Rate equation for the degradation of nitrobenzene by ‘Fenton-like’ reagent

Advances in Environmental Research 7(2003) 583–595

1093-0191/03/$ - see front matter� 2002 Elsevier Science Ltd. All rights reserved.PII: S1093-0191Ž02.00024-2

Rate equation for the degradation of nitrobenzene by ‘Fenton-like’reagent

Miguel L. Rodriguez , Vitaliy I. Timokhin , Sandra Contreras , Esther Chamarro ,a b c c

Santiago Esplugas *c,

Universidad de Los Andes, Escuela Basica de Ingenierıa, La Hechicera, Merida, Venezuelaa ´ ´ ´Department of Physical Chemistry, Institute of Physical Chemistry, National Academy of Sciences of Ukraine, 3a Naukova Street,b

79053, Lviv, UkraineDepartament d’Enginyeria Quimica i Metallurgia, Universitat de Barcelona, Marti i Franques 1, 08028, Barcelona, Spainc ´

Accepted 2 April 2002

Abstract

This paper describes the effect of temperature and initial concentration of H O , Fe(II), PhNO and dissolved2 2 2

oxygen on the degradation rate of PhNO in homogeneous aqueous solution by ‘Fenton-like’ reagent2

(wH O x 4wFe(II)x ). The oxidation productso-, m- and p-nitrophenol were found as intermediates in the ratio2 2 o o

1:1.3–2.8:1.4–2.7 as compared with PhNO when conversion of the latter was less than 25%. This fact suggests that2

hydroxylation of PhNO was promoted by HO radicals. The reaction was investigated in a completely mixed-batch•2

reactor under a wide range of experimental conditions: pH;3.0; 278–318 K; 1.5-wH O x -26.5 mM; 0.04-2 2 o

wFe(II)x -1.1 mM; 0.3-wPhNO x -2.5 mM; and 0-wO x -1.4 mM. The activation energy for the degradation ofo 2 o 2 o

PhNO was determined to be 59.7 kJ mol . The degradation rate of PhNO follows pseudo-first-order kinetics. They12 2

results of this study demonstrate that the degradation rate of PhNO in the ‘Fenton-like’ system could be predicted2

with sufficient precision by the equationR s1.05=10 exp(y59.7yRT)wH O x wFe(II)x wPhNO x . The11 0.68 1.67 y0.32D 2 2 o o 2 o

second-order rate constant for the overall rate of H O decomposition by Fe(III ) was found to be 0.83 M s aty1 y12 2

298 K. The value of the steady-state HO radical concentration in the ‘Fenton-like’ reaction was found to be•

;10 M, as estimated by two independent methods.y13

� 2002 Elsevier Science Ltd. All rights reserved.

Keywords: Chemical degradation; Hydroxyl radical; Kinetics; Hydrogen peroxide; Iron salts

1. Introduction

Nitroaromatic compounds are widely used as rawmaterials in many industrial processes, such as thepreparation of pesticides, explosives, textiles and paper.Consequently, these compounds are found as waterpollutants as a result of their release in industrialwastewater(Beltran et al., 1998; Sarasa et al., 1998).Nitroaromatic compounds are not only important as

*Corresponding author. Tel.:q34-93-402-1290; fax:q34-93-402-1291.

E-mail address: esplugas@angel.qui.ub.es(S. Esplugas).

constituents in industrial waste streams, but also ashazardous wastes that have contaminated soils, ground-water, sludges and other hazardous solid wastes regulat-ed under the Resource Conservation and Recovery Act(US EPA, 1980; Sittig, 1985). Remediation of waste-waters containing these pollutants is very difficult, sincethey are usually resistant to biological degradation(O’Connor and Young, 1989).Oxidation processes that involve the generation of

the highly reactive hydroxyl radical(HO ) are of current•

interest for the destruction of organic pollutants insurface and groundwaters, and industrial wastewaters.

584 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Generation of HO radicals by the dark reaction of•

H O with ferrous salt(known as Fenton’s reagent) has2 2

been the subject of numerous studies during the lastdecade(Arnold et al., 1995; Pignatello and Day, 1996;Chen and Pignatello, 1997; Hislop and Bolton, 1999).The oxidizing species generated in the Fenton reactionhave been discussed by many investigators, but are stillcontroversial (Walling, 1975; Stubbe and Kozarich,1987; Bossmann et al., 1998; MacFaul et al., 1998;Goldstein and Meyerstein, 1999; Kremer, 1999). Therecognition of the HO radical as the active intermediate•

is not yet universal, and doubts as to its very existencein the system have even been raised(Bossmann et al.,1998; Kremer, 1999). In other studies of Fenton’sreagent, it is generally considered that the reactionbetween H O and Fe(II) in acidic aqueous medium2 2

(pH F3) produces HO radicalswEq. (1)x and can•

involve the steps presented belowwEq. (1)–Eq. (6)x(Haber and Weiss, 1934; Barb et al., 1951a,b; Walling,1975). The rate constants are reported at 298 K inM s for a second-order reaction rate(Lin andy1 y1

Gurol, 1998; De Laat and Gallard, 1999).

