shaker at 150 rpm.

Oxidation in dairy farm wastewater

To take advantage of the merits of the MnO₄⁻ candles for on-site treatment, a laboratory-scale experiment must mimic natural conditions as much as possible. Therefore, we collected dairy farm wastewater in the province of Nakhon Ratchasima, Thailand. There were more than 70 milking cows in the feedlot. The farm was served by a local groundwater pit located adjacent to the milking area. Large volumes of water are used daily to flush the milking parlor, which generates wastewater containing manure. This wastewater receives no treatment and flows naturally to the adjacent lagoon watershed. This lagoon water is later used in a nearby agricultural field. Wastewater samples were collected from three places: (1) the discharge ditch (DC), (2) the nearby waste lagoon (LG), and (3) a groundwater pit (GWP). The pH and temperature were simultaneously monitored on-site using a multiparameter water quality meter (YSI 556) and the data are reported in Table 2. The samples were stored in a 2-L bottle and kept at 4°C until transported to the central laboratory at the Suranaree University of Technology (Nakhon Ratchasima, Thailand) for chemical characteristic analysis using the EPA standard method. The main components of all samples are summarized in Table 2. The estrogenic hormones were also determined using analytical methods as described earlier.

Table 2.

Characteristics of Water from Dairy Farm and Analytical Methods

			Value^a
Items	Abb.	Unit	DC	LG	GWP	Method & Reference
Chemical parameters
Estradiol	E2	μg-E2/L	9.2	4.1	<LOD	HPLC analysis with SPE (Wang et al., 2006)
pH^a	—	—	7.25	6.98	6.42	YSI 556^a
Temperature^a	T	°C	29.5	30.5	18.0	YSI 556^a
Conductivity	—	μS/cm	2,470	1,273	975	YSI 556^a
Ammonia	NH₃	mg-NH₃-N/L	159	0.01	0.13	Titration method (EPA 350.2)
Nitrite nitrogen	NO₂⁻	mg/L	—	0.02	0.01	Colorimetric method (EPA 353.2)
Nitrate nitrogen	NO₃⁻	mg/L	0.4	0.1	2.15	Cadmium reduction method (EPA 353.2)
Total Kjeldahl nitrogen	TKN	mg/L	268	8	1.4	Kjeldahl method (EPA 351.3)
Total phosphorous	TP	mg-P/L	3	6.1	1.79	Ascorbic acid method (APHA 4500-P)
Total nitrogen	TN	mg-N/L	268.4	8.12	3.56	(EPA 353.2)
Total organic carbon	TOC	g/100g	0.06	0.03	0.001	Organic fertilizer analysis (APHA 5310-P)
Total suspended solids	TSS	mg/L	12,950	26,250	527	Dried at 103–105°C (EPA 160.2)
Biological oxidant demand	BOD	mg-BOD/L	1,140	960	3.43	5 Day BOD test (EPA 5210B)
Chemical oxidant demand	COD	mg-COD/L	5,140	1,628	8.91	Close reflux, Colorimetric method (EPA 5220D)

Data monitored on-site using the YSI 556 multiparameter water quality meter.

DC, discharge; GWP, groundwater pit; LG, lagoon.

To quantify the ability of the MnO₄⁻ candles to treat E2 under natural conditions, each aqueous matrix containing 3 mg/L of E2 was filtered and treated with a selected candle in a 200-mL vessel. After candles were placed in water, samples were collected periodically for 180 min. At each sampling, 1 mL of sample was quenched with ascorbic acid as discussed earlier.

Flow-through candle system

To determine if the candle could treat multiple discharge events, fresh batches of E2-spiked filtered discharge water were prepared. The flow-through candle system consisted of the following: (1) an E2 solution reservoir (500 mL); (2) a peristaltic pump model no. BT 100 2J (Hebei, China); (3) Masterflex Viton^® tubing (Coleparmer, Vernon Hills, IL); (4) a modified 130-mL chromatographic column equipped with a stop valve; and (5) a 5-L beaker for collecting treated water (Fig. 1).

