Peroxymonosulfate Oxidation of

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Rapid Acceleration of Ferrous Iron/Peroxymonosulfate Oxidation of Organic Pollutants by Promoting Fe(III)/Fe(II) Cycle with Hydroxylamine Jing Zou, Jun Ma, Liwei Chen, Xuchun Li, Yinghong Guan, Pengchao Xie, and Chao Pan Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/es4019145 • Publication Date (Web): 13 Sep 2013 Downloaded from http://pubs.acs.org on September 16, 2013

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Environmental Science & Technology

1 2

Rapid Acceleration of Ferrous Iron/Peroxymonosulfate Oxidation of Organic

3

Pollutants by Promoting Fe(III)/Fe(II) Cycle with Hydroxylamine

4

Jing Zou,1 Jun Ma,1,2,* Liwei Chen,3 Xuchun Li,4 Yinghong Guan,5 Pengchao Xie,1 and

5

Chao Pan1

6 7

1

8

Technology, Harbin 150090, P.R. China

9

2

State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of

National Engineering Research Center of Urban Water Resources, Harbin 150090, P.R.

10

China

11

3

12

China

13

4

School of the Environment, Nanjing University, Nanjing 210093, P.R. China

14

5

School of Water Conservancy and Construction, Northeast Agriculture University,

15

Harbin 150030, P.R. China

College of Chemical Engineering, Nanjing Forestry University, Nanjing 210037, P.R.

16 17 18 19

Prepared for Environ. Sci. Technol.

20 21 22

Corresponding author: *

Email: [email protected]

23

Tel: (+86)-0451-8628-2292

Fax: (+86)-0451-8628-3010

24 25 26 27 1

ACS Paragon Plus Environment


1

Abstract

2

The reaction between ferrous iron (Fe(II)) with peroxymonosulfate (PMS) generates

3

reactive oxidants capable of degrading refractory organic contaminants. However, the

4

slow transformation from ferric iron (Fe(III)) back to Fe(II) limits its widespread

5

application. Here, we added hydroxylamine (HA), a common reducing agent, into

6

Fe(II)/PMS process to accelerate the transformation from Fe(III) to Fe(II). With benzoic

7

acid (BA) as probe compound, the addition of HA into Fe(II)/PMS process accelerated

8

the degradation of BA rapidly in the pH range of 2.0-6.0 by accelerating the key reactions,

9

including the redox cycle of Fe(III)/Fe(II) and the generation of reactive oxidants. Both

10

sulfate radicals and hydroxyl radicals were considered as the primary reactive oxidants

11

for the degradation of BA in HA/Fe(II)/PMS process with the experiments of electron

12

spin resonance and alcohols quenching. Moreover, HA was gradually degraded to N2,

13

N2O, NO −2 , and NO3− , while the environmentally friendly gas of N2 was considered as

14

its major end product in the process. The present study might provide a promising idea

15

based on Fe(II)/PMS process for the rapid degradation of refractory organic contaminants

16

in water treatment.

17 18 19 20 21 22 23 24 25 26

2


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1

Introduction

2

Recently, there has been increasing interest in the application of sulfate radicals

3

( SO•− 4 )-based advanced oxidation processes (SR-AOPs) to degrade or mineralize

4

refractory organic contaminants in water treatment and groundwater remediation.1-9

5

SO•− has been known as a strong oxidant for its high standard redox potential 4

6

(2.5-3.1V),10 which is comparable to that of hydroxyl radicals ( •OH , 1.9-2.7V).11

7

Moreover, SO•− is more efficient than •OH to degrade some refractory organic 4

8

contaminants for its selective oxidation capacity.10 As one of the most important way to

9

1-2, 4, generate SO•− 4 , the activation of peroxymonosulfate (PMS) has been widely studied.

10 11

8, 12-18

PMS could be activated by UV,4,

12

transition metals,1-2,

8, 17-18

and heterogeneous

12

catalysts13-16 to generate SO•− 4 . Although UV is an efficient way to activate PMS with

13

quantum yield of 1.04,4 the intrinsic drawback of

14

application. Among the common transition metal ions, Co(II) has been found to be the

15

2 best activator of PMS to generate SO•− 4 . Although heterogeneous cobalt activators have

16

been used to control the concentration of dissolved cobalt ions, the adverse effect of

17

leaching cobalt ions should be carefully considered prior to practical application.15-16

18

Owing to the environmentally friendly nature and the advantages of cost effectiveness

19

and high activity, ferrous iron (Fe(II)) has been commonly selected as the activator of

20

PMS to generate SO•− in practical application.8, 19 4

high cost limit its widespread

21

Unfortunately, Fe(II)/PMS process has some intrinsic drawbacks, including the low