Fe IIqH OŽ . 2 2y •™Fe III qHO qHO k s63 (1)Ž . 1

Fe III qH OŽ . 2 2q •™Fe IIqH qHOO k s0.002y0.01 (2)Ž . 2

•Fe IIqHOŽ .y 8™Fe III qHO k s3=10 (3)Ž . 3

• •HOqH O ™HOOqH O2 2 27k s2.7=10 (4)4

•Fe IIqHOOŽ .y 6™Fe III qHOO k s1.2=10 (5)Ž . 5

•Fe III qHOOŽ .q 3™Fe IIqH qO k -2=10 (6)Ž . 2 6

The hypothesis of Haber and Weiss(1934) that theFenton reaction involves the formation of HO radicals•

as the actual oxidants has been proved by many tech-niques, including EPR spectroscopy. Although a consid-erable number of investigators, using the electronparamagnetic resonance(EPR) spin-trapping technique,have found evidence for the formation of HO radicals•

from Fenton’s reagent(Dixon and Norman, 1964; Buett-ner, 1987; Rosen et al., 2000), it has also been reportedby others (Rush and Koppenol, 1987; Rahhal andRichter, 1988) that this species is not the only oxidizingintermediate, but that some type of high-valent iron-oxointermediates also exist(Groves and Watanabe, 1986;Kean et al., 1987; Sychev and Isak, 1995; Bossmann etal., 1998; Kremer, 1999). Using EPR spin-trapping,three types of oxidizing species(free HO , bound HO• •

and high-valence iron species, which is probably a ferryl

ion, Fe _O) were detected by Yamazaki and PietteIV

(1991). In the present work, the principal oxidant isassumed to be HO radical, but others, such as iron-oxo•

species, cannot be ruled out.Nitrobenzene is one of the most representative nitroar-

omatic compounds present in several wastewaters andit is considerably soluble at room temperature. A numberof studies on the degradation of PhNO in aerated2

aqueous solutions by Fenton’s reagent have been report-ed (Lipczynska-Kochany, 1991, 1992), although theyare actually ‘Fenton-like’ processes becausewH O x 4wFe(II)x . However, in a system containing2 2 o o

PhNO , H O and Fe(II), the effects of temperature and2 2 2

initial concentrations of these components and dissolvedoxygen on the degradation rate of PhNO have not been2

investigated in detail. Nitrobenzene half-lives of 360min at wPhNO x s0.1 mM, wH O x s3.9 mM,2 o 2 2 o

wFeCI x s0.035 mM(Lipczynska-Kochany, 1991) and2 o

of 250 min atwPhNO x s0.1 mM, wH O x s8.0 mM,2 o 2 2 o

wFeCI x s0.035 mM(Lipczynska-Kochany, 1992) have2 o

been determined, but these were studied only at roomtemperature and were reported without any experimentaldetails. In spite of the numerous studies of Fenton’sreagent, the chemistry and kinetics of PhNO oxidation2

has not been well elucidated. Due to its importance asan environmental pollutant, this compound was selectedfor a study on its chemical degradation by ‘Fenton-like’reagent. The study was conducted with the aim ofdetermining some intermediate products and kineticparameters for PhNO removal in pure water. Knowl-2

edge of the kinetic information is necessary for betterthe planning of better methods of degradation of thepollutant, as well as for mechanistic studies.

2. Experimental section

2.1. Materials and reagents

Nitrobenzene 99%(Probus), o-, m- andp-nitrophenol)99%(Merck), hydrogen peroxide 30%(Merck), oxy-gen 99.99%(AlphaGas), nitrogen 99.99%(AlphaGas),acetonitrile 99.8%, isocratic grade for HPLC(Merck),FeSOØ7H O 98%(Panreac) and sodium hydrogen sul-4 2

fite solution 40% wyv (Panreac) were used as received.All solutions of PhNO , H O and ferrous salt were2 2 2

prepared in Millipore water.

2.2. Degradation experiments

All experiments were conducted in a thermostattedbatch glass reactor(1 l) equipped with a magneticstirrer in the absence of light, and under air pressure ofapproximately 1=10 Pa (wO x s0.27 mM). Kinetic5

2 o

experiments on PhNO degradation were initiated by2

adding a known amount of H O into the reactor, which2 2

contained PhNO and FeSO solutions under vigorous2 4

585M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

magnetic stirring. The reactions were carried out in non-buffered conditions and pH values for all experimentsdecreased during the reaction from 3.5 to 2.5. In thisrange, the values of the overall rate constants for theH O decomposition by Fe(II) and Fe(III ), and conse-2 2

quently the PhNO degradation rate, are relatively inde-2

pendent of pH within experimental error(De Laat etal., 1999; Hislop and Bolton, 1999). Samples wereperiodically withdrawn during kinetic experiments,quenched with sodium hydrogen sulphite solution 40%wyv to avoid further reactions, and used for analysis.For the experiments with oxygen and nitrogen, gases

were bubbled into the solution for 30 min before theaddition of H O and continuously during the reaction.2 2

It is important to establish that the disappearance ofPhNO in this batch reactor proceeds by the ‘Fenton-2

like’ reaction, i.e. to exclude the possibility of vapori-zation of PhNO from water. To this end, a control2

experiment was carried out with prolonged heating(2h) at 318 K without Fenton’s reagent. Control experi-ments were also conducted for sparged samples. Theseexperiments showed that absolutely no change wasobserved in PhNO initial concentration(HPLC2

analysis).