FIG. 1.

Schematic diagram of flow-through candle system.

E2 concentrations of 3 mg/L in the reservoir were prepared in filtered discharge water at least 24 h before use. Then, E2-containing discharge water was pumped at a flow rate of 10 mL/min into the column. A stop valve was initially closed to accumulate water above the valve for approximately two-thirds of the column height (Fig. 1). We slowly opened the stop valve so that the effluent flow rate was maintained at the influent flow rate (±0.5 mL/min), thus giving a constant volume in the column at any time. The flow-through experiment started (T = 0 min) when we placed a 0.64-cm fresh candle in the water and simultaneously emptied the effluent reservoir to the influent reservoir. Hence, cycling the 0.5 L of discharge water through the flow-through system took ∼50 min. We assumed that there were three events of discharge water from the milking parlor, so three cycles of 0.5 L were designed for our flow-through experiment. The system was rested for 60 min after each discharge event before the next; therefore, one experiment took 330 min. At each sampling time, we sampled from the effluent reservoir to analyze the E2 and MnO₄⁻ concentrations. To verify that the flow-through system was not causing a decrease in the E2 concentrations, we performed an additional experiment where three events of 0.5 L of E2-spiked solution were passed through the column with an absence of a MnO₄⁻ candle.

Results and Discussion

E2-MnO₄⁻ kinetic experiments

Reactions of E2 (11.01 μM) with an excess of MnO₄⁻ (84.03–252.1 μM) occurred in a distilled water background with pH ∼6. The E2 degradation rate increased when more MnO₄⁻ was added into the solution. The E2 concentration was reduced by 70% in 10 min after treating with 84.03 μM MnO₄⁻ (k = 164.5 × 10⁻³/min), while complete E2 transformation (100%) was achieved when the MnO₄⁻ concentration was more than 210.08 μM (k = 406.7 × 10⁻³/min) (Fig. 2A). It is noteworthy that when the initial concentrations of each reactant were not greatly different (11.01 μM E2 vs. 10.5 μM MnO₄⁻), the E2 degradation rate was considerably lower (i.e., 5.9 × 10⁻³/min; data not shown).

FIG. 2.

(A) Loss of E2 when treated with various concentrations of MnO₄⁻; (B) Plot of pseudo-order rate constants and various initial concentrations of MnO₄⁻; (C) Loss of E2 (initial concentrations ranging from 1.84 to 12.85 μM) when treated with MnO₄⁻ at 84.03 μM; (D) Plot of initial rates and various concentrations of E2 when treated with 84.03 μM MnO₄⁻.

While second-order expressions are commonly used to describe contaminant destruction rates by MnO₄⁻, if MnO₄⁻ is in excess, the reaction can also be described by a pseudo-first-order expression (Huang et al., 1999; Siegrist et al., 2002). Then, the general rate equation can be written as: \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}r = - \frac { 1 } { \alpha } \frac { d [ E2 ] } { dt } = k [ E2 ] ^ { \alpha } [ MnO_4^ - ] ^ { \beta } \tag { 1 } \end{align*} \end{document}

where α is a reaction order with respect to E2, β is a reaction order with respect to MnO₄⁻, r is a reaction rate (i.e., \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$r = k_{obs} [ E2 ] ^{ \alpha}$$ \end{document} ), k is a second-order rate constant, and k_obs is a pseudo-order rate constant. Varying the initial concentration of MnO₄⁻ and measuring k_obs by fitting the results into a pseudo-first-order equation using regression analysis resulted in a wide range of E2 destruction rates (i.e., k_obs = 164.5 × 10⁻³ to 490.5 × 10⁻³/min) (Fig. 2A). The value of β with respect to MnO₄⁻ can be expressed in log–log form as follows: \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}log \ k_{obs} = log \ k + \beta \ log \ [ MnO_4^ - ] _o \tag{2}\end{align*} \end{document}