22

degradation efficiency for refractory organic contaminants, the slow transformation from

23

Fe(III) to Fe(II), and the accumulation of ferric oxide sludge, which limit its widespread

24

application.8,

25

researchers have tried to add chelating agents, including citrate, pyrophosphate, and

26

(S,S)-ethylenediamine-N,N'-disuccinic acid trisodium salt (EDDS), to stabilize iron ions

19

Many efforts have been made to alleviate these drawbacks. Previous

3



1

in solution and accelerate the generation of reactive oxidants.19 Heterogeneous iron

2

activators were used to avoid the accumulation of ferric oxide sludge and extend the

3

optimum pH range.13, 15 UV12 and electrochemistry17 were introduced to accelerate the

4

generation of reactive oxidants and the transformation from Fe(III) to Fe(II). However,

5

more efforts are needed to further alleviate the intrinsic drawbacks of Fe(II)/PMS process

6

for its widespread application.

7

However, the aforementioned major drawbacks of Fe(II)/PMS process could be finally

8

attributed to the slow transformation from Fe(III) to Fe(II). As a result, some reducing

9

agents could be added to improve the degradation efficiency of organic contaminants by

10

accelerating the transformation from Fe(III) to Fe(II). Hydroxylamine (HA) is a

11

well-known reducing agent for its role of reducing Fe(III) to Fe(II) in the determination

12

of total dissolved iron with the spectrophotometric method using 1,10-phenanthroline.20-21

13

However, HA was commonly considered as one of the antioxidants of reactive oxidants

14

for its strong reducing property.22 Although HA has been introduced into Fenton process

15

to enhance the generation of •OH by accelerating the redox cycle of Fe(III)/Fe(II),23 to

16

the best of our knowledge, the addition of HA into Fe(II)/PMS process has never been

17

reported. It has been reported that both SO•− and •OH were generated in Fe(II)/PMS 4

18

process.8 Based on the strong ability of reducing Fe(III) to Fe(II)20-21 and low rate

19

10 constants with •OH 11 and SO•− , in this paper, HA was added into Fe(II)/PMS process 4

20

to improve the degradation efficiency of the probe compound by accelerating the key

21

reactions, such as the redox cycle of Fe(III)/Fe(II) and the generation of reactive oxidants.

22

Owing to the relatively simple structure and stable property with common oxidants but

23

10 high rate constants with •OH 11 and SO•− , benzoic acid (BA) was selected as the 4

24

radical probe compound in this study.

25

The aim of this study is to investigate the enhancement of the degradation efficiency of

26

BA in Fe(II)/PMS process with the addition of HA and specifically focus on (i) the effect

27

of HA concentration, (ii) the effect of initial pH, (iii) the role of HA, (iv) the 4


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1

identification of primary reactive oxidants, and (v) the end degradation products of HA in

2

the process.

3

Experimental Section

4

Materials.

Oxone®

(KHSO5·0.5KHSO4·0.5K2SO4,

PMS),

hydroxylamine

5

hydrochloride (HA, 99.999%), benzoic acid (BA, 99.5%), phenol, tetrachlorophenol

6

(4-CP), p-chlorobenzoic acid (p-CBA), p-hydroxybenzoic acid (p-HBA), and

7

5,5-dimethyl-1-pyrolin-N-oxide (DMPO) were of ACS reagent grade and supplied by

8

Sigma-Aldrich, Inc. Nitrobenzene (NB), ferrous sulfate, perchloric acid, sodium

9

hydroxide, sodium sulfate, sodium chloride, and Rhodamine B were of analytical reagent

10

grade and purchased from Sinopharm Chemical Reagent Co., Ltd. Tert-butyl alcohol

11

(TBA) was of guaranteed reagent grade and supplied by Tianjian Chemical Reagent Co.,

12

China. Methanol (supplied by Tedia) and methyl tert-butyl ether (MTBE, purchased from

13

Ficher) were of HPLC grade.

14 15

All of these chemicals were used as received without further purification. All solutions were prepared with ultrapure water produced by a Milli-Q Biocel ultrapure water system.

16

Experimental Procedure. All experiments were performed in 150 mL triangular

17

flasks with a constant stirring rate at 25±0.5 °C. Each 100 mL reaction solution with

18

desired concentrations of BA, ferrous iron, and HA, was prepared with ultrapure water

19

and adjusted to the desired initial pH with perchloric acid and sodium hydroxide. The

20

desired PMS dosage was then added to start the reaction. Samples were withdrawn at

21

predetermined time intervals and quenched with excess pure methanol before analysis.

