Effect of Chloride and Sulfate Ions on the

Effect of Chloride and Sulfate Ions on the Photoreduction Rate of Ferric Ion in UV Reactor Equipped with a Low Pressure Mercury Lamp Truong Giang Le1, Ngoc Tung Nguyen1, Quang Trung Nguyen2, Joseph De Laat3, Hai Yen Dao*, 1 1

Institue of Chemistry, VAST, 18 Hoang Quoc Viet Road, Hanoi, Vietnam Institute of Environmental Technology, 18 Hoang Quoc Viet Road, Hanoi, Vietnam 3 Université de Poitiers, Laboratoire de Chimie de l’Eau et Microbiologie de l’Eau (CNRS UMR 6008), Ecole Nationale Supérieure d’Ingénieurs de Poitiers, 40 Avenue du Recteur Pineau, 86 022 Poitiers Cedex, France 2

Abstract: This study aims to demonstrate the effects of pH and of inorganic anions on the photoreduction of Fe(III). Effective quantum yields for the production of OH radicals from iron(III) hydroxo species in pechlorate medium and in the present of chloride and sulfate anions were determined in the pH range 1-3. Experiments were carried out in batch at 25°C with a low-pressure mercury vapor lamp emitted at 253.7 nm. The method was based on measuring the pseudo-first-order rate constant of the photoreduction of Fe(III)-complexes in which tert-butanol scavenged the OH at an identical rate. The results showed that the apparent quantum yield of photoreduction of Fe(III) increased when the pH increases, the values passed from 0.063 at pH 1 to 0.1 at pH 3 ([Fe(III)]0 = 1 mM and 3 mM). The effects of Cl- or SO42- on the photoreduction of Fe(III) depended on pH and on ion concentrations. The apparent quantum yield values increased in the present of anion Cl -, whereas those values decreased in the present of anion SO42-. For the Photo Fenton system, the effects of Cl - or SO42- were found to depend on pH, on the concentrations of the inorganic anions and to decrease the rate of decomposition of H2O2. Keywords: photoreduction, quantum yield, Fe(III)-complexes, tert-butanol, Cl-, SO42-

Introduction Of the advanced oxidation processes (AOPs), the Fenton’s reaction (Fe(II)/H2O2) is an effective method of organic pollutant oxidation. In the absence of light and complexing ligands other than water, the most accepted mechanism of decomposition of H 2O2 by ferrous and ferric ions in acid homogeneous aqueous solution, involves the formation of hydroxyperoxyl (HO2/O2-) and hydroxyl radicals OH (1). It is now well-known that the combination of the Fenton’s reaction with UV light (also called as the Photo-Fenton process) largely enhances the rate of oxidation organic compounds (2, 3). The enhancement of reaction rates is due to the photoreductions of Fe(III) species which lead to a regeneration of Fe(II) and to the production of hydroxyl radicals OH. The rate of oxidation of organic pollutants by the photo-Fenton process depends on various parameters such as UV intensity, irradiation wavelength, quantum yields of photoreduction of Fe(III) species, concentrations of reactants, the reactivity of OH radicals with the organic and inorganic solutes present in the *Corresponding author; E-mail address: [email protected] ISSN 1203-8407 © 2014 Science & Technology Network, Inc.

water and pH of solution – a condition often explained as a dependence on specific iron species. Table B.1 and Figure B.1 give the values of the equilibrium constants for the formation of dissolved iron(III) species and the distribution of dissolved species as a function of pH, respectively. It has been reported that inorganic ions like chloride and sulfate ions may decrease the efficiency of advanced oxidation processes because hydroxyl radicals can be scavenged by chloride and sulfate to form less reactive inorganic radicals (4-7). In the case of the dark Fenton reaction, chloride and sulfate ions form complexes with iron(III) complexes (Table B.2 and Figure B.2) (8, 9). De Laat et al. (8-10) pointed out that chloride and sulfate also decrease the overall reaction rates of degradation of organic compounds by the Fenton reaction because unreactive chloro- and sulfatocomplexes are unreactive toward H2O2. In the case of the photo-Fenton reaction, the presence of chloride and sulfate ions may also affect the overall reaction rates because chloro and sulfatoiron(III)-complexes have not the same spectra with free iron(III) species (Figure 1 and Table B.3). J. Adv. Oxid. Technol. Vol. 17, No. 2, 2014

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Photoreactor

6000 -1 -1 e (M .cm )

253.7nm

4000

The photo-reactor consists of a cylindrical vessel with the low pressure mercury vapor lamp placed in axial position with quartz sleeve protection (Figure A.1 and Table A.1 in the Supporting Information). The target solution was fed to the reactor, mixed well and saturated with N 2 in order to avoid formation of peroxyl radicals. The temperature was kept at 25.0  0.5 °C. Before starting of experiment, UV lamp was turned on during at least 20 minutes to reach stable irradiation. Through the process, samples were withdrawn at timed intervals.

313 nm

Fe(O H) 2+

Fe(SO 4 )2 FeSO 4 + FeCl 2 +

2000

FeCl 2+

Fe 3+

0 220

270

320 370 Wavelength (nm)

420

Chemical Actinometry

Figure 1. UV/visible spectra of Fe(III) complexes.

It has also been observed that UV irradiation of chloro and sulfato complexes leads to the formation of ferrous ion, of hydroxyl radicals and of other inorganic radicals such as chlorine atoms and sulfate radicals via direct photolysis reactions or by indirect reactions between hydroxyl radicals and chloride or sulfate ions (4-7, 11, 12).

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N

Commercially available high-purity Fe(ClO4)3., 9 H2O (Aldrich, low chloride), HClO 4 (Riedel-de Haën, ACS reagent), NaClO4 (Sigma) and NaOH (Labosi, Analynorm 1 M) were used without further purification. All aqueous solutions were prepared with purified water using Milli-Q water. Solution of iron(III) were prepared in the following way: the required amount of Fe(ClO 4)3 was dissolved in acidified water. The ionic strength was adjusted to the desired value with NaClO4. The iron solutions were freshly prepared each time before use. The experimental conditions were controlled as follows: 25 °C, I = 0.2 M (adjusted with NaClO4), tert-Butanol: 53-106 mM; different concentrations of anions: (Cl-: 100 mM, 200 mM); (SO42- : 5 mM, 33 mM). pH: 1-3 (adjusted by HClO4). For each experiment, Tables E.1, E.2, E.3 reports the initial concentration of FeIII, NaCl, Na2SO4 ionic strength, the pH of solution, the fraction of FeIII present as hydroxocomplexes, chlorocomplexes and sulfatocomplexes. 306

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This study was designed to investigate the effects of chloride and sulfate concentrations on the overall rates of photoreduction of Fe(III), to compare the reaction rates with those obtained in perchlorate solutions and the efficiency of Photo-Fenton process.

Material and Methods Preparation of Solutions

The photonic fluxes emitted by the low pressure Hg lamp were determined by actinometry using H2O2 and ferrioxalate (6mM/ H2SO4 0.1 N) as actinometers (13). The intensity of the UV lamp was found to be equal to (4.52 ± 0.126) 10-6 E.s-1 (from H2O2 actinometry) and (5.14 ± 0.098) 10-6 E.s-1 (from ferrioxalate actinometry). (More detail in the Section C - Supporting information).