2.3. Selection and range of experimental variables

Several series of experiments on PhNO degradation2

by ‘Fenton-like’ reagent(wH O x 4wFe(II)x ) were2 2 o o

conducted by varying the temperature(278–318 K), theinitial concentrations of H O , Fe(II) and PhNO , and2 2 2

dissolved oxygen(0–1.4 mM). Table 1 summarizes thevalues of these operating variables in a group of exper-iments. The initial concentration of PhNO was tested2

from 0.37 to 2.46 mM. Hydrogen peroxide and Fe(II)concentration ranges were chosen to achieve the highestPhNO degradation percentage in a time interval from2

5 to 120 min. The wH O x ywPhNO x , wH O x y2 2 o 2 o 2 2 o

wFe(II)x and wPhNO x ywFe(II)x ratios varied from 2.1o 2 o o

to 16.7, 6.8 to 102.8 and 0.8 to 16.4, respectively.

2.4. Analytical methods

Qualitative and quantitative determinations ofPhNO , o-, m- and p-nitrophenols were carried out by2

high-performance liquid chromatography(HPLC) usingauthentic standards. Chromatography was performedusing a Spherisorb ODS 2(5 mm, 25 cm=4.6 mm i.d.)column, mobile phase of acetonitrile–water(40:60 vyv), pH 2.5 at a flow rate of 1 ml min and a UVy1

detector(Waters 996 photodiode array detector). ThePhNO ,o-, m- andp-nitrophenol were detected at 267,2

282, 277 and 321 nm, respectively. Under these condi-tions, the retention times of the compounds were asfollows (min): 7.5 (p-nitrophenol), 7.8(m-nitrophenol),11.8 (o-nitrophenol) and 13.8(PhNO ). Solution pH2

was measured during the reaction using a GLP 22 pHmeter(Grison).

3. Results and discussion

3.1. Intermediates in the degradation of PhNO by2

HPLC

To clarify the reaction pathways for PhNO after the2

dark ‘Fenton-like’ reaction, we worked on the identifi-cation of reaction products by HPLC. Identification ofnitrophenols as reaction intermediates suggests that thedegradation of PhNO by the ‘Fenton-like’ process takes2

place by reaction with HO radical, which is assumed•

to be the main oxidant. Reaction of PhNO with HO•2

radicals to form HO –PhNO adducts(nitrohydroxycy-•2

clohexadienyl radicals) has a significantly high rateconstant:

• •HOqPhNO™o-, m- and p-isomeric HO2

yPhNO adducts (7)2

with k s3=10 M s at 298 K(Hoigne, 1997). It9 y1 y17

is worth mentioning that hydroxycyclohexadienyl radi-cals have been observed by UV and EPR methods inthe reaction of benzene with HO radicals(Walling,•

1975). When the reaction is run in the presence ofoxygen, the HO –PhNO adducts formed react prefera-•

2

bly with dissolved oxygen(Kunai et al., 1986; Bohnand Zetzsch, 1999) compared to HOyHOO radicals or• •

Fe(II)yFe(III ) ions due to their very low steady-stateconcentration, thus producing nitrophenols and HOO•

radicals:

•o-, m- and p-isomeric HOyPhNO adducts2•qO ™o-, m- and p-nitrophenols qHOO2

(8)

In the absence of oxygen, the HO –PhNO adducts•2

formed react preferably with Fe(II) or Fe(III ) (Kolthoffand Medelia, 1949a,b; Walling, 1975), eventually lead-ing to the formation of isomeric nitrophenols as reactionintermediates.When conversion of PhNO was 15–25%, the ratio2

of the o-, m- and p-isomers was 1:1.3–2.8:1.4–2.7. Athigh conversion of PhNO(80–90%), the ratio was2

found to be 1:0.6–1.0:0.8–1.2. As can be observed, theformation of theo-isomer is increased at higher conver-sion of PhNO , which can be expected if theo-2

nitrophenol degradation rate is lower than that of them- and p-isomers. In fact, it has been shown byLipczynska-Kochany(1991) that theo-nitrophenol deg-radation rate is lower than that ofp-nitrophenol. Besides,we found that theo-nitrophenol degradation rate islower than that ofm-nitrophenol under the same exper-imental conditions.