By plotting log k_obs versus \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$log \ [ MnO_4^ - ] _o$$ \end{document} , the value of β in this regression was 1.02 ± 0.05 (r² = 0.99) and indicates that the reaction was first order with respect to MnO₄⁻ (Fig. 2B). Likewise, by varying the initial concentration of E2 with 84.03 μM MnO₄⁻ and measuring the reaction rate, the value of α with respect to E2 can be determined using a log–log form in the following relationship \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$r = k_{obs} [ E2 ] ^{ \alpha}$$ \end{document} ; therefore, the α can be determined from: \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}log \ r_o = log \ k_{obs} + \alpha \ log \ [ E2 ] _o \tag{3}\end{align*} \end{document}

Initial reaction rates (r_o) were calculated by approximating the tangent to the concentration–time curve (Casado et al., 1986) (Fig. 2C). Figure 2B and D revealed that the initial reaction between E2 and MnO₄⁻ is second order (i.e., α = β = 1) with a rate constant (k″) of 59.0 ± 1.07 M⁻¹ s⁻¹, demonstrating that the rate of E2 degradation is second order overall and first order with respect to MnO₄⁻ and E2. It is noteworthy that, also at pH ∼6, the second-order rate constant observed in this study is about twice of that reported by Jiang et al. (2012), but our rate is still in the range of other phenolic compounds (Waldemer and Tratnyek, 2006). As per the literature review, only a limited number of studies were focused on the rates of EDCs. One of those EDCs was reported to have a second-order rate constant with respect to MnO₄⁻ and contaminants were bisphenol A (28.5 ± 1.1 M⁻¹ s⁻¹ at pH 7, 20°C; Zhang et al., 2013), diclofenac (1.57 ± 0.02 M⁻¹ s⁻¹ at pH 7, 20°C, Cheng et al., 2015), and estrone or E1 (44.5 ± 0.94 M⁻¹ s⁻¹ at pH 5.8, 25°C, Shao et al., 2010). Compared to the destruction rates of various contaminants as reported by Waldemer and Tratnyek (2006), our E2 second-order rate constant can be grouped as chemicals with high reactivity with MnO₄⁻ as the rate is second only to the phenol groups.

Controlled-release permanganate candles

Release of the MnO₄⁻ concentration from three different sizes of candles was carried out in an attempt to select the most suitable candle size for successive experiments. We tested the release rate for 28 days and reported the results using three timelines, 0.2, 7, and 28 days (Fig. 3). The concentration of MnO₄⁻ increased with a longer release time until all MnO₄⁻ had been exhausted from the candles. By comparison, the same type of candle (aged vs. aged, fresh vs. fresh), MnO₄⁻ concentration increased proportionally with increased candle length indicating there was an even distribution of MnO₄⁻ in the entire candle. In the initial phase (0.2 days), all types of fresh candles (i.e., 1.28, 0.64, 0.32 cm in length) yielded approximately a 10-fold higher MnO₄⁻ concentration than the aged candles (Fig. 3). This initial large flux of candle is not surprising as it has been observed by other researchers (Kang et al., 2004; Rauscher et al., 2012). In addition, Christenson et al. (2012) and Kambhu et al. (2012) reported that, with the same candle size, contaminant degradation kinetics mimicked the dissolution patterns of the candles. Although the mass of MnO₄⁻ became limiting sooner when candles were aged, we would obtain a more linear decline in the degradation kinetics of contaminant. Therefore, using fresh candles was possible in the field application.

FIG. 3.

MnO₄⁻ release characteristics from the fresh and aged slow-release candles at 0.2, 7, and 28 days. Bars indicate sample standard deviations (n = 3).