22

Alcohols quenching experiments with TBA and methanol were performed by adding

23

desired alcohols into the reaction solution before the addition of PMS. Chloride ions and

24

sulfate ions were introduced into the reaction solution in the form of sodium chloride and

25

sodium sulfate, respectively. There were no significant inhibition effect on BA

26

degradation when the concentrations of chloride ions and sulfate ions were not more than 5



1

0.4 mM as shown in the Supporting Information (SI) Text S1, Figure S1, and Figure S2.

2

Electron spin resonance (ESR) experiments were performed using DMPO as

3

spin-trapping agent, whose detailed parameters and procedure are shown in SI Text S2.

4

All experiments were independently repeated at least 2 times, and the average values

5

along with one standard deviation (±SD) were provided in Figures.

6

Analytical Methods. The concentrations of BA, NB, phenol, 4-CP, p-CBA, and

7

p-HBA were measured with a high performance liquid chromatography (Waters 1525),

8

and the detailed parameters are shown in SI Text S3. The pH value was measured with a

9

pH meter (Ultrabasic 7 from Denver Instrument). The concentration of PMS was

10

measured at 556 nm by a modified spectrophotometric method using Rhodamine B.24 The

11

concentration of ferric iron was measured at 300 nm25 with an UV-Vis spectrometer

12

(Varian Carry 300), and the detailed procedure is shown in SI Text S4. HA was first

13

derived to acetoxime by reacting with acetone, and then measured by a gas

14

chromatography (Agilent 7890A) equipped with a flame ionization detector,26 whose

15

detailed procedure is shown in SI Text S5. The concentrations of total dissolved nitrogen

16

(TDN) and total organic carbon (TOC) were measured by a Multi N/C 3100 analyzer

17

(Jena, Germany) equipped with a total nitrogen monitor (Jena TNM-1, Germany). The

18

concentration of Kjeldahl nitrogen was measured with the traditional indophenol

19

method.27 The concentrations of

20

chromatograph (Dionex ICS-3000), and the detailed procedure is shown in SI Text S6.

21

The concentration of dissolved N 2 O was measured with a gas chromatograph (Agilent

22

6890N) equipped with an electron capture detector, a headspace sampler (Agilent G1888),

23

and a GS-CarbonPLOT capillary column (30 m × 0.32 mm × 3.00 µm), based on the

24

method of static headspace gas chromatographer.28

25

Results and Discussion

26

NO −2

and

NO3− were measured by an ion

Degradation Efficiency of BA in HA/Fe(II)/PMS Process. Figure 1 shows the 6


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1

degradation of BA in HA/Fe(II)/PMS process. As can be seen, less than 6% of BA was

2

degraded in 15 minutes in Fe(II)/PMS process. Such low degradation efficiency could be

3

interpreted with the low Fe(II) concentration and the slow transformation from Fe(III) to

4

Fe(II).2 Surprisingly, more than 80% of BA was degraded in 15 minutes in Fe(II)/PMS

5

process with the addition of HA. Meanwhile, it should be noted that only about 6% of BA

6

was degraded in HA/PMS process. As shown in Figure 1 and SI Figure S3, with the same

7

low concentration of transition metals (as low as 10.8 µM), the degradation efficiency of

8

BA in HA/Fe(II)/PMS process was as high as that in Co(II)/PMS process, while the latter

9

process has been considered as one of the most efficient way to activate PMS.2

10

(Insert Figure 1)

11

Effect of HA Concentration on BA Degradation in HA/Fe(II)/PMS Process. To

12

investigate the role of HA, the effect of HA concentration on the degradation of BA in

13

HA/Fe(II)/PMS process was studied. As shown in Figure 2, increased degradation of BA

14

was observed with the increase of HA concentration in the range of 0.0 to 0.4 mM, then

15

the increase in HA concentration resulted in a decrease in BA removal. It should be noted

16

that HA here mainly existed in the form of NH3OH+ at pH 3.0 (about 99.9% in molar

17

ratio as shown in SI Figure S4) with pKa1 = 5.9629. The degradation of BA was actually

18

regulated by NH3OH+ with the addition of HA into Fe(II)/PMS process. Although the

19

increase of NH3OH+ concentration could accelerate the transformation from Fe(III) to

20

Fe(II), a large amount of the generated reactive oxidants could be quenched by excess

21

10 NH3OH+ ( k = 1.5 × 107 M -1s -1 for SO•− and k < 5.0 × 108 M -1s -1 for •OH 11) as the 4

22

concentration of NH3OH+ was high enough. Hence, in order to improve the degradation

23

of probe compounds at the highest extent and reduce the cost, a proper dosage of HA

24

should be selected in the application.