586 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Table 1Effect of wH O x , wFe(II)x , wPhNO x , wO x and temperature on the pseudo-first-order rate constant(k ) and degradation rate of2 2 o o 2 o 2 o obs

PhNO (R ) with ‘Fenton-like’ reagent2 D

Run wPhNO x2 o wH O x2 2 o wFe(II)xo T k =103obs R =10 (M s )6 y1D Da

(mM) (mM) (mM) (K) (s )y1

Eq. (10) Eq. (17)(%)

Effect of wH O x2 2 o

6 1.74 0.55 0.45 0.47 4.45 3.54 0.82 0.67 0.76 13.43 0.82 6.37 0.26 298 1.44 1.18 1.14 3.44 6.99 1.54 1.27 1.21 4.77 13.24 2.02 1.66 1.88 13.39 26.45 1.96 1.61 – –

Effect of wFe(II)xo

8 0.13 0.38 0.31 0.38 22.64 0.82 6.99 0.26 298 1.54 1.27 1.21 4.710 0.52 4.74 3.90 3.91 0.311 1.04 12.4 10.2 12.5 22.5

Effect of wPhNO x2 o

17 0.43 3.24 1.39 1.50 7.94 0.82 6.99 0.26 298 1.54 1.27 1.21 4.718 1.63 0.62 1.01 0.97 4.019 2.46 0.32 0.78 0.85 9.0

Effect of wO x2

22b 1.72 1.41 1.21 14.24c 0.82 6.99 0.26 298 1.54 1.27 1.21 4.721d 1.51 1.24 1.21 2.4

Effect of T15 278 0.20 0.16 0.21 31.313 288 0.66 0.54 0.53 1.94 0.82 6.99 0.26 298 1.54 1.27 1.21 4.714 308 3.18 2.61 2.65 1.516 318 5.03 4.13 5.52 33.7

1 0.81 2.43 0.049 298 0.058 0.047 0.037 21.32 0.85 7.96 0.26 298 1.98 1.67 1.31 21.612 0.82 4.41 0.26 298 1.55 1.27 0.89 29.920 0.37 1.55 0.086 298 0.30 0.11 0.09 18.2

DscR wEq. (10)xyR wEq. (17)xc=100y wEq. (10)x.a RD D D

Under nitrogen-saturated conditions(wO x (0 mM).b2 o

Under air-saturated conditions(wO x (0.27 mM).c2 o

Under oxygen-saturated conditions(wO x (1.4 mM).d2 o

Comparison of isomeric ratios can be used to verifyif the HO radical is the attacking species in different•

chemical systems. The isomeric distributions of nitro-phenols formed during the degradation of PhNO by2

‘Fenton-like’ reagent, photolysis and radiolysis are pre-sented in Table 2. These ratios reflect the extent of theattack by the HO radical at different positions of the•

aromatic ring and are in agreement with those expectedfor a very reactive homolytic reagent with some electro-philic character. As can be observed from Table 2, theratio of theo-, m- andp-isomers is essentially the sameunder the different methods of degradation. This fact

supports the hypothesis of the generation of HO radicals•

in the ‘Fenton-like’ reaction.

3.2. Degradation of PhNO by ‘Fenton-like’ reagent:2

kinetic studies

The linear regression analysis of the data representingPhNO concentration vs. reaction time indicates that in2

a first approach, the oxidation reaction Eq.(7) can bedescribed by pseudo-first-order kinetics with respect toPhNO concentrationwEq. (9)x:2

w x w xln PhNO y PhNO syk t (9)2 2 o obsŽ .

587M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Table 2The isomeric distribution of nitrophenols formed during the action of Fenton reagent, photolysis and radiolysis on the aqueoussolutions of PhNO2

Source of the attacking Ratio of theo-, m- Referencespecies andp-isomers

‘Fenton-like’ reagenta 1:1.3–2.8:1.4–2.7 This work‘Fenton-like’ reagentb 1:0.6–1.0:0.8–1.2 This work‘Fenton-like’ reagent 1:1.3:1.9 Norman and Radda, 1962‘Fenton-like’ reagent 1:0.8:1.8 Loebl et al., 1949Photolysisc 1:0.6:0.7 Lipczynska-Kochany, 1992Radiolysis Loebl et al., 1950pH 2 1:0.9:1.0pH 6 1:1.0:1.0

Radiolysis(pH 5.5) Matthews and Sangster, 1967Air 1:0.6:0.7Nitrogen 1:0.7:1.8

Conversion of PhNO was 15–25%.a2

Conversion of PhNO was 80–90%.b2

In the presence of H O .c2 2

Fig. 1. Effect of initial concentration of H O on the degradation of PhNO . First-order plots for PhNO degradation at pH;3.02 2 2 2

by ‘Fenton-like’ reagent:wPhNO x , 0.82 mM; wFe(II)x , 0.26 mM; wH O x , (�) 1.74, (j) 3.54, (m) 6.37,(h) 6.99,(n) 13.242 o o 2 2 o

and(e) 26.45 mM at 298 K.

where k represents the pseudo-first-order rate con-obs

stant. In this case, a plot of ln(wPhNO xyxwPhNO x ) vs.2 2 o

time in every experiment must lead to a straight linewith slope of k . Figs. 1–5 show these plots forobs

experiments in which the temperature and initial con-centration of H O , Fe(II) and PhNO , and dissolved2 2 2

oxygen were varied. As it can be observed, points liein satisfactory straight lines with correlation coefficientsgreater than 0.92. It is worth noting that thek valuesobs

(Table 1) were calculated from experimental data cov-ering 76–99% removal of PhNO . In experiments car-2

ried out in duplicate,k varied by less than 10%. Allobs

these results support the pseudo-first-order kineticsassumed.Nitrobenzene degradation rates(R ) were calculatedD

by Eq. (10) (Table 1):

w x w xR syd PhNO ydtsk PhNO (10)D 2 obs 2 o

588 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Fig. 2. Effect of initial concentration of Fe(II) on the degradation of PhNO . First-order plots for PhNO degradation at pH;3.02 2

by ‘Fenton-like’ reagent:wPhNO x , 0.82 mM; wH O x , 6.99 mM; wFe(II)x , (�) 0.13, (j) 0.26, (m) 0.52 and(h) 1.04 mM at2 o 2 2 o o