Our test continued to investigate the ability of a MnO₄⁻ candle to treat E2 under varied pH conditions and in the discharge water in later experiments. To verify that the MnO₄⁻ must be valid over a longer period, we monitored the MnO₄⁻ concentration for all candle sizes up to 28 days and found an increase in MnO₄⁻ at a slower release rate. Although the biggest size of MnO₄⁻ candles was able to release a concentration of MnO₄⁻ of up to 800 mg/L (fresh candles) and 500 mg/L (aged candle) (Fig. 3), the need for oxidant replenishment was only designed to comply with three discharge events in 1 day. Therefore, we selected the 0.64-cm candle for the rest of study. We also selected the aged candle, unless otherwise stated, to avoid the unnecessarily large flux that may build up manganese dioxide (MnO₂) on the surface (Christenson et al., 2012). In addition, a lower concentration of MnO₄⁻ was preferred in this study. Several researchers have demonstrated that a MnO₄⁻ concentration of not more than 100 mg/L would be able to remove most EDCs in aqueous solution (Jiang et al., 2012; Zhang et al., 2013; Cheng et al., 2015); hence, the higher concentration of MnO₄⁻ is unnecessary. In addition, we expected the concentration of MnO₄⁻ after dilution in adjacent lagoons not to exceed the typical concentration used in the water treatments (∼2 mg/L; Hu et al., 2009).

Treatment of E2 using permanganate candles

Permanganate candle efficiency

To determine the efficiencies of using permanganate candles, a 20 mg/L solution of E2 was prepared in distilled water at least 24 h before use. We used 0.64-cm aged candle as they can continuously release MnO₄⁻ for 180 mg/L over 7 days (Fig. 3), which should be enough to reduce E2 and other competitive chemicals found in the dairy farm wastewater within three events of discharge water from the milking parlor. The results showed that at 7 days, this MnO₄⁻ candle was able to release MnO₄⁻ at only 70 mg/L (Fig. 4). These results indicate that a larger oxidant candle may be needed to overcome the natural oxidant demand of the wastewater.

FIG. 4.

Temporal changes in MnO₄⁻ and E2 concentration (20 mg/L) when treated with aged MnO₄⁻ candles. Bars indicate sample standard deviations (n = 3).

We further reported E2 removal results as short term (i.e., 0–12 h) and long term (i.e., 1–7 days) (Fig. 4). Using MnO₄⁻ candles slowed the degradation of E2 compared with MnO₄⁻-spiked solution due to the slower input of MnO₄⁻. For example, when the candle released MnO₄⁻ to the solution equivalent to 10 mg/L, 90.5% of E2 was removed in 12 h (Fig. 4), while it took only 20 min to reach the same removal percentage in the MnO₄⁻-spiked solution (Fig. 1A). After this short term, the E2 degradation rate was much slower, and the E2 concentration was lower than our LOD by day 10 (3 μg/L in the distilled water, and 4 μg/L in the dairy farm wastewater). Although the initial concentrations were different from the earlier experiment, these results confirmed that if MnO₄⁻ candles generate and sustain MnO₄⁻ concentrations, degradation of E2 and other competitive compounds would still be initiated within a few hours. It was noted that a large flux of oxidant was diminished by using an aged candle to avoid mixing results from the uneven release of MnO₄⁻ as reported elsewhere (Kang et al., 2004). According to our MnO₄⁻ release experiment, MnO₄⁻ concentrations would be more consistent over the longer period (i.e., 28 days) and would reach ∼270 mg/L (Fig. 3). Several researchers have demonstrated that these concentrations would be sufficient to remove most prominent EDCs in aqueous solution (Shao et al., 2010; Jiang et al., 2012; Wu et al., 2012; Zhang et al., 2013; Cheng et al., 2015).