25

(Insert Figure 2)

26

Effect of Initial pH on BA Degradation in HA/Fe(II)/PMS Process. It has been

27

reported that the degradation of 2-chlorobiphenyl was greatly affected by the increases in 7



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1

initial pH in Fe(II)/PMS process due to the changes in Fe(III) speciation, catalyst

2

precipitation, and self-dissociation of PMS through non-radical pathways.8 Although

3

Fe(II) species are readily soluble in a wide pH range (2.0-9.0), Fe(III) ions begin to

4

precipitate in the form of ferric oxyhydroxides above pH 3.0,30 which may redissolve

5

with difficulty and are less reactive to activate PMS. Meanwhile, the distribution ratio of

6

NH3OH+ was controlled by solution pH as shown in SI Figure S4. However, owing to the

7

potential role of quenching reactive oxidants and/or complexing with Fe(III) or Fe(II), no

8

buffering agents were introduced to stabilize solution pH during our experiments. Initial

9

solution pH was adjusted using sodium hydroxide or perchloric acid, and the variation of

10

solution pH is shown in SI Figure S5.

11

As shown in Figure 3, an obvious increasing of BA degradation was observed with the

12

increase of initial pH in the range of 2.0 to 3.0. This could be attributed to the formation

13

of Fe(OH)2 with the increase of initial pH, which has been reported to be more reactive

14

than Fe(II) ions in Fenton process.31 The degradation of BA was slightly affected by the

15

increase of initial pH in the range of 3.0 to 5.5, but it was obviously inhibited by

16

increasing the initial pH from 5.5 to 7.0 in HA/Fe(II)/PMS process. It has been reported

17

that Fe(III) ions begin to precipitate in the form of ferric oxyhydroxides above pH 3.0.30

18

Moreover, NH2OH becomes the dominant existing form of HA as pH increases up to 6.0

19

(SI Figure S4). Hence, most of the generated reactive oxidants would be consumed by

20

NH2OH via side reactions for the high rate constants between NH2OH and reactive

21

oxidants ( k = 9.5 ×109 M -1s -1 for •OH

22

compared with the conventional Fe(II)/PMS process, HA/Fe(II)/PMS process had a much

23

higher degradation efficiency of BA in the pH range of 2.0-7.0 (Figure 3 and SI Figure

24

S6). Therefore, with the increase of initial pH, more HA would exist with the form of

25

NH2OH and Fe(II) would be continuously transformed to Fe(OH)2 and ferric

26

oxyhydroxides, which could be the major causes to the variation of the degradation

27

efficiency of BA with the increase of initial pH in the range of 2.0 to 7.0.

11

, k = 8.5 × 108 M -1s -1 for SO•− 4

8


10

). Besides,

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1

(Insert Figure 3)

2

Role of HA in HA/Fe(II)/PMS Process. Based on the aforementioned data and

3

analysis, it could be inferred that the addition of HA into Fe(II)/PMS process might

4

greatly accelerate the transformation from Fe(III) to Fe(II) and the generation of reactive

5

radicals. To further explore the role of HA, the variations of the concentrations of Fe(III),

6

PMS, and HA in HA/Fe(II)/PMS process and Fe(II)/PMS process were both measured

7

and shown in Figure 4.

8

As shown in Figure 4, nearly all Fe(II) (more than 99%) was promptly oxidized to

9

Fe(III) within 10 seconds in Fe(II)/PMS process when the dosage of Fe(II) was 10.8 µM.

10

This phenomenon could be interpreted by the fast reactions between Fe(II) and PMS

11

(equations 1 and 2)32 and the slow transformation from Fe(III) to Fe(II). In

12

HA/Fe(II)/PMS process, nearly 97% of Fe(II) (10.5 µM) was promptly transformed to

13

Fe(III) within 10 seconds, and then Fe(III) concentration kept relatively constant in 12.5

14

minutes but then decreased slowly. It should be noted that more than 90% of 10.8 µM

15

Fe(III) was transformed to Fe(II) within 10 seconds when mixed with 0.40 mM HA (data

16

not shown). Therefore, the steady Fe(III) concentration in HA/Fe(II)/PMS process should

17

be in dynamic equilibrium and would be changed with the variations of the residual

18

concentrations of PMS and HA. It means that the recovery of Fe(II) in Fe(II)/PMS

19

process was strongly accelerated with the addition of HA. Meanwhile, the variation of

20

Fe(III) concentration in HA/Fe(II)/PMS process was heavily affected with the added

21

PMS concentration as shown in SI Figure S7. Owing to the successive recovery of Fe(II)

22

and the excess of PMS, reactive oxidants would be continuously generated via equations

23

1 and 2, consequently enhancing the degradation of BA in HA/Fe(II)/PMS process.