The concentration of hydrogen peroxide in stock solutions of H2O2 was determined by iodometric titration. Concentrations of hydrogen peroxide in solutions containing Fe(II) or Fe(III) were determined spectrophotometrically using the TiCl4 method (14) and a molar absorption coefficient of 724 M-1 cm-1 for the titanium peroxocomplex. Fe(II) was determined by photometric method using 1,10-orthophenanthroline (15) as the colorproducing reagent which can form a red complex that absorbs at 510 nm. The molar extinction coefficient of Fe(ophen)32+ was found to be e510 = 11050 L. mol–1 . cm–1. pH measurements were conducted using a pH meter PHM240. All the spectra were recorded on a SAFAS 190 DES spectrophotometer. Calculations: Distributions of ferric complexes were made with MINEQL+ software (16). Equilibrium constants were obtained from the literature (17) and corrected for differences in ionic strength.

Results and Discussion Photoreduction of Ferric Species: The Fe(III)/h/ tert-Butanol/ N2 System Photolysis of Fe(III) in the presence of tert-butanol and in the absence of dissolved oxygen represents an appropriate experimental method to determine the quantum yield of photoreduction of Fe(III), because

T. G. Le et al.

d [  OH ] = kFe(II) [Fe(II)][OH] dt with kFe(II) = 2.7 108 M-1 s-1,

d [  OH ] = kTert-butanol [tert-butanol][OH] dt with kTert-butanol = 6.8 108 M-1 s-1, A calculation indicates that more than 99 % of the hydroxyl radicals would be scavenged by tert-butanol under our experimental conditions.

Effect of pH on the Rate of Photoreduction of Aquahydroxocomplexes of Fe(III) To demonstrate the important role of pH in the photoreduction of Fe(III), a series of experiments were conducted at a given concentration of Fe(III) (1 or 3 mM) and at different pH values ( 0.7-3.0). The experiments were carried out in perchlorate medium with an ionic strength of 0.2 M, at 25 °C, 53 mM of tert-butanol. Under these conditions, nearly all the  OH radicals generated in solution were completely scavenged by tert-butanol. The production of Fe(II) at various times of irradiation was monitored (Figure 3a). Figure 3a showed that the formation rate of Fe(II) followed a zero-order reaction kinetic and increased when pH increased. Furthermore, experiments were carried out pushing the conversion of Fe(III) into Fe(II) up to 40%, without noticing any change in the zero-order rate of Fe2+ generation. Therefore, if Fe2+

[FeIII] : 3mM; pH 1

1.5 0.001

[[FeIII]: 3 mM; pH 1; [Tert -B] = 5- 200 mM

0.0008

1.0

[FeII], M

kapp (10-7 .M.s -1 )

the organic radicals resulting from the oxidation of tert-butanol do not react with Fe(III) and Fe(II) (18, 19). Base on its simple behavior, the Fe(III)/hν/tertbutanol/N2 system had been selected in this work to assess the photoreduction quantum yields of Fe(III) species. The aqueous solutions of Fe(III) were saturated with N2 in order to avoid formation of peroxyl radicals. Photochemical experiments conducted in the present work ([Fe(III)]0 = 3 mM, pH = 1.05) in which the concentrations of tert-butanol was varied (0-200 mM), showed that the production of Fe(II) followed a zero-order kinetics because the photons emitted at 254 nm were totally absorbed by Fe(III) and that the zero-order rate constants of formation of Fe(II) followed a typical saturation curve reaching a plateau for tert-butanol concentration higher than 25 mM (Figure 2 and Table D.1-Supporting information). By comparing the rates of reaction of OH radicals by Fe(II) and by tert-butanol,

0.0006 [Fe(II)] = f(t) 0.0004

0.5

0.0002

Without Tert-Butanol

0 0

2000

4000 Time (s) 6000

0.0 0

50

100 150 Tert-Butanol (mM)

200

Figure 2. Effect of tert-butanol concentration on the rate constant of the photoreduction of Fe(III).

was involved in reactions as a reactant, shifting from zero concentration to high conversion values should entail differences. Particularly, the zero order kinetics was not observed because the effective formation of Fe2+ started depending on the present amounts of ferrous ion. This meant that there was no evidence for reactions involving Fe2+ as a reactant. The apparent reaction rate constants of photoreduction of Fe(III) (kapp) can be determined by the zero-order kinetic expression: [Fe(II)]t - [Fe(II)]0 = [Fe(II)] = kapp.t = app.

I Ia t = app. 0 .t VT VT

where, VT represents the volume of solution and app, the apparent quantum yield of photoreduction of Fe(III). The kapp values, determined from the slopes of [Fe(II)] = f(t) straight lines were reported in Table E.1 – Supporting information. As expected, the rate of photoreduction of Fe(III) or the apparent quantum yields of photoreduction increased when the pH increased (Figure 3b). The values of app calculated with I0 = 4.52 10-6 E.s-1 (determined from actinometry with H2O2) rose from 0.061 at pH ≤ 1 (1.41 10 -7 M.s-1) to 0.098 at pH  3. By using the value of I0 determined from ferrioxalate actinometry (I0 = 5.14 10-6 E.s-1), the values of app ranged from 0.054 at pH ≤ 1 to 0.088 at pH  3. The change of app with pH can be explained by the fact that the various aquahydroxocomplexes of Fe(III) have different photoreactivities. As illustrated in the Figure B.1 (Supporting information), Fe3+ is the most predominant monomeric of Fe(III) (more than 98% of Fe(III)) at pH < 1. FeOH 2+ represents less than 2% of Fe(III) at pH = 1 and approximately 50% at pH 3. J. Adv. Oxid. Technol. Vol. 17, No. 2, 2014

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0.15

k app = f(pH)

3.0

Molar fraction of Fe(III) species

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0.5

2.5

0.4

0.13

pH 2 .4 7 p H 1. 9 5

0.3

p H 1. 5 8 p H 1. 1

0.2

[FeII] : 1mM

Φapp

[Fe(II)], mM

kapp (10-7 .M. s -1 )

pH 2 .9 4

0.1

2.0

Time (s)

0 0

1000

0.11

1.0 Fe 3+

0.8 0.6 0.4

1.5

0.0

1

pH

1.5

2

2.5

3

Calculated with Io = 4.52.10-6

0.07

[Fe(III)] = 1mM

Fe(O H)2 + [Fe 2 (O H)2 ] 4+

0.2

0.09

2000

FeO H 2+

Calculated with Io = 5.14.10-6

[Fe(III)] = 3 mM

0.05

1.0 0.5

1.0

1.5 pH 2.0

2.5

0.5

3.0

1.0

1.5

pH

2.0

2.5

3.0

Figure 3. Rate constants and quantum yields of photoreduction of Fe(III) as a function of pH. 0.8

0.5 III

0.4 FeII (mM)

Fe II (mM)

[Cl-] = 100 mM

0.6

0.4 pH 2.85 pH 2.56 pH 2.20 pH 1.61 pH 1.16

0.2

0.0 0

500

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[FeIII] =1 mM

[Fe ] = 1 mM;

1000 1500 Time (s)

2000

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[SO42-]= 33 mM

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0.3 0.2

pH 1.15 pH 2.90 pH 1.65 pH 2.65 pH 2.05

0.1 0

0

2000

4000 Time(s)

6000

8000

Figure 4. Effect of pH on the formation of ferrous ion in the presence of chloride ion (100 mM) and sulfate ion (33 mM) Experimental conditions given in table E2, E3.