Effect of initial pH

We further investigated the effect of initial pH on the oxidation rate of E2 after receiving MnO₄⁻ from the candle at five different pHs (3, 5, 7, 9, and 11). The final pH values were not measured. The k_obs values of the E2 degradation were in the following order pH 7 < pH 5 < pH 3 < pH 9 <<< pH 11, ranging between 0.02 and 0.74/min (r² = 0.99), indicating that E2 oxidation by MnO₄⁻ was significantly influenced by the pH (Fig. 5). This is not uncommon for the oxidation of micropollutants by various states of manganese (Mn(VII) to Mn(II)) (Lu et al., 2011). Our finding was similar to Shao et al. (2010) who also reported that the rate of oxidation of estrone (E1), an oxidized form of E2, followed the order of pH 9.4 > pH 7.5 > pH 2.5 > pH 4.6 > pH 6.6. Although their target contaminant was different than ours, our k_obs was four times less than that reported in Shao et al. (2010). This is due to the use of aged slow-release MnO₄⁻ in our experiments, resulting in the aqueous solution receiving much less MnO₄⁻.

FIG. 5.

Degradation of E2 at different pHs (i.e., 3, 5, 7, 9, 11) when treated with aged MnO₄⁻ candles.

In our study, a slight increase in the k_obs value was observed at twofold as the solution became acidified (Fig. 6), which can be explained from the generation of Mn²⁺ that subsequently was oxidized by excess MnO₄⁻ to form manganese dioxide (MnO₂) (Gates-Anderson et al., 2001). In acidic conditions, the redox potential of MnO₄⁻ is more than twice as much (E° = 1.70 V) than in alkaline condition (E° = 0.59 V) (Shao et al., 2010). This, in turn, accelerated the oxidation rate of E2 as MnO₂, which has been reported by several researchers to oxidize a wide range of emerging contaminants, including bisphenol F, lincosamide, and steroid estrogens (Xu et al., 2008; Chen et al., 2010; Lu et al., 2011). Similar results were observed for the E2-MnO₄⁻ reaction, where the reaction rate increased after the pH decreased as a result of excess humic acid (Xu et al., 2008). At pH >9, our results showed that the k_obs value was ∼40-fold higher compared with the environmental pH (i.e., pH 7), demonstrating the high influence of pH on the E2-MnO₄⁻ reaction (Fig. 6). Although Gates–Anderson et al. (2001) stated that highly reactive hydroxyl radicals may be generated and directly oxidize organic contaminants in strongly basic conditions, we believe that in this case, deprotonation is responsible as the pKa of E2 is 10.7 (Lewis and Archer, 1979). Once the solution pH was beyond the pKa of E2, the E2 in the uncharged state rapidly changed. These results from the pH experiment suggested that E2 removal by slow-release MnO₄⁻ candles can be facilitated in both acidic and basic conditions.

FIG. 6.

Plot of k_obs versus initial solution pH when treated with aged MnO₄⁻ candle.

Oxidation in dairy farm wastewater

To expand the results obtained with the E2-MnO₄⁻ in distilled water, we determined whether the MnO₄⁻ candles could treat E2 in dairy farm wastewater that was collected from the dairy feedlot. All samples were analyzed for E2 concentration and the results are reported in Table 2. The SPE recoveries of the E2 from solution matrices were in the following order: discharge water (DC) < lagoon water (LG) < groundwater (GW) < distilled water (DW) (i.e., 57%, 63%, 82%, 89%, respectively). The highest E2 concentration of 9.2 μg/L was from the discharge water sample. We also believe that most of the E2 concentration had already flowed to the receiving water or possibly been adsorbed because of its high K_ow. Yet, the E2 concentration from all sampling areas could be considered insignificant because of our CAFO size. In addition, E2 concentrations in the discharge water were still orders of magnitude lower compared to the following laboratory experiment (0.009 mg/L vs. 3 mg/L).