24 25 26

(Insert Figure 4) 3+ − HSO5− + Fe2+ → SO•− 4 + Fe + OH

k1 = 3.0 × 104 M -1s -1 32

HSO5− + Fe2+ → OH + Fe3+ + SO 42− 9


(1) (2)


1

Page 10 of 27

NH 3OH + + Fe3+ → Fe 2+ + nitrogenous products

2

HSO5− + OH → SO5•− + H 2 O k4 = 1.7 × 107 M -1s -1 11

3

HSO5− + SO•− → SO5•− + HSO 4− 4

4 5

k5 < 105 M -1s -1 10

NH 3OH + + OH → OH − + nitrogenous products k6 < 5.0 × 108 M -1s -1 11 NH 3OH + + SO•− → SO 24 − + nitrogenous products 4

k7 = 1.5 × 107 M -1s -1 10

(3) (4) (5) (6) (7)

6

Owing to the successive recovery of Fe(II) and the continuous generation of reactive

7

oxidants, PMS would be continuously decomposed via equations 1, 2, 4, and 5, resulting

8

in a much higher PMS consumption in HA/Fe(II)/PMS process than that in Fe(II)/PMS

9

process. This hypothesis was verified by the measured variation of PMS concentration in

10

HA/Fe(II)/PMS process and Fe(II)/PMS process. As shown in Figure 4, less than 4% of

11

PMS was decomposed within 15 minutes in Fe(II)/PMS process, while more than 98% of

12

PMS was decomposed within 15 minutes in HA/Fe(II)/PMS process. The results further

13

suggest that the addition of HA into Fe(II)/PMS process significantly accelerated the

14

recovery of Fe(II) and the generation of reactive oxidants, and thus improved the

15

decomposition of PMS and the degradation of BA. Meanwhile, owing to the successive

16

recovery of Fe(III) and the continuous generation of reactive oxidants, HA would be

17

continuously decomposed via equations 3, 6, and 7 in HA/Fe(II)/PMS process. As shown

18

in Figure 4, more than 60% of HA was decomposed within 15 minutes. The results

19

further indicate that the addition of HA into Fe(II)/PMS process should accelerate the

20

transformation from Fe(III) to Fe(II) and the generation of reactive oxidants. Moreover,

21

the variations of the concentrations of PMS and HA also confirmed that the steady Fe(III)

22

concentration was in dynamic equilibrium in HA/Fe(II)/PMS process.

23

Identification of Primary Reactive Oxidants in HA/Fe(II)/PMS Process. It has

24

•− been reported that three different reactive oxidants (i.e., SO•− 4 , •OH , and SO 5 ) can be

25

generated for the catalyst-mediated decomposition of PMS.2, 4, 33 As shown in equations 1, 10


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1

2, 4, and 5, all the three reactive oxidants could be generated in HA/Fe(II)/PMS process.

2

Owing to the high rate constants with SO•− ( k = 2.5 × 107 M -1s -1 10) and •OH 4

3

( k = 9.7 × 108 M -1s -1 11), methanol is an effective quencher for both SO•− and •OH . Due 4

4

to the high rate constant with •OH ( k = 6.0 ×108 M -1s -1 10) and the much slower rate

5

constant with SO•− ( k = 8.0 × 105 M -1s -1 11), TBA is an effective quencher for •OH but 4

6

•− not for SO•− is relatively inert towards TBA and methanol for the 4 . Meanwhile, SO 5

7

low rates with alcohols ( k ≤ 103 M -1s -1 34). Based on these properties, the quenching

8

experiments with methanol could allow us to differentiate between the contribution of

9

SO•− and SO•− 5 4 / •OH on the degradation of BA, while the quenching experiments with

10

TBA could allow us to differentiate between the contribution of SO•− and •OH . 4

11

Figure 5 shows the inhibition effect of TBA and methanol on the degradation of BA in

12

HA/Fe(II)/PMS process. The addition of 10 mM methanol (250 times of the initial BA

13

concentration) almost completely inhibited the degradation of BA, which excluded the

14

contribution of SO•− on the degradation of BA. Meanwhile, the degradation efficiency 5

15

of BA was decreased from more than 80% to 30% with the addition of 10 mM TBA.

16

Based on the inhibition effect of TBA and methanol on BA degradation, it could be

17

concluded that the primary reactive oxidants were SO•− and •OH in HA/Fe(II)/PMS 4

18

process. Although SO•− was commonly accepted as one of the primary reactive oxidants, 4

19

it was controversial whether •OH was one of the primary reactive oxidants in

20

Fe(II)/PMS process.2,

21

HA/Fe(II)/PMS process, nitrobenzene (NB) was selected as the indicator of •OH for the

22

high rate constant with •OH ( k = 3.9 ×109 M -1s -1 11) and a much lower rate constant with

23

SO•− ( k ≤ 106 M -1s -1 10). As shown in SI Figure S8, more than 60% of NB was degraded 4

24

within 15 minutes, while NB degradation was almost completely inhibited with the

25

addition of 10 mM TBA. The results further confirm that •OH was one of the primary

26

reactive oxidants in HA/Fe(II)/PMS process.