Effect of pH on the Rate of Photoreduction of Fe(III) in the Presence of Chloride and Sulfate

increases whereas the rate decreases in the presence of sulfate ions,

The rates of photoreduction of Fe(III) (1 mM) in the presence of perchlorate (200 mM), chloride (100 and 200 mM), and sulfate ions (5 and 33 mM) have been also compared. All the experiments were carried out with tert-butanol (53 and 106 mM). The data show that the formation rate of Fe(II) in the presence of chloride and sulfate ions also followed zero-order kinetics. For each experiment, the apparent rate constant was reported in Figure 4 and Tables E.2, E.3. Figure 5 also give the values of the apparent quantum yields of photoreduction of Fe(III) which were calculated by using the values of Io determined with H2O2 actinometry (I0 = 4.52 10-6 E s-1) and with ferrioxalate actinometry (I0 = 5.14 10-6 E s-1). Under the experimental conditions used in this work, the data show that: - the formation rate of Fe(II) in the presence of perchlorate and chloride increases when the pH

- the formation rate of Fe(II) is faster in the presence of chloride than in the presence of perchlorate, and slower in the presence of sulfate than in the presence of perchlorate,

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- the rate of formation of Fe(II) at a given pH depends on the concentration of chloride or of sulfate ions because an increase in the concentration of chloride or of sulfate enhances the concentration of chloro- or of sulfato-iron(III) complexes (Figures 6 and 7). Figure 6 showed that the apparent rate constant of photoreduction of Fe(III) decreased when the fraction of Fe(III)-chlorocomplexes increased that is to say when the pH of solution decreased. In the case of sulfate ions, a decrease of the apparent rate constant of photoreduction of Fe(III) was also observed when the fraction of Fe(III)-sulfatocomplexes increased (Figure 7).

T. G. Le et al. 5.0

Cl- : 100 mM

0.17

Cl- : 100 mM

Cl- : 200 mM Perchlorate medium

SO4 2- : 33mM

0.13

SO4 2- : 5mM

Φ app

kapp (10-7 .M. s -1 )

0.13

SO4 2- : 5mM

3.0

Cl- : 100 mM

0.17

SO4 2- : 33mM

Caculated with Io Feox = 5.14 10-6

SO4 2- : 5mM

Φ app

Cl : 200 mM Perchlorate medium

4.0

Caculated with Io H2O2 = 4.52.10-6

Cl- : 200 mM Perchlorate medium

-

SO4 2- : 33mM

0.09

0.09

0.05

0.05

2.0 1.0 0.01

0.01

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1.0

1.5

pH

2.0

2.5

3.0

5b) Values of  app determined with Io = 4.52 10-6 E s-1 (H2O2 actinometry).

5a) Apparent rate constants of photoreduction of Fe(III). Io = 4.52 10-6 E s-1 (H2O2 actinometry).

0.5

1.0

1.5

pH

2.0

2.5

3.0

5c) Values of  app determined with Io = 5.14 10-6 E s-1 (Ferrioxalate actinometry).

Figure 5. Change in the apparent rate constants and quantum yields of photoreduction of Fe(III) as a function of pH and in the presence of chloride (100, 200 mM), perchlorate (200 mM) and sulfate (5, 33 mM) [FeIII]= 1mM; [tert-butanol] = 53-106 mM; I = 0.2-0.3 M.

5

100 [FeIII] : 1mM

[Fe III] : 1mM

-

Σ Fe free+ hydroxocomplexes (Cl :100 mM)

4

80 kapp 10-7 s -1

Σ FeCl2+ + FeCl2 + (Cl : 200 mM)

%

60 40

3 2 Cl- : 100mM

Σ Fe free + hydroxocomplexes (Cl-:200 mM)

20

2+

1

Cl- : 200mM

0

Cl- : 100mM

+

Σ FeCl + FeCl2 (Cl : 100 mM)

0

20 30 40 50 60 Cl- : 200mM (FeCl2+ + FeCl2+ ) % Linearof(Cl: 6a) Speciation of Fe(III) as a function of pH 6b) kapp = f(fraction Fe(III) complexed by Cl-) 100mM) (Chloride ions : 100 or 200 mM) Linear (Cl- : Figure 6. Effect of chloride ions on the speciation of Fe(III) and change of k app values as a function of the fraction of Fe(III) complexed 200mM)

1

1.5

2 pH

2.5

3

0

10

by chloride.

Effect of Chloride Concentrations at pH  1.1 To clarify the effects of chloride concentration on the reaction rate, another series of experiments was conducted by changing the concentration of chloride at pH  1.1. Experiments were carried out at pH  1.1 in order to have only Fe3+ ion as the unique form of free Fe(III). The fraction of iron(III) present as of iron(III)-chlorocomplexes increases from 0% to 60% when the concentration of chloride ions increases from 0 to 200 mM. The data showed that an increases in concentrations of chloride (in the range 0-100 mM) led to an increase of the rate constant of photoreduction of Fe(III). But when the concentrations of chloride were higher than 150 mM, the rate constants decreased slightly (Figure 8, Table E.4). It is difficult to explain the effect of chloride ions on the rate of photoreduction of Fe(III) because the addition of chloride ions initiates many reactions:

- formation of Fe(III)-chlorocomplexes and these chlorocomplexes have not the same UV/Visible spectra as the spectra of free Fe(III) specie. These chlorocomplexes may also be more or less photoreactive than free Fe(III) species, - formation of chlorine atoms from the photolysis of Fe(III)-chlorocomplexes, - formation of chlorine atoms and of dichlorine anion radicals as secondary radicals by reaction of hydroxyl radicals with chloride ion. These inorganic radicals are less reactive than hydroxyl radical toward tertbutanol (Table 1). Nadtochenko and Kiwi (20) studied the laser photolysis at 347 nm of solutions of Fe (III) ( 1 mM) with chloride (2 M). They observed the formation of Cl2-, explained that Cl2- is formed via the reaction of HO with Cl- and they also showed that both FeCl 2+ and FeCl2+ chlorocomplexes can lead to the formation of Cl• atoms: J. Adv. Oxid. Technol. Vol. 17, No. 2, 2014

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80

FeSO4 + ( SO4 2- : 33mM)

kapp 10-7 s -1

% Complexes

60

[FeIII ] : 1mM pH: 1.0 - 2.8

FeSO4 + ( SO4 2- : 5 mM)

40 Fe(SO4 )2 - (SO4 2- : 33mM)

20

2[SO Series5 4 ] : 33 mM

3

2Series6 [SO 4 ] : 5 mM

2

Poly. (Series6) Linear (Series5)

1 -

2-

Fe(SO4 )2 (SO4 : 5 mM)

0

0 0.5

1.0

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2.0

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0

pH

20

40 60 80 (FeSO4 + + Fe(SO4 )2 - ) %

100

7b) kapp = f(fraction of FeIII complexed by SO42-)

7a) Speciation of Fe(III) as a function of pH (Sulfate ions: 5 or 33 mM).