Treating spiked E2 in filtered environmental matrices using slow-release MnO₄⁻ candles resulted in a slight change in the MnO₄⁻ concentration and an apparent decrease in the first-order degradation compared to the reaction in distilled water following the k_obs order k_DC > k_LG > k_GWP > k_DW (i.e., 0.18, 0.10, 0.03, and 0.02/min, respectively) (Fig. 7). Unlike groundwater and distilled water, the E2 reactions in the discharge water and lagoon water were complete (>95% removal) in 60 min. The faster degradation in dairy farm wastewater was similar to a previous report that highly active manganese intermediates may increase the oxidation activity during the MnO₄⁻ oxidation (Jiang et al., 2009). Although it seemed that total organic constituents such as manure-borne chemicals enhanced the overall removal of E2 from aqueous water and were a contributing factor in the depletion of MnO₄⁻, we believe that the decrease in E2 degradation was also likely a result from the adsorption effect that occurred simultaneously.

FIG. 7.

Temporal changes in E2 concentration in different environmental matrices when treated with aged MnO₄⁻ candles.

The key environmental variables influencing the rate and extent of E2 degradation with MnO₄⁻ are (1) temperature, (2) pH, and (3) natural oxidant demand (Siegrist et al., 2001). Based on our evaluations, these criteria were met at this dairy feedlot for all variables (Table 2). The degradation rates of contaminant with MnO₄⁻ usually decreased as the temperature decreased (Albano et al., 2010; Chokejaroenrat et al., 2011). In our study, the farm was served by a local groundwater pit with a temperature of 18°C, which is lower than water temperatures from all locations (29.5°C). Although the water temperature in this area can be varied from 10 to 35°C dependent upon the season, this temperature discrepancy (18–29.5°C) may not cause a significant impact to the degradation of E2. It should also be noted that the E2 degradation rate could be raised twice as much when the discharge water is acidified to a pH below 5 (Fig. 6). Similar suggestions have been made by other researchers to reduce the pH of wastewater before treatment (Frontistis et al., 2011). In addition, using slow-release MnO₄⁻ candles will be beneficial because conventional use of MnO₄⁻ may be insufficient for samples with high organic carbon such as untreated discharge water.

Flow-through candle system

In the final part of this study, the flow-through candle system was used to investigate the ability of a MnO₄⁻ candle to treat E2-spiked dairy discharge water that was flushed from the milking parlor in three discharge events. A MnO₄⁻ candle was positioned halfway between the water level and the bottom of the chromatographic column (Fig. 1), which mimicked how the candle would be placed in the discharge ditch. This time, we used a fresh candle in the system to ensure that when discharge water with high organic constituents is present, the MnO₄⁻ concentration would be sufficient to treat the target compounds.

Temporal changes in E2 concentrations showed that E2 reduced from the solution when treated with a MnO₄⁻ candle as expected (Fig. 8). In the first discharge event, E2 was reduced by ∼39% after 50 min. The E2 reduction in the first event was much slower than any batch scale experiment due to two reasons: (1) the presence of the organic carbon, and (2) the initial water that passed the stop valve untreated. On the other hand, the E2 concentration in the last event was evident with E2 removal of 92%. The MnO₄⁻ candles generated a MnO₄⁻ concentration of 30 mg/L in the effluent beaker for the first cycle. These higher MnO₄⁻ concentrations resulted in higher and faster E2 degradation. Because of the continuous release of MnO₄⁻ that accumulated in the effluent reservoir, the MnO₄⁻ concentration was 65 mg/L at the end of the experiment (i.e., 330 min), making the reduction of E2 in this last event considerably faster compared to the first two events. The additional experiment, which was run without a MnO₄⁻ candle, showed that E2 losses after passing through the flow-through column were ∼3–8% (data not shown). This finding was similar to Rauscher et al. (2012) who found that loss of PAHs during the flow-through system was merely 5%, while 100% removal was observed during the oxidation from the slow-release MnO₄⁻ candle.

FIG. 8.

Degradation of E2 following three cycles of pumping dairy feedlot discharge water through the flow-through system.