27

8,

17,

35

To further verify the contribution of •OH

(Insert Figure 5) 11


in


1

The deduction that SO•− and •OH were the primary reactive oxidants in 4

2

HA/Fe(II)/PMS process could be further consolidated if SO•− and •OH could be 4

3

detected directly. Therefore, an attempt with ESR experiments was made to detect SO•− 4

4

and •OH . DMPO was selected as the spin-trapping agent in ESR experiments. SO•− 4

5

and •OH could be detected by measuring the signals of DMPO-OH adducts and

6

DMPO-SO4 adducts, respectively.36-37

7

As shown in Figure 6, the special hyperfine coupling constants (a(N) 1.49 mT, a(H)

8

1.49 mT, all±0.05 mT, obtained by simulation) were consistent with that of DMPO-OH

9

adducts, while the special hyperfine coupling constants (a(N) 1.38 mT, a(H) 1.02 mT, a(H)

10

0.14 mT, a(H) 0.08 mT, all±0.05 mT, obtained by simulation) were in accordance with

11

that of DMPO-SO4 adducts. Meanwhile, the intensity of DMPO radical adducts signals in

12

HA/Fe(II)/PMS process was much stronger than that in Fe(II)/PMS process. Owing to the

13

fact that the intensity of DMPO radical adducts signals is proportional to the

14

concentrations of reactive oxidants, the concentrations of reactive oxidants in

15

HA/Fe(II)/PMS process should be much higher than that in Fe(II)/PMS process. The

16

results further confirm that both SO•− and •OH were generated in HA/Fe(II)/PMS 4

17

process and the addition of HA into Fe(II)/PMS process accelerated the generation of

18

reactive oxidants. Moreover, it should be noted that the intensity of DMPO-OH adducts

19

signals was much stronger than that of DMPO-SO4 adducts signals in HA/Fe(II)/PMS

20

process and Fe(II)/PMS process. This phenomenon could be interpreted with the fast

21

transformation from DMPO-SO4 adducts to DMPO-OH adducts via nucleophilic

22

substitution (SI Scheme S1).37

23

(Insert Figure 6)

24

End Degradation Products of HA in HA/Fe(II)/PMS Process. Since HA is a kind of

25

toxic compound,38 its degradation is necessary to study in HA/Fe(II)/PMS process. As

26

shown in Figure 4, HA was gradually decomposed in HA/Fe(II)/PMS process. Owing to

27

the low concentrations of reactive oxidants, HA should be mainly decomposed via 12


Page 12 of 27

Page 13 of 27


1

equation 3 by reacting with the successive recovery of Fe(III) in HA/Fe(II)/PMS process.

2

N2, N2O, NO −2 , and NO3− have been reported as the major degradation products of HA

3

when reacting with Fe(III) via equations 8-13.38-41

4

Fe3+ + NH 3OH + → Fe2 + + NH 2 O• + 2H +

(8)

5

NH 2 O• + NH 2 O• → N 2 + 2H 2 O

(9)

6

Fe3+ + NH 2 O• → Fe 2+ + NHO + H +

(10)

7

NHO + NHO → N 2 O + H 2 O

(11)

8

5Fe3+ + NH 2 O• + 2H 2 O → 5Fe 2+ + NO3− + 6H +

12)

9

NO3− + NH 3OH + → NO − + NO 2− + H 2 O + 2H +

(13)

10

In order to identify the final degradation products of HA in HA/Fe(II)/PMS process,

11

the concentrations of total dissolved nitrogen (TDN), Kjeldahl nitrogen, N2O, NO −2 , and

12

NO3− were measured. However, the concentration of Kjeldahl nitrogen was too low to be

13

detected. Due to the relatively high solubility of nitrogen oxides,42 the drop of TDN

14

concentration could be attributed to the generation of N2, which means [N-TDN]0 --

15

[N-TDN] = [N-N2] ([N-TDN] and [N-N2] means the concentration of nitrogen contained

16

in TDN and N2, respectively; [TDN] and [N2] means the concentration of TDN and N2,

17

respectively). As shown in SI Figure S9, the added concentration of N-HA, N-N2O,

18

N-NO −2 , N-NO3− , and N-N2 after 15 minutes (406.81 ± 2.38 µM) was equal to the initial