Figure 7. Effect of sulfate ions on the speciation of Fe(III) and change of k app values as a function of the fraction of Fe(III) complexed by sulfate ions.

k app 10-7 s -1

2.5 200 2 150 1.5 Fe3+

0.5

50

FeCl 2+

FeCl2 +

0 0

50

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100 150 [Cl ], mM

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1

100

N

0 200

Figure 8. Rate constants of photoreduction of ferric and dependence of the speciation of Fe(III) as a function of the concentration of chloride ion (pH  1.1). h

FeCl2+  [Fe2+…Cl]  Fe2+ + Cl h

FeCl2+  [FeCl+…Cl]  FeCl+ + Cl David and David (21) studied the photolysis of ferric perchlorate or ferric chloride with •OH radical scavengers (methyl methacrylate, benzoic acid) and chloride. They observed a decrease of quantum yield when the concentration of chloride ions increased ([Cl-] = 0-1 M, pH 2). They concluded that photolysis of Fe(III)-chlorocomplexes did not lead to the formation of chlorine atoms. Machulek et al. (4) studied the photolysis of Fe(III) at pH 1.0-3.3 and at 355 nm with increasing concentrations of chloride ions (0.05-0.75 M). The change in the concentration of Cl2- was followed by spectrophotometric measurements at 340 nm. They concluded that in the FeIII/Cl-/UV system, chlorine 310


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atoms were formed either directly by the photolysis of chlorocomplexes Fe(III) or indirectly from HOClradicals formed by reaction of OH with Cl-. Because of the complexity of the reaction pathway and uncertainty of many constants, the individual quantum yields of photoreduction of Fe(III) chlorocomplexes will be explained more detail in the next paper concerning quantum yield of each species and modeling.

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Effect of Sulfate Concentrations at pH  1 Additional experiments at pH  1 were also carried out for the Fe(III)/SO42-/UV system. Experiments were carried out with solutions containing 1 mM of FeIII, tert-butanol (53-106 mM), and with concentrations of sulfate ions ranging from 0 to 50 mM (Table E.5). As shown in Figure 9, the apparent rate of photoreduction of Fe(II) decreased from 1.463 10 -7 M-1 s-1 to 1.041 10-7 M-1 s-1 when then concentration of sulfate increased from 0 to 50 mM or when the total fraction of Fe(III) sulfate as sulfatocomplexes increased from 0 to 87.5%. The decrease in the rate of photoreduction of Fe(III) with increasing concentrations of sulfate ions suggested that Fe(III)-sulfatocomplexes are less photoreactive than free Fe(III) species. These data are consistent with previous works. Benkelberg and Warneck (11) studied the oxidation of benzene (2 mM) by photoreduction of ferric perchlorate (0.77 mM) with sodium sulfate (10 mM) at pH 2 (I = 0.035 M), and tert-butanol (0.2 M) as a scavenger of OH (I = 0.035 M). Under these conditions, 89, 2.3 and 8.7% of Fe (III) are in the forms of FeSO4+, FeOH2+ and Fe3+, are respectively. From the rates of formation of

T. G. Le et al. Table 1. Formation and reactivity of Cl/Cl2- and SO4- radicals

Reaction

Rate constant (M-1 s-1)

Reference

1

Cl- + HO  ClOH-

k+ = 4.3 (± 0.4) 109 k-3 = 6.1 (± 0.8) 109

(24)

2

ClOH- + H+  Cl + H2O

k+ = 2.6 (± 0.6) 1010 k- = 3.6 (± 0.4) 1010

(12)

3

Cl + Cl- Cl2-

k+ = 7.8 (± 0.8) 109 k- = 5.2 (± 0.3) 104

4

Cl2- +

-

H2O  ClOH + H + Cl +



5

Cl + tert-butanol  Products

6

Cl2- + tert-butanol  Products

7 8

-



OH + tert-butanol  Products -



-

HSO4 + OH SO4 + H2O -



1.3 10

(17)

3.0 10

8

(25)

2

7 10 (pH 1) 6.8 10

8

3.5 10

5

9

SO4 + H2O  H + SO4 + OH

6.6 102

10

SO4- + OH-  SO42- + OH

1.4 107

-

+

SO4-

 S2O8

2-

2-

11

SO4 +

12

SO4- + H2O2  SO42- + H+ + HO2

13

SO4- + tert-butanol   CH2C(CH3)2OH + SO42- + H+

phenol, the absolute SO42- quantum yields were determined and found to rise from 0.0016 at 350 nm toward 0.0079 at 280 nm. These quantum yield values were about 40 to 60 times smaller than the quantum yields of photoreduction FeOH 2+. According to the authors, this large difference was attributed to the fact that the Coulomb attraction between Fe2+ and SO4are much stronger than between Fe2+ and OH. Therefore, it is more difficult for SO4- to leave the solvent cage than for OH.

Photo-Fenton: Effect of Chloride and Sulphate Ions on the Degradation of H2O2 In order to confirm the above result concerning the activities of Fe(III) species or Fe(III)-complexes, a series of experiments of decomposition of H 2O2 in the presence of hydroxyl radicals scavenger ([tert-butanol] = 106 mM) was carried out. Under these experimental conditions, all released hydroxyl radicals are trapped by tert-butanol and ferrous ions released into the medium react with H2O2. The experimental conditions and results were presented in (Figure 10 and Table E.6). The results showed that the rate of decomposition of H2O2 (1mM) was faster in the presence of chloride and slower in the presence of sulfate (Figure 10a). In all cases, the rate of decomposition of H 2O2 followed a zero-order rate law.

(12)

3

(26) (22)

(17, 27)

2.7 108 1.2 107 8.0 105 – 9.1 105

(28)

In the absence of hydroxyl radicals scavenger, the results in Figure 11 shows that the rate of decomposition of H2O2 was faster in perchlorate medium and slower in the presence of chloride ions and sulfate ions. Under the condition used, the rate of the decomposition of H2O2 was first-order with respect to H2O2. And the rate of decomposition of H2O2 by Fe system (III)/H2O2/UV was much faster than the system Fe (III)/H2O2 (Table 2). The data demonstrated that, the effect of chloride and sulfate anions on the rates of decomposition of H2O2 with, and without tert-butanol were not identical. These differences were not well explained because the presence of chloride and sulfate ions in medium leads to many reactions involving Cl•-, Cl2•- and SO4•radicals. Therefore, a modeling study is needed in order to assess these possibilities. Additional research is in progress in order to predict the effects of chloride and sulfate ions on the overall rates of decomposition of H2O2 by Fe(III)/ H2O2/UV system.

Effects of Chloride and Sulfate Ions on the Decomposition of Diethylene Glycol by Fe(III)/H2O2/UV In order to examine the effects of chloride and sulfate ions on the efficiency of photo Fenton, a competitive oxidation experiment of DEG by Fe(III)/ J. Adv. Oxid. Technol. Vol. 17, No. 2, 2014

311

1.5

150

1.2

120

kapp

0.9

90 60 30

Fe(SO4 )2 -

0.0 0

20

40 60 % Total complexes

80

% Complexes

0.6

150

kapp 1.2

120

0.9

Fe3+

90

FeSO4+

0.6

60

0.3

30 Fe(SO4)2-

0.0

0 100

0 0

9a) kapp values as a function of sulfate concentration.