298 K.

Fig. 3. Effect of initial concentration of PhNO on the degradation of PhNO . First-order plots for PhNO degradation at pH;3.02 2 2

by ‘Fenton-like’ reagent:wFe(II)x , 0.26 mM; wH O x , 6.99 mM; wPhNO x , (�) 0.43, (j) 0.82, (m) 1.63 and(h) 2.46 mM ato 2 2 o 2 o

298 K.

589M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Fig. 4. Effect of the concentration of dissolved oxygen on the degradation of PhNO . First-order plots for PhNO degradation at2 2

pH ;3.0 by ‘Fenton-like’ reagent:wPhNO x , 0.82 mM; wFe(II)x , 0.26 mM; wH O x , 6.99 mM; wO x , (�) 0, (j) 0.27 and(n)2 o o 2 2 o 2 o

1.4 mM at 298 K.

Fig. 5. Effect of temperature on the degradation of PhNO . First-order plots for PhNO degradation at pH;3.0 by ‘Fenton-like’2 2

reagent:wPhNO x , 0.82 mM; wFe(II)x , 0.26 mM; wH O x , 6.99 mM;T, (�) 278, (j) 288, (m) 298, (h) 308 and(e) 318 K.2 o o 2 2 o

590 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Careful analysis of our reaction conditions forPhNO degradation, taking into account the reaction2

mechanism proposed by Walling for Fe(II) with H O2 2

wEq. (1)–Eq. (6)x, shows why it is a ‘Fenton-like’process. The rate constant for the reaction of ferrousions with H O is high(k s63 M s at 298 K)y1 y1

2 2 1

(De Laat and Gallard, 1999), and therefore Fe(II)oxidizes to Fe(III ) in a few seconds in the presence ofexcess H O (Table 1; 6.8FwH O x ywFe(II)x F102.8).2 2 2 2 o o

If Fe(II) oxidation is a pseudo-first-order process, thetime required for the conversion of 99% of Fe(II) toFe(III ) can be calculated by Eq.(11):

w xt sln100y k H O (11)99 1 2 2 oŽ .and was found to be 2.6–46 s(Table 1). Therefore,PhNO degradation by Fenton’s reagent performed in2

this study(at wH O x 4wFe(II)x ) is simply a Fe(III )y2 2 o o

H O catalyzed process. Fenton’s reagent with an excess2 2

of H O with respect to Fe(II) is known as a ‘Fenton-2 2

like’ reagent(Pignatello, 1992). It is worth noting thatsome publications describe the degradation of PhNO2

by Fenton’s reagent(Lipczynska-Kochany, 1991, 1992),although the reactions were actually proceeding by a‘Fenton-like’ process.

3.3. Steady-state HO radical concentration•

If a steady-state HO radical concentration(wHO x )• •ss

is assumed, thewHO x value in such a system can be•ss

estimated by Eq.(12):•w xHO sk yk M (12)Ž .ss obs 7

The steady-state HO radical concentration occurs only•

when the H O concentration is relatively constant2 2

during the experiment. Furthermore, thewHO x concen-•ss

tration is governed by both its formation rate, dependenton wH O x and wFe(II)x , and the scavenging rate. In2 2 o o

this study, the time for H O complete decomposition2 2

(t ) in the ‘Fenton-like’ reaction could be estimated byd

Eq. (13) { assuming thatwFe(III )x(wFe(II)x and theo

limiting step for this process is Eq.(2)} :

w xw x w xt s H O y k H O Fe II (13)Ž .d 2 2 o d 2 2 oŽ .owherek is the second-order rate constant for the overalld

rate of decomposition of H O by Fe(III ); k was found2 2 d

to be 0.47 M s at 298 K and pH 3.0(De Laat andy1 y1

Gallard, 1999). Calculations according to Eq.(13) showthat the H O concentration does not significantly2 2

change, even up to complete degradation of PhNO2,

which can be expected for a ‘Fenton-like’ process. Forexample, for run 5, the time for H O complete decom-2 2

position was found to be 136 min, with the rate constant(k ) determined within a 60 min interval. In this case,obs

the H O was decomposed less than 45%, i.e. more2 2

than 55% of H O was present in the solution after2 2

complete degradation of PhNO . It can be concluded2

that the decrease in the initial H O concentration does2 2

not have a significant effect on the HO radical steady-•

state concentration. Furthermore, the linearity of theln(wPhNO xywPhNO x ) vs. time plots for all the runs2 2 o

(Figs. 1–5) gives evidence of a HO radical steady-state•

concentration. ThewHO x values were calculated using•ss

Eq. (12) and were found to be very low(10 –10y14 y12

M), which is in agreement with data reported for theoxidation of atrazine(;10 M) by ‘Fenton-like’y13

reagent{ Fe(III )yH O } (Gallard and De Laat, 2000).2 2

The wHO x values were calculated assuming that the•ss

rate of production of HO radicals in the ‘Fenton-like’•

processwlimiting step, Eq.(2)x is equal to the decayrate of the radicals by reaction with PhNOwEq. (7)x.2