Given that 39% of E2 was removed during the first cycle, the complete oxidation of manure-borne chemicals along with other organic matter likely needs two important factors: (1) larger concentration of MnO₄⁻, and (2) longer contact time. Moreover, given that these chemicals will be dictated by the MnO₄⁻ concentrations generated from dissolution of the candles, MnO₄⁻ concentrations could be increased by changing the diameter, number, or formulation of MnO₄⁻ candles (i.e., KMnO₄ mass) placed in a conduit to treat the incoming discharge. Waldemer and Tratnyek (2006) stated that, to maintain maximum degradation rates, the MnO₄⁻ concentration should remain at least fivefold excess of natural oxidant demand and the target contaminant. While these modifications may need to be optimized for the hydrological and environmental conditions to effectively treat natural E2 or any manure-borne chemicals, the relatively low cost of manufacturing MnO₄⁻ candles combined with their efficacy in removing organic chemicals indicates that this technology offers a potentially low-cost and low-maintenance approach to treating E2-contaminated water.

Conclusion

Slow-release permanganate candles (paraffin-MnO₄⁻) were developed and tested at the laboratory scale to remove 17β-estradiol (E2) from environmental matrices (i.e., discharge water) collected from a dairy farm in the province of Nakhon Ratchasima, Thailand. We studied the E2-MnO₄⁻ kinetic experiments, effect of initial pH, and the efficacy of MnO₄⁻ candles to remove E2 in the flow-through system. We determined differences in the MnO₄⁻ release from freshly prepared MnO₄⁻ candles versus washed (i.e., aged) candles over different time scales. The results showed that the aged MnO₄⁻ candle with a 4.6:1 ratio using 23 g KMnO₄ and 5 g paraffin wax, with a diameter of 0.5 cm and a length of 0.64 cm, was able to effectively treat E2 from all environmental aqueous solutions. Kinetic experiments verified that the E2-MnO₄⁻ reaction was first order with respect to MnO₄⁻ and the initial E2 concentration (second-order rate: 59.0 ± 1.07 M⁻¹ s⁻¹). The reactions were pH dependent and followed the order pH 7 < pH 5 < pH 3 < pH 9 < pH 11. Using the candle in the flow-through system, we confirmed that nearly 92% of E2 removal occurred in the last discharge event (∼330 min). These laboratory-scale results provided proof-of-concept that slow-release MnO₄⁻ candles could be used as an economical approach to treat wastewater from a dairy feedlot before discharge to receiving watersheds.

Footnotes

Acknowledgment

This work was funded by National Science and Technology Development Agency (NSTDA), in a new researcher grants sponsored by Ministry of Science of Technology. Appreciation is expressed to Ms. Supansa Meechuay, the owner of Jumnian Farm for providing access to the dairy feedlot in Nakhon Ratchasima, Thailand. We would also like to thank Ms. Jiraporn Khoksunan, Ms. Siritorn Chaithep, and Mr. Putthipong Limprasert, who participated in this research under the Suranaree University of Technology undergraduate Environmental Engineering project, for assistance in collecting discharge water.

Author Disclosure Statement

No competing financial interests exist.

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Oxidation of 17β-Estradiol in Water by Slow-Release Permanganate Candles

Abstract

Abstract

Introduction

Materials and Methods

Standards and reagents

E2-MnO4− kinetic experiments

Controlled-release permanganate candles

Treatment of E2 using permanganate candles

Evaluation of MnO4− candle efficiency

Effect of initial pH

Oxidation in dairy farm wastewater

Flow-through candle system

Results and Discussion

E2-MnO4− kinetic experiments

Controlled-release permanganate candles

Treatment of E2 using permanganate candles

Permanganate candle efficiency

Effect of initial pH

Oxidation in dairy farm wastewater

Flow-through candle system

Conclusion

Footnotes

Acknowledgment

Author Disclosure Statement

References

E2-MnO₄⁻ kinetic experiments

Evaluation of MnO₄⁻ candle efficiency

E2-MnO₄⁻ kinetic experiments