19

concentration of N-HA ([N-HA] means the concentration of nitrogen contained in HA;

20

[HA] means the concentration of HA). It means that the end degradation products of HA

21

should be N2O, NO −2 , NO3− , and N2 in HA/Fe(II)/PMS process. Meanwhile, the

22

environmentally friendly gas of N2 should be the main end degradation product of HA in

23

the process because N-N2 concentration was much higher than the concentrations of

24

N-N2O, N-NO −2 and N-NO3− . In previous study of Bengtssen,39 N2 was also considered

25

as the primary end product of HA when HA concentration was higher than Fe(III) 13



1

concentration. As shown in Figure 4, HA concentration was always much higher than

2

Fe(III) concentration during the reaction in HA/Fe(II)/PMS process.

3

However, it should be noted that nearly 40% of HA still existed after PMS was

4

completely decomposed in HA/Fe(II)/PMS process as shown in Figure 4. Increasing the

5

dosage of PMS from 320 µM to 960 µM, 400 µM HA could be completely decomposed

6

within 6 hours as shown in SI Figure S10. In addition, due to the presence of residual

7

PMS, the concentration of TDN measured by the Multi N/C 3100 analyzer (Jena,

8

Germany) was inaccurate. So, only the variation of the concentrations of N-N2O,

9

N-NO −2 and N-NO3− , the degradation products of HA, was shown in SI Figure S10.

10

Meanwhile, as shown in SI Figure S11, the removal of TOC was accelerated within 6

11

hours with the increase of PMS concentration in HA/Fe(II)/PMS process. Hence,

12

increasing the dosage of PMS might be a proper method to decompose the residual HA in

13

HA/Fe(II)/PMS process. In addition, with 0.96 mM PMS, the possible intermediate

14

products of BA, including phenol, 4-CP, p-CBA, and p-HBA were detected. However,

15

only p-HBA was detected during our experiments, which was shown in SI Figure S12.

16

Technical Implication. Fe(II)/PMS process, an environmentally friendly SR-AOP, has

17

advantages in degrading some refractory organic contaminants for the strong and

18

8, 18 selective oxidation capacity of SO•− Owing to the slow transformation from Fe(III) 4 .

19

to Fe(II), high Fe(II) dosage is always needed to improve the degradation of refractory

20

organic contaminants, consequently enhancing the accumulation of ferric oxide sludge.19

21

In HA/Fe(II)/PMS process, the redox cycle of Fe(III)/Fe(II) was strongly accelerated with

22

the addition of HA, so that a low Fe(II) concentration (10.8 µM) was enough to degrade

23

the probe compound rapidly in the pH range of 2.0-6.0 (more than 80% of BA could be

24

degraded in 15 minutes as shown in Figure 1). Such low concentration of Fe(II) exists

25

extensively in groundwater.43-44 Meanwhile, since both SO•− and •OH were generated 4

26

rapidly in HA/Fe(II)/PMS process, most of the refractory organic contaminants could be

27

10 degraded efficiently for the selective oxidation ability of SO•− and the strong oxidation 4

14


Page 14 of 27

Page 15 of 27


1

ability of •OH 11. Although HA itself is a kind of toxic compound,38 it could be

2

completely decomposed by increasing the dosage of PMS. Meanwhile, the major final

3

degradation product of HA was found to be the environmentally friendly gas of N2.

4

Hence, the addition of HA into Fe(II)/PMS process might be a promising way to degrade

5

refractory organic contaminants rapidly in water treatment.

6

However, this paper has just discovered an interesting phenomenon and proposed a

7

preliminary interpretation that the addition of HA into Fe(II)/PMS process could

8

significantly improve the degradation efficiency of BA by accelerating the key reactions,

9

including the redox cycle of Fe(III)/Fe(II) and the generation of reactive oxidants.

10

Although the addition of HA into Fe(II)/PMS process could efficiently alleviate the

11

accumulation of ferric oxide sludge and significantly accelerate the degradation of

12

organic compound, further investigations to relieve the negative effect of residual HA

13

and its degradation products (i.e., N2O, NO −2 , and NO3− ) should be performed prior to

14

any recommendation for practical application.

15

Acknowledgments

16

The authors greatly thank J. M. Shen for his help with ESR operation and J. J. Yang, J.

17

Zhao, G. Wen, and S. Y. Yue for discussions. This research was supported by the Funds

18

for Creative Research Groups of China (Grant No. 51121062), the Natural Science

19

Foundation of China (Grant No. 51178134), and the National Science & Technology

20

Pillar Program of China (Grant No. 2012BAC05B02).

21

Supporting Information Available

22

Texts S1-S6, Scheme S1, and Figures S1-S12. This information is available free of charge

23

via the Internet at http://pubs.acs.org.