The iron dynamics can be extremely complex in the‘Fenton-like’ reaction because of a shift from Fe(II) toFe(III ), along with precipitation of the iron ions asamorphous oxyhydroxides. These iron speciation chang-es can affect the stoichiometry of HO radical genera-•

tion. As a result, these systems cannot be considered asbeing in steady state. Given that all the experimentswere carried out at pH ;3.0 and withwH O x 4wFe(II)x , no iron precipitates were observed.2 2 o o

Several papers are available concerning H O loss2 2

and HO production in the reaction of H O with iron•2 2

salts and oxides(Lin and Gurol, 1998; Lindsey andTarr, 2000). The formation of HO radicals in pure water•

was studied using chemical probes(benzoic acid and1-propanol), and a linear increase with H O concentra-2 2

tion was found(Lindsey and Tarr, 2000). In our study,the steady-state HO radical concentration also showed•

a linear dependence on the H O initial concentration:2 2

•w xHO s4.30ssy11 y13w x=10 H O q1.48=10 M (14)Ž .2 2

with constant concentrations of Fe(II) and PhNO at2

298 K.

3.4. Effect of wH O x on the degradation rate (R )2 2 o D

The H O concentration effect onR can be observed2 2 D

in runs 3–7 and 9, with constant Fe(II) and PhNO2concentrations(Table 1 and Fig. 6a). Initially, it can beexpected that, as the molar ratio of H O to pollutant is2 2

increased, there are more HO radicals available to attack•

the aromatic structure, and therefore the degradationreaction rate must increase. In the present case, this rateincrease was obtained for H O initial concentrations in2 2

the range 1.74–13.24 mM. However, higher H O con-2 2

591M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Fig. 6. Effect of wH O x , wFe(II)x , wPhNO x and wO x on the degradation rate of PhNO(R ) by ‘Fenton-like’ reagent at pH2 2 o o 2 o 2 o 2 D

;3.0, 298 K. (a) Effect of wH O x : wFe(II)x , 0.26 mM; wPhNO x , 0.82 mM. (b) Effect of wFe(II)x : wH O x , 6.99 mM;2 2 o o 2 o o 2 2 o

wPhNO x , 0.82 mM.(c) Effect of wPhNO x : wH O x , 6.99 mM;wFe(II)x , 0.26 mM.(d) Plot of k vs. 1ywPhNO x : wH O x , 6.992 o 2 o 2 2 o o obs 2 o 2 2 o

mM; wFe(II)x , 0.26 mM.o

592 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

centrations(wH O x )13.24 mM, run 9) led to a similar2 2 o

PhNO removal rate. This suggests that the PhNO2 2

degradation rate becomes insensitive to H O concentra-2 2

tion when the H OyPhNO molar ratio is higher than2 2 2

16 (Fig. 6a). At high H O concentration it becomes2 2

the main sink for HO radicals according to Eq.(4).•

Therefore, the effect of a high H O concentration can2 2

be attributed to the fact that it is competing withPhNO wEq. (7)x for HO radicals wEq. (4)x, thus•

2

decreasing the steady-state concentration and thePhNO oxidation rate. As a result, the steady-state2

concentration of HO radicals and the PhNO degrada-•2

tion rate at the plateau should be constant.The rate order for H O(wH O x -13.24 mM) was2 2 2 2 o

determined to be 0.68 from a logR vs. logwH O x plot.D 2 2 o

3.5. Effect of wFe(II)x on degradation rate (R )o D

The initial concentration of Fe(II) (runs 4, 8, 10 and11 in Table 1) has a significant influence on thePhNO degradation rate. At constant H O and PhNO2 2 2 2

concentrations,R increases with the Fe(II) concentra-D

tion (see Fig. 6b). Given that wH O x4wFe(II)x, the2 2

metallic ion is consumed during the first seconds(1.4–23 s; see above), which leads to fast production ofHO radicals, and consequently to rapid PhNO degra-•

2

dation. After a few seconds, Eq.(2) becomes thelimiting step, which regenerates Fe(II) ions. Hydroxylradical formation from this reaction becomes very slowand PhNO degradation proceeds via a ‘Fenton-like’2

reagent.The rate of H O decomposition by Fe(III ) has been2 2

extensively studied(Barb et al., 1951a,b; Walling andWeill, 1974; Kremer and Stein, 1977; De Laat andGallard, 1999; Gallard et al., 1999; Gallard and DeLaat, 2000). Many kinetic models derived from hypo-thetical mechanisms have been tested(Barb et al.,1951a,b; Walling and Weill, 1974; Kremer and Stein,1977). Recently, De Laat and Gallard(1999) demon-strated that the rate of H O decomposition by Fe(III )2 2

could be very accurately predicted by a kinetic modelthat takes into account the rapid formation and slowerdecomposition of Fe(III )–hydroperoxy complexeswFe (HO ) and Fe (OH)(HO ) x. For 5-wH O x yIII 2q III q

2 2 2 2 o

wFe(III )x -500, a second-order kinetic law describeso

the H O initial decomposition rate(De Laat and Gal-2 2

lard, 1999):y1w xw x w xyd H O ydtsk H O Fe III M s (15)Ž . Ž .2 2 d 2 2 o

wherek is the second-order rate constant for the overalld

rate. Thek value was found to be 0.47 M s aty1 y1d

298 K and pH 3.0(De Laat and Gallard, 1999). Therate constant for hydroxyl radical generationwEq. (1)xis more than two orders of magnitude higher than therate constant for H O decomposition by Fe(III ). There-2 2

fore, after a few seconds the latter reaction becomes thelimiting step.