% Fer. species

FeSO4 +

0.3

1.5

kapp(10-7 .s -1 )

kapp (10-7 .s -1 )

T. G. Le et al.

10

20

30 40 [SO4 2-]0 (mM)

50

9b) kapp values as a function of the fraction of Fe(III) complexed by sulfate.

Figure 9. Rate constants of photoreduction of Fe(III) and dependence of the speciation of Fe(III) as a function of the concentration of sulfate ion (pH  1.1).

18

kapp (10-7 s -1 )

H2 O2 (mM)

0.8 0.6 SO4 2- : 33 mM

0.4

t o

-

ClO4 : 100 mM

0.2 Cl- : 100 mM

0.0 0

1000

2000 Time(s)

o D 3000

y p

[H2 O2 ] = 1 mM ;[FeIII] = 1mM 15 [Tert -butanol] = 106 mM

[FeIII] : 1mM; [H2 O2 ] : 1mM; pH: 2; Tert-Butanol : 106 mM

1.0

N

4000

10a) Decomposition of H2O2 by Fe(III)/H2O2/UV ([H2O2]0 = 1mM, O2 < 0.1 mg/L)

o C

12 9

-

[Cl ] = 100 mM

6 3

[ClO4 -] = 100 mM 2-

[SO4 ] = 33.3 mM

0 1

1.5

2

pH

2.5

3

10b) Plot of the zero-order rate constant as a function of pH

Figure 10. Decomposition of H2O2 by FeIII/UV/H2O2 in the presence of hydroxyl radical scavenger and in the absence of oxygen.

UV/H2O2 was performed. DEG has been chosen because it is an aliphatic and nonphotolyzable compound. Its degradation was only initiated by •OH radical. The rate constant of the reaction of •OH with DEG is relatively high 2.1 109 M-1s-1 (22). The experiments were conducted at pH = 3, T = 25 ± 1°C, [DEG]0 = 2.08 mM, [Fe(III)]0 = 0.2 mM, [H2O2]0 = 20 mM, I = 0.2 M. All experiments were carried out in the presence of dissolved oxygen to promote the formation of peroxyl radicals ROO• and the oxidation of DEG to CO2. During the reaction, the concentrations of H2O2, Chemical Oxygen Demand ([COD]0 = 330 mg O2/L) and the absorbance at 254 nm were determined (Figure 12). At the end of the reaction, the residual total organic carbon ([TOC]0 = 100 mg C/L) was measured to confirm the values of COD (Table 3). 312


The results (Figs. 12a-12c) show that the rates of decomposition of H2O2 in perchlorate medium and chloride medium (0-100 mM) were similar in the presence of DEG. Whereas, the rate was slower in the presence of sulfate. End of the reaction, analyses showed that H2O2 was completely depleted in samples contained perchlorate and chloride ions and that the H2O2 removal was nearly equal to 65% in the presence of sulfate. Mineralization rates of DEG in perchlorate medium or in the presence of a low concentration of chloride (1 mM) are in the same order of magnitude. For a chloride concentration of 100 mM, the rate was slower and the oxidation stopped when the concentration of H2O2 residual becomes zero. In the presence of sulfate (33.33 mM), the rate of oxidation of DEG is similar to that observed in the presence of 100 mM of chloride

T. G. Le et al. Table 2. Decomposition of H2O2 by Fe (III)/H2O2/UV and Fe (III)/H2O2 in the absence of hydroxyl radical scavenger

FeIII (mM)

H2O2 (mM)

pH

I (M)

Anion

[Anion] (mM)

% Fe Complexes

% Fe Free

3 3 3

20 20 20

1.15 1.09 1.12

0.2 0.2 0.2

ClO4ClSO42-

100 100 33

0 79.2 95.8

100 21.8 4.2

k’app Photo- Fenton (M-1. s-1) 4.52 10- 4 3.23 10- 4 1.81 10- 4

kapp Fenton (M-1. s-1) 5.73 10-5 3.03 10-5 1.61 10-5

k’app / kapp 5.8 10.1 11.2

Table 3: The percentage of Fe(III) complexes, COD and TOC values at the end of reaction time

No

Anion

[Anion] (mM)

%Fe(III) Complexes

COD end of experiment (mg O2 / L)

TOC end of experiment (mg C / L)

D1 D2 D3 D4

Perchlorate Chloride Chloride Sulfate

200 10 100 33.33

0 1.33 12.96 69.35

3.2 (99 %) 23.1 (93 %) 114.7 (65 %) 72.2 (78 %)

3.3 (97 %) 4.9 (95 %) 44.1 (55 %) 14.3 (86 %)

0.020

-

0.015

Cl : 100 mM -

ClO4 : 100 mM

0.010 FeIII/H2 O2 / UV : 2-

0.005

SO4 : 33 mM -

Cl : 100 mM

Ln ([H2 O2 ]t/[H2O2 ]o)

SO4 2- : 33 mM

[H2O2 ] (M)

0.0

Without UV :

SO4 2- : 33 mM

-0.2

Cl- : 100 mM

Without UV :

ClO4 - : 100 mM

-0.4 FeIII/H2 O2 / UV :

-0.6

SO4 2- : 33 mM Cl- : 100 mM

-0.8

ClO4 - : 100 mM

-1.0

ClO4 - : 100 mM

0.000

-1.2 0

2000

4000 6000 Time (s)

8000

11a) [H2O2] = f(t)

0

1000

2000 3000 Time (s)

4000

5000

11b) ln ([H2O2]t/ [H2O2]o) = f(t)

Figure 11. Effect of sulfate and chloride ions on the rate of decomposition of H 2O2 by FeIII/H2O2 and FeIII/UV/H2O2 in the absence of hydroxyl radical scavenger at pH 1; [Fe III] = 3 mM.

but it continues throughout the duration of the experiment because the H2O2 residual is still high.. The COD and TOC removal at the end of the reaction time are consistent (Table 3). The UV absorbance at 254 nm decreases during the reaction due to the decrease of the concentration of H2O2 (H2O2 e = 18.6 M-1 cm-1 at 254 nm). It should be noted that the oxidation of DEG to CO2 potentially pass through the formation of oxalic acid (23). The evolution of the UV absorbance showed no accumulation of oxalic acid. Because it would lead to a significant increase in absorbance at 254 nm due to the formation of ions ferrioxalate. In conclusion, the results showed that the rate of the oxidation of DEG is slower in the presence of sulfate. The decrease in the rate of degradation of DEG could be mainly attributed to the formation of

sulfate complexes of iron (III) ( 70%, Table 3) that has lower photo-reactivity than hydroxocomplexes of iron (III). If we compare the efficiency of the photo Fenton process by examining the relationship between COD removal and H2O2 consumed, the results show that the performance of the process in sulfate and perchlorate medium are identical. These results are consistent with the fact that sulfate radicals formed during the photoreduction of Fe (III) is completely changed to OH radicals via the reaction 10 Table1. Whereas, the rate of degradation of DEG decreased in the presence of high concentration of chloride ions. These results could be explained by the fact that the radical Cl2•- represents the dominant radical in the reaction medium. As radicals Cl2•- are less reactive than the •OH radicals, a considerable fraction of radicals Cl2•- reacted with other species J. Adv. Oxid. Technol. Vol. 17, No. 2, 2014