24

References

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enhancement on Fenton oxidation by addition of hydroxylamine to accelerate the ferric

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15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32

20


Page 20 of 27

Page 21 of 27


100

Ct/C0 (%)

80

BA in HA/Fe(II)/PMS BA in Fe(II)/PMS BA in HA/PMS BA in Co(II)/PMS

60

40

20

0 0 1 2 3 4 5 6

3

6

9

12

15

Reaction Time (min) Figure 1. Degradation of BA in HA/Fe(II)/PMS process. Conditions: [HA]0 = 0.40 mM (no addition for Fe(II)/PMS and Co(II)/PMS process), [Fe(II)]0 = 10.8 µM for HA/Fe(II)/PMS and Fe(II)/PMS process (Co(II)]0 = 10.8 µM for Co(II)/PMS process), [PMS]0 = 0.32 mM, [BA]0 = 40 µM, pH0 = 3.0 (pH0 = 8.0 for Co(II)/PMS process), 25 °C. Error bars represent the standard deviation from at least duplicate experiments.

7 8

21



Page 22 of 27

100

Ct/C0 (%)

80

60

40

0.00 mM 0.05 mM 0.10 mM 0.20 mM

20

0 0 1 2 3 4 5 6 7 8 9 10 11 12 13

3

0.30 mM 0.40 mM 6

0.60 mM 1.00 mM 9

12

15

Reaction Time (min) Figure 2. Effect of HA concentration on BA degradation in HA/Fe(II)/PMS Process. Conditions: [HA]0 = 0.00-1.00 mM, [Fe(II)]0 = 10.8 µM, [PMS]0 = 0.32 mM, [BA]0 = 40 µM, pH0 = 3.0, 25 °C. Error bars represent the standard deviation from at least duplicate experiments.

22


Page 23 of 27


100

80

Ct/C0 (%)

60

40

20

0 2 1 2 3 4 5

3

4

5

6

7

Initial pH Figure 3. Effect of initial pH on BA degradation in HA/Fe(II)/PMS process. Conditions: [HA]0 = 0.40 mM, [Fe(II)]0 = 10.8 µM, [PMS]0 = 0.32 mM, [BA]0 = 40 µM, pH0 = 2.0-7.0, 25 °C, and reaction time (T) = 15 minutes. Error bars represent the standard deviation from at least duplicate experiments.

6 7 8 9

23



Page 24 of 27

12 100

9

Fe(III) in Fe(II)/PMS Fe(III) in HA/Fe(II)/PMS 60 6 40

CFe(III) (µM)

Ct/C0 (%)

80

3

PMS in Fe(II)/PMS PMS in HA/Fe(II)/PMS HA in HA/Fe(II)/PMS

20

0

0 0

3

6

9

12

15

1

Reaction Time (min)

2 3 4 5 6

Figure 4. Effect of reaction time on Fe(III) concentration and Ct/C0 of PMS and HA in HA/Fe(II)/PMS process. Conditions: [HA]0 = 0.40 mM (no addition for Fe(II)/PMS process), [Fe(II)]0 = 10.8 µM, [PMS]0 = 0.32 mM, [BA]0 = 40 µM or [methanol]0 = 10 mM, pH0 = 3.0, 25 °C. Error bars represent the standard deviation from at least duplicate experiments.

7 8 9 10 11 12

24


Page 25 of 27


100

Ct/C0 (%)

80

60

40 no alcohols 10mM Methanol 10mM TBA

20

0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

3

6

9

12

15

Reaction Time (min) Figure 5. Inhibition effect of radical scavengers on BA degradation in HA/Fe(II)/PMS process. Conditions: [HA]0 = 0.40 mM, [Fe(II)]0 = 10.8 µM, [PMS]0 = 0.32 mM, [BA]0 = 40 µM, [TBA]0 = 10 mM or [methanol]0 = 10 mM, pH0 = 3.0, 25 °C. Error bars represent the standard deviation from at least duplicate experiments.

25



Page 26 of 27

a.

Intensity

b.

c.

349

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19

350

351

352

353

354

Magnetic Field Strength (mT) Figure 6. ESR spectra obtained from (a) ultrapure water, (b) Fe(II)/PMS process, and (c) HA/Fe(II)/PMS process with the existence of DMPO (★ represents •OH adduct and █ represents SO•− adduct). Conditions: [HA]0 = 10 mM, [Fe(II)]0 = 50 µM, [PMS]0 = 32 4 mM, [DMPO]0 ≈ 0.1 N, pH0 = 3.0, 25 °C.

20 21 26


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Peroxymonosulfate Oxidation of

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