The rate order for Fe(II) (wFe(II)x -1.04 mM) waso

determined to be 1.67 from a logR vs. logwFe(II)xD o

plot.

3.6. Effect of wPhNO x on degradation rate (R )2 o D

The effect of nitrobenzene concentration can beobserved in runs 4 and 17–19 at constant H O and2 2

Fe(II) concentrations(Table 1). The degradation rateshows non-linear dependence on the initial PhNO2

concentration(Fig. 6c) and is somewhat decreased whenits concentration is increased. If it is assumed that theproduction rate of HO radicals is equal to the decay•

rate by reaction with PhNO wEq. (7)x, i.e.2

k wFe(III )xwH O x sk wPhNO x , there must be a lin-d 2 2 o obs 2 o

ear dependence ofk vs. 1ywPhNO x wEq. (16)x. Inobs 2 o

other words, thek value will be inversely proportionalobs

to the PhNO initial concentration while all other para-2

meters are kept constant:y1w xw x w xk sk Fe III H O y PhNO s (16)Ž . Ž .obs d 2 2 o 2 o

Fig. 6d shows the linear regression analysis for Eq.(16), giving a slope(k wFe(III )xwH O x ) of 1.5=10y6

d 2 2 o

M s . Taking wFe(III )x(wFe(II)x s0.26 mM andy1o

wH O x s6.99 mM (Table 1), we calculatedk s0.832 2 o d

M s at 298 K. In previous work carried out by Dey1 y1

Laat and Gallard(1999), the second-order overall rateconstant for H O decomposition was found to be 0.472 2

M s at 298 K and pH 3.0.y1 y1

The rate order for PhNO was determined to be2

y0.32 from a logR vs. logwPhNO x plot. SincekD 2 o obs

is inversely proportional to the initial PhNO concentra-2

tion (Fig. 6d), this species degradation rate should beconcentration-independentwEq. (16)x, i.e. the PhNO2degradation rate order should be zero. The deviationfrom zero with such unusual order(y0.32) can ariseas a result of several successive steps of differentstoichiometries interacting in the reaction of H O with2 2

Fe(II) in the presence of PhNO .2

3.7. Effect of wO x on degradation rate (R )2 o D

It has been observed that dissolved oxygen(0-wO x -1.4 mM) has a very weak effect on the PhNO2 o 2

degradation rate(Table 1) When oxygen is replaced by.

nitrogen, the PhNO degradation rate suffers a decrease2

of approximately 12%, which is close to the precisionin the determination ofk . With this slight influence,obs

the rate order for molecular oxygen will be zero. Thesedata can indicate that organic radicals R and ROO do• •

not affect the HO radical steady-state concentration and•

oxygen is probably not involved in the first step ofPhNO oxidation mediated by hydroxyl radicals. Thus,2

the oxygen effect may not be apparent when observingPhNO loss by a ‘Fenton-like’ process. It may be2

involved in subsequent steps, i.e. further reactions ofthe intermediate organic radicals. For example, theHO –PhNO adduct reacts faster with oxygen than with•

2

HO or HOO radicals, producing nitrophenols according• •

to Eq. (8) in the presence of dissolved oxygen. In the

593M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

Fig. 7. Plot of lnR vs. 1yT for the degradation of PhNO(pHD 2

;3.0; wPhNO x , 0.82 mM;wH O x , 6.99 mM;wFe(II)x , 0.262 o 2 2 o o

mM).

absence of oxygen, the HO –PhNO adducts or organic•2

radicals resulting from secondary reactions may eitherdimerize or react with Fe(II) or Fe(III ) (Walling, 1975;Sobkowiak et al., 1992; Bandara et al., 1997). Thereaction of HO –PhNO adducts with Fe(II) or Fe(III )•

2

may also lead to the formation of isomeric nitrophenols.In addition, other hydroquinoneyquinone-type inter-

mediate by-products can be oxidized or reduced byFe(II) or Fe(III ) and, in the absence of molecularoxygen, these reactions may have a significant effect onthe overall rate of formation and distribution of finalproducts(Kolthoff and Medelia, 1949a,b; Norman andRadda, 1962; Kunai et al., 1986; Bohn and Zetzsch,1999). This aspect of the effect of oxygen was notstudied in this work.It is worth noting that the pH in all the experiments

decreased during the reaction, going from 3.5 to 2.5,i.e. the reactions proceeded in acid medium. At suchpH values, Fe(II) cannot be rapidly oxidized to Fe(III )by molecular oxygen. Indeed, in reports by Lin(1997),and Lin and Gurol(1998), no significant effect ofdissolved oxygen on the H O decomposition rate was2 2

observed in the heterogeneous catalytic reaction withgranular-size goethite(a-FeOOH) particles in aqueoussolution under various experimental conditions.