313

T. G. Le et al. 350

Perchlorate Sulfate 33,33 mM Chloride 10 mM Chloride 100 mM

0.020

COD (mgO2 /L)

0.015

[H2 O2] (M)

Perchlorate Sulfate 33,33 mM Chloride 10 mM Chloride 100 mM

300

0.010

250 200 150 100

0.005

50 0.000

0 0

1000

2000

3000

4000

5000

6000

0

1000

2000

0.9

350

0.8

300

0.7 0.6 Perchlorate 0.5

Sulfate 33,33 mM

0.4

Chloride 10 mM Chloride 100 mM

1000

2000

250 200

o C

150 100

t o

0.3 0

3000

3000

4000

5000

Time (s)

COD removal (mg/l)

Absorbance at 254nm (1 cm-cell)

Time(s)

4000

Time (s)

5000

N

y p

Perchlorate Sulfate (33,33 mM) Chloride (100 mM)

50

0

0

5

10

15

20

25

H2O2 consumed (mM)

Figure 12. Evolution of H2O2 concentrations, COD and absorbance at 254 nm observed in the oxidation of diethylene glycol ([DEG] = 2.08 mM) by Fe (III)/H2O2/UV in the presence of O2.

o D

such as H2O2 and Fe2+ ions. These side reactions may explain the lower efficiency of the photo Fenton process when we observed the relation between COD removal and H2O2 consumed (Figure 12d).

Conclusion

In the first part of this study, the experiments of photoreduction of Fe(III) in perchlorate medium at 253.7 nm was given. While taking the photonic flux value determined by chemical actinometry of hydrogen peroxide, the experimental results showed that the rate of photoreduction of ferric iron (pH < 3) depends on the pH, the concentration of chloride and sulfate ions. In perchlorate medium the rate of photoreduction increases significantly with increasing pH from 1 to 3, the values pass from 0.063 at pH 1 to 0.1 at  pH 3 ([Fe(III)]0 = 1 mM and 3mM). This increase in quantum yield is in good agreement with the bibliographical data. In the presence of chloride ions, the results showed a significant increase in the rate of photoreduction of 314


ferric iron, whereas a decrease was observed in the presence of sulfate ions. In the second part, the results showed that the effects of Cl- or SO42- on the degradation of H2O2 were found to depend on pH and on the concentrations of the inorganic anions and to decrease the rate of decomposition of H2O2. The reaction rate markedly decreased in the presence of sulfate and chloride but are approximately 4.5 to 11 times faster than that in absence of UV irradiation in the same experimental conditions employed. The results could be mainly attributed to the fact that the apparent quantum yield for the production of Fe(II) depends on the distribution of the Fe(III) complexes, to the reactivity of Fe(II) and Fe(III) species with soluble substrates (H2O2) and to the formation of inorganic radicals (Cl2- and SO4-) which are less reactive than hydroxyl radicals. In order to have more deepened interpretation of the experimental results, this study showed that some complementary studies are necessary in order to determine quantum yields of photoreduction of chloro

T. G. Le et al.

and sulfato complex of Fe(III) at 254 nm and at other wavelengths emitted by the low pressure mercury vapor lamps (> 254 nm) because the contribution of these secondary radiations to the photoreduction of Fe(III) is not negligible. The determination of quantum yields requires, on the other hand, a better knowledge of the absorption spectra of the species of Fe(III) (especially Fe(OH)2+ and dimer) and a judicious choice of the experimental method such as using the optical bench in the place of a low pressure mercury vapor lamp would also make it possible to determine quantum yields of photoreduction in a more precise way.

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Acknowledgements

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This work was financially supported by Viet Nam national foundation for science and technology development (NAFOSTED) with the project code “104.03.25.09” and A-water project (Daninda).

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(13) (14) (15)

(16) (17)

References (1) De Laat, J.; Gallard, H. Environ. Sci. & Technol. (2) (3) (4)

(5) (6) (7) (8) (9) (10) (11)

1999, 3, 2726-2732. Ruppert, G.; Bauer, R.; Heisler, G. J. Photochem. Photobiol. A. 1993, 73, 75-78. Pignatello, J.; Sun, Y. ACS Symp. Ser. 1993, 518, 77105. Machulek Jr., A.; Vautier-Giongo, C.; Moraes, J.E. F.; Nascimento, C.A.O.; Quina, F.H. Photochemestry and Photobiology 2006, 82, 208-212. Machulek et al. Environ Sci Technol. 2007, 41(24), 8459-63. Machulek et al. Photochemical & Photobiological Sciences 2009, 8, 985-981. Devi, L. et al. Journal of Molecular Catalysis A: Chemistry 2013, 374, 125-131. De Laat, J.; Le Truong, G. Environ Sci Technol. 2005, 39, 1811-1818. De Laat, j.; Le Truong, G. Applied Catalysis B. 2006, 66, 137-146. Le Truong, G.; De Laat, J.; Legube, B. Water research 2004, 38, 2384-2394. Benkelberg, H.J.; Warneck, P. J. Phys. Chem. 1995, 99, 5214-5221.

(20) (21) (22)

(23) (24) (25) (26) (27) (28)

Yu, X.Y.; Barker, J.R. J. Phys. Chem. A. 2003, 107, 1313-1324. Nicole, I.; De Laat, J.; Dore, M. Water Research. 1990, 24, 157-168. Eisenberg, G.M. Indus. Eng. Chem. 1943, 15, 327328. AFNOR NF T 90-101, 1988. Détermination de la demande en oxygène, (DCO) – Méthode par le bichromate de potassium. Schecher and McAvoy, W.D.; Schecher and D.C. Environ. Urban Syst. 1992, 16, 65-76. Martell, A.E.; Smith, R.M. Critical Stability Constants; Vol. 3, Plenum Press: New York, 1977. Carey J.H.; Langford C.H. Can. J. Chem. 1975, 53, 2436-2440. Lopes, L. PhD. Transfert de charge des aquahydroxycomplexes de Fe(III) : Mécanismes d’oxydation radicalaire de composés aliphatiques hydroxylés photoinduite par l’ion hexaaqua Fe(III), réactivité des radicaux peroxyle, rendements quantiques de photoréduction. Doctorat de l’Université de Poitiers, 2000. Nadtachenko, V.; Kiwi, J. Inorg. Chem. 1988, 37, 5233-5238. David, F.; David, P.G. J. Phys. Chem. 1976, 80, 579583. Buxton, G.V.; Greenstock, C.L.; Helman, W.P.; Ross, A.B. J. Phys. Chem. Ref. Data.1988, 17, 513886. McGinnis, B.D.; Adams, V.D.; Middlebrooks, E. J. Wat. Res, 2000, 34, 2346-2354. Jayson, G.G.; Parsons, B.J.; Swallow, A.J. Journal of Chem. Soc., Faraday Trans. 1973, 69, 1597. Mertens, R.; von Sonntag, C. J. Photochem. Photobiol. A. 1995, 85, 1-9. Hasegawa, K.; Neta, P. J. Phys. Chem. 1987, 82, 854-857. Neta, P.; Huie, R.E.; Ross, A.B. J. Phys. Chem. Ref. Data, 1988, 17, 1027-1247. Clifton, C.L.; Huie, R.E. Int. J. Chem. Kinet. 1989, 21, 677-687.