3.8. Effect of temperature on degradation rate (R )D

Degradation of PhNO was carried out at five differ-2

ent temperatures(from 278 to 318 K) (Table 1; runs 4and 13–16). As can be expected, temperature exerts astrong effect on the PhNO degradation rate, which is2

increased at high temperature due to an increment inthe pseudo-first-order rate constant. The data exhibitArrhenius-type behavior with an activation energy of59.7 kJ mol wEq. (17)x, calculated from the usualy1

logR vs. 1yT plot (Fig. 7):D

4 y1R s3.2=10 expy59.7 kJ mol yRTŽ .Dy1M s (17)Ž .

It is interesting to note that the activation energy issomewhat higher than values measured for ferrous ion-catalyzed decomposition of H O in different media2 2

(sulfuric or perchloric acid) of 39.5 and 40.8 kJ moly1

(Hardwick, 1957).

3.9. Rate equation for the degradation of PhNO2

The rate equation can be expressed in a simple wayas:

ba c dw xw x w x w xR sA H O Fe III PhNO O expyE yRTŽ . Ž .D 2 2 o 2 o 2 o aoy1M sŽ .

whereE is the activation energy of PhNO degradationa 2

(59.7 kJ mol ). The exponentsas0.68,bs1.67,csy1

y0.32 andds0 represent the reaction orders for H O ,2 2

Fe(II), PhNO and molecular oxygen, respectively. The2

coefficient A can be calculated on the basis of theexperimental PhNO degradation ratewEq. (10); Table2

1x at the temperature determined and H O , Fe(II) and2 2

PhNO initial concentrations. This coefficient was deter-2

mined to be 1.05=10 as an average value from 1911

experiments(Table 1). Given all these parameters, therate equation for PhNO degradation via ‘Fenton-like’2

reagent becomes:

1.6711 0.68 y0.32w xw x w xR s1.05=10 H O Fe III PhNOŽ .D 2 2 o 2 ooy1=expy59.7yRT M s (18)Ž .Ž .

The degradation rates of PhNO calculated using Eq.2

(17) are presented in Table 1, together with thosepredicted from Eq.(10). As shown in Table 1, there isgood agreement between experimental and calculateddegradation rates. Thus, empirical Eq.(17) can be usedfor the prediction of PhNO degradation rates with2

sufficient precision(ca. 13.4%) under a wide range ofexperimental conditions: pH;3.0; 278–318 K; 1.5-wH O x -26.5 mM; 0.04-wFe(II)x -1.1 mM; 0.3-2 2 o o

wPhNO x -2.5 mM; and 0-wO x -1.4 mM.2 o 2 o

4. Conclusions

Nitrobenzene removal has been investigated by apply-ing ‘Fenton-like’ reagent(wH O x 4wFe(II)x ), which2 2 o o

was found to be an appropriate method to efficiently

594 M.L. Rodriguez et al. / Advances in Environmental Research 7 (2003) 583–595

remove this compound from aqueous solutions. Thestudy and mathematical treatment of typical operatingparameters in this system resulted in the followingstatements related to the kinetics of the reaction:

● The degradation rate of PhNO can be expressed as2

a pseudo-first-order reaction with respect to PhNO2

concentration. This reaction rate becomes insensitiveto H O concentration whenwH O x )13.24 mM.2 2 2 2 o

The effect of high H O concentration is attributed2 2

to the fact that this reagent competes with PhNO2

wEq. (7)x for the HO radicalswEq. (4)x, leading to a•

decrease in the HO radical steady-state concentra-•

tion, and consequently in the PhNO degradation2

rate.● Identification of nitrophenols as reaction intermedi-

ates suggests that the degradation of PhNO takes2

place by reaction with HO radicals as the oxidant•

species. The kinetic results and initial product distri-bution under a wide range of experimental conditionssupport the presence of HO radicals in this process.•

● Although the chemistry of Fenton’s systems involvesa rather complex mechanism, an attempt has beenmade to fit experimental results to a simplistic kineticmodel and empirical kinetic equationwEq. (17)x.According to the results, this equation quantitativelypredicts the PhNO degradation rate under a wide2

range of experimental conditions with sufficient pre-cision. Furthermore, our study provides additionalinsight into mechanistic and kinetic factors control-ling the PhNO degradation rate in a ‘Fenton-like’2

process.

Acknowledgments

M. Rodriguez expresses his gratitude to the ULA-CONICIT collaboration(Venezuela) for financial sup-port and V. Timokhin thanks the Ministerio de Educationy Cultura(Spain) for a NATO Science Fellowship. Theauthors wish to express their gratitude for the financialsupport given by the Ministerio de Ciencia y Tecnologıa´(Spain, Project AMB 99-0442). We gratefully acknowl-edge Profs M.L. Kremer(Hebrew University, Israel),M. Bekbolet (Bogazici University, Turkey) and N.Restrepo(Universidad del Valle, Cali, Colombia) fortheir helpful discussions during the preparation of thispaper.

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