Received for review December 12, 2013. Revised manuscript received March 3, 2014. Accepted March 11, 2014.


315

T. G. Le et al.

SUPLEMENTARY MATERIAL Section A: Reactor used in the present work

N2

pH T NaOH HClO4

Solution à étudier Solution

y p

Lampe UV lampUV

Thermostated batch Water 25.0  0.2 °C

t o 0 1

N

o C

Magnetic Barreau Bar aimanté

AgitateurAgitator magnétique Magnetic

Figure A.1: UV reactor used in the present work

o D

Table A.1: Description of the UV reactor

Type of reactor

Batch reactor

Volume of irradiated solution (V, L)

1.95

Internal diameter of the reactor (cm)

9.4

Thickness of annular (cm)

3.45

UV lamp (low pressure Hg lamp) Temperature (° C)

Vilber Lourmat, T-6C, 6 watt 25.0 ± 0.1

- Recirculating pump for thermostated water (q  200 L/h; 25.0  0.5 °C). - Solution of sodium hydroxide or of perchloric acid - pH transmitter (OPM 223; Endress + Hauser) equipped with a pH sensor and a temperature sensor for pH and temperature measurements and temperature compensation.

316


T. G. Le et al.

Section B: Supporting information B. Introduction: Table B.1: Hydrolysis equilibrium constants for iron(II) and iron(III) in perchlorate medium at 25 oC Log Ko

Equilibria

Log K = f(I)

Reference

I = 0 ( 25 °C)

MINEQL+

2.04 I

Milburn et Vosburgh, 1955

(I = 0) -2.19

Fe3+aq  Fe(OH)2+ + H+aq

-2.17 Fe3+aq  Fe(OH)2+ + 2H+aq

-5.67

2Fe3+aq  Fe2(OH)24+ + 2 H+aq + 2H2O

 2.17   5.67 

-12

Fe + H2O  Fe OH + H II

+

Fe2+ + 2H2O  FeII(OH)20 + 2 H+

-20.6

Equilibria

I=0

Fe + H2O  FeOH + H 2+

+

Fe + 2 H2O  [Fe(OH)2] + 2H 3+

+

3+

2 Fe

+

+ 2 H2O  [Fe2(OH)2]

4+

Fe + H2O  [FeOH] + H 2+

+

+

Fe + 2H2O  Fe(OH)2 + 2H 2+

Turner et al. (1981)

MINEQL+

I = 0 ( 25 °C)  12 

3.066 I 1  3.81 I

 0.07I


I = 0 ( 25 °C)

MINEQL+

1.022 I

 0.01I


 0.34I


+

-9.5

3+

 0.016I

Knight et Sylva, 1975

- 9.5 2+

3.066 I 1  3.71 I

-3.09 -2.95

Fe3+aq  Fe(OH)3 + 3 H+aq

 0.01I

1  2.4 I

+

 9.5 

1  2.1 I

 20.6 

1.0221 I 1  2.01 I

I = 0.1 M I = 0.2 M I = 0.3M I = 0.4 M I = 0.5 M I = 1 M

-2.19

-2.63

-2.72

-2.77

-2.79

-2.79

-2.72

-5.67

-6.33

-6.47

-6.54

-6.56

-6.57

-6.47

-2.95

-2.95

-2.95

-2.95

-2.95

-2.95

-2.95

-9.50

-9.72

-9.77

-9.79

-9.8

-9.8

-9.77

-20.6

-20.8

-20.8

-20.9

-20.9

-20.9

-20.8


317

T. G. Le et al. 100

100 [Fe(III)] = 1 mM I = 0.2 M

3+

Fe 80

[Fe(III)] = 3 mM I = 0.2 M

3+

Fe 80

60 (%)

(%)

60

40

40

Fe(OH)2+ Fe2 (OH)2 4+

20

Fe(OH)2+

20

Fe(OH)2 +

Fe2 (OH)2 4+

0

Fe(OH)2 +

0

0

0.5

1

1.5

2

pH

2.5

3

0

0.5

a) [Fe (III)]T = 10-3 M and I = 0.2 M

1

1.5

pH

2

2.5

3

b) [Fe (III)]T = 3 10-3 M and I = 0.2 M

Figure B.1: Distribution of aquahydroxycomplexes of Fe(III) as a function of pH

y p

(Calculated with the constants provided by MINEQL+)

o C

Table B.2: Hydrolysis constants for the formation of iron(III) complexes with chloride and sulfate ions at 25 oC ( De Laat and Le, 2005, 2006)

t o

Reaction Fe + Cl  FeCl 3+

-

2+

Fe + 2Cl  3+

-

Fe(Cl)2+ +

o D

100

MINEQL+

0.37 2.32 3.28 1.18

MINEQL+ MINEQL+ MINEQL+ MINEQL+

% Complexes

FeCl2

+

40 2+

FeOH2+

0 0.2 0.3 [Cl-], M

FeSO4 +

60 40 Fe3+

Fe(SO4 )2 -

20

FeCl

0.1

0.79

80

60

0

MINEQL+

[FeIII] = 1 mM; pH = 1; [SO4 2-] = 0 - 200 mM

Fe3+

20

0.68

100

[FeIII] = 1 mM; pH = 1; [Cl-] = 0 - 500 mM

80

% Complexes

-

Reference

N

Fe + Cl  FeCl Fe3+ + SO42-  FeSO4+ Fe3+ + 2 SO42-  Fe(SO4)2Fe2+ + SO42-  FeSO4 2+

log K (I = 0.2 M)

0.4

a) Speciation of Fe(III) by chloride ions

0 0.5

0

0.04

0.08 0.12 2[SO4 ], M

0.16

b) Speciation of Fe(III) by sulfate ions

Figure B.2: Speciation of Fe(III) by chloride an sulfate ions at pH 1 ([Fe(III)]0 = 1mM) (Caculated by MINEQL+) 318


0.2

T. G. Le et al.

Table B.3: Molar extinction coefficients of iron(III) species (ε, M-1 cm-1). Wavelength (nm)

Fe3+

Fe(OH)2+

Fe(OH)2+

Fe2(OH)24+

200 210 220 230 240 250 254 260 280 300 313 320 330 350 380 390 400 440 480 500

3417 3250 3355 3805 4040 3710 3310 2800 830 135 40 19 6

5170 5140 4155 2850 1720 1065 983 1040 1830 2160 1905 1695 1350 703 194 75 45

3475 3655 3605 3935 4125 3980 3735 3445 2135 1605 1493 1465 1425 1335 1035 920 815 510 335 280

8715 10630 12130 13090 13975 14710 14880 14900 13590 11355 10505 10285 10080 9410 6390 5460 4490 2125 1195 825

Table B.4: Molar extinction coefficients and chloro and sulfatocomplexes of iron(III) (ε, M-1.cm-1) (De Laat and Le, 2005, 2006) Wavelength (nm)

FeCl2+

FeCl2+

FeSO4+

Fe(SO4)2-

200 210 220 230 240 250 254 260 280 300 313 320 330 350 380 400

6420 5082 5862 5400 4197 3119 2542 2050 886 830 1186 1409 1629 1355 346 62

2850 5272 5792 5800 4714 3388 2908 2399 1138 1149 1527 1782 2023 1793 685 286

3393 4061 4355 4275 3733 2870 2443 2071 1704 2140 2003 1804 1397 607 90 17

5933 4652 5001 4474 3586 2773 2426 2280 2724 3609 3548 3234 2500 1140 252 78


319

T. G. Le et al.

Results C. Determination of the photonic flux emitted by the UV lamp The photonic fluxes emitted by the low pressure Hg lamp used in all photochemical experiments described here were determined by actinometry using H2O2 and ferrioxalate (6mM/ H2SO4 0.1 N) as actinometers (Nicole et al., 1990). The first method was based on the photolysis of H2O2: h

H2O2  H2O + 0.5 O2 and the second one, on the photoreduction of ferrioxalate into iron(II) : h

2 Fe3+ + C2O42-  2 Fe2+ + 2 CO2 Figures 4 and Table 5 presents the results obtained from the decomposition of hydrogen peroxide with an initial concentration of H2O2 ranging from 50 µM to 100 mM. The values of incident photon flux (Table C.1) were calculated using the general equation of photolysis: Ln(10-Do – 1) = Ln(10-Dt – 1) – 2,302 e  H 2O 2  with

Do = e H 2 O 2  [H2O2]0 and Dt = e H 2 O 2  [H2O2]t k* = 2,302 e  H 2O 2 

I0 V

t o

N

y p

I0 t = Ln(10-Dt – 1) – k* t V

o C

For experiments in extremely concentrated solution with total absorption of photons (D o and Dt >> 2), the

o D

photon flux can be calculated from the following equation: [H2O2]0 – [H2O2]t =  H 2O 2

I0 t = k.t V

with

k =  H 2O 2

I0 V

For experiments conducted with high concentrations of H2O2 ([H2O2]0 = 100 mM), all photons are absorbed by the solution and 99 % of photons should be absorbed for an optical path length of 1.07 cm which is very lower than the annular path length of the photochemical reactor used (l = 3.45 cm). For these conditions, the zeroorder kinetic equation with respect to the H 2O2 concentration can be used to calculate I 0 (Figure C.1(b)). The results lead to a mean value of I 0 equal to 4.64 10-6 E s-1 (SD = 0.097 10-6) (Table C.1).

320


T. G. Le et al.

Table C.1: Determination of photon flux (I0) emitted by the lamp Lourmat Vibert, 6 watt pH

H2O2 V l Do Reaction (mM) (l) (cm) (t=0) time (s)

k (M-1.s-1) ([H2O2] =f(t))

Io* k* Io ln(10^Do -1) = ( calculation (calculation based on k) based on k*) f(t)

3.45 6.56 3.45 6.56

10500 12420

2.34 10-6 2.43 10-6

3.4510-4 3.6010-4

4.56 10-6 4.75 10-6

4.56 10-6 4.75 10-6

3.45 6.52 3.45 6.50

10800 7800

2.43 10-6 2.38 10-6

3.6010-4 3.5110-4

4.75 10-6 4.63 10-6

4.75 10-6 4.64 10-6

1.18 101.53 1.95 3.45 6.52 1.12 101.53 1.95 3.45 6.52

15900

2.42 10-6

3.5710-4

4.71 10-6

4.72 10-6

28920

2.31 10-6

3.4210-4

4.51 10-6

4.52 10-6

1.10 101.25 1.95 3.45 6.50 5.53 101.25 1.95 3.45 6.50

17400

2.32 10-6

3.4210-4

4.5210-6

4.52 10-6

10500 2.39 10-6 Mean Standard deviation

3.5210-4

4.6610-6 4.64 10-6 0.097 10-6

4.65 10-6 4.64 10-6 0.096 10-6

4.01 1.13

51.39 1.95 3.45 3.30 51.25 1.95 3.45 3.29 52.85 1.95 3.45 3.39

7200 5880

2.21 10-6 2.27 10-6

3.27 10-4 3.35 10-4

4.30 10-6 4.42 10-6

4.32 10-6 4.43 10-6

7500

2.24 10-6

3.32 10-4

4.37 10-6

4.38 10-6

51.53 1.95 3.45 3.31 49.70 1.95 3.45 3.19

5340

2.33 10-6

3.45 10-4

4.55 10-6

4.56 10-6

8220

2.28 10-6

3.38 10-4

4.45 10-6

4.47 10-6

9720

2.28 10-6

3.38 10-4

4.45 10-6

4.47 10-6

1.12 5.40

50.20 1.95 3.45 3.22 51.67 1.95 3.45 3.32 50.00 1.95 3.45 3.21

8100 9720

-6

2.30 10 2.29 10-6

-4

3.41 10 3.41 10-4

-6

4.49 10 4.47 10-6

4.51 10-6 4.50 10-6

5.40

50.00 1.95 3.45 3.21

9600

2.28 10-6

3.52 10-4

4.45 10-6

4.65 10-6

5.40 5.00

51.25 1.95 3.45 3.29 50.00 1.95 3.45 3.21

16800 6600

2.26 10-6 2.23 10-6

3.43 10-4 3.29 10-4

4.41 10-6 4.35 10-6

4.53 10-6 4.35 10-6

5.00

47.00 1.95 3.45 3.02

10980

2.27 10-6

3.36 10-4

4.43 10-6 4.43 10-6 0.066 10-6

4.44 10-6 4.47 10-6 0.094 10-6

4.23 102.23 1.95 5.10 102.22 1.95 5.15 101.53 1.95 1.09 101.25 1.95

1.08 1.09 1.07 1.15

Mean Standard deviation 5400

3.34 10-4

4.41 10-6

1.11

10.27 1.95 3.45 0.66 10.30 1.95 3.45 0.66

5400

3.34 10-4

4.41 10-6

5.20

10.00 4.00 6.25 1.16

21900

3.05 10-4

4.56 10-6 4.46 10-6 0.0860 10-6

1.10

Mean Standard deviation 2.1

5.12 1.95 3.45 0.33

2730

3.57 10-4

4.72 10-6

5.00

0.05 1.95 3.45 0.0032

7200

3.30 10-4

4.36 10-6

4.78

0.05 1.95 3.45 0.0032

8700

3.37 10-4

Mean

4.45 10-6 4.41 10-6

Standard deviation

0.063 10-6

Mean

4.52 10-6

Standard deviation

0.126 10-6


321

T. G. Le et al. 0.12

20

H2 O2 = 100 mM

[H2 O2 ] = 100 mM

0.10

10

[H2 O2] (M)

Ln(10^DO -1)

15 [H2 O2 ] = 50 mM

5 [H2 O2 ] = 10 mM

0

Simulated by GEPASI

0.08

Experiments

0.06 0.04

[H2 O2 ] = 5 mM

y = -2.3127E-06x

0.02

-5

R2 = 9.9894E-01

[H2 O2 ] = 0.05 mM

-10

0.00

0

2000

4000 6000 8000 Irradiation time (s)

10000

0

a) Photolysis of H2O2. Experimental results and simulation performed by Gepasi with I0 = 4.52 10-6 E s-1. 12 10 [H2 O2 ], mM

40

20 10 0 0

2000

o D

4000 6000 Time (s)

8000

N

10000

8

t o

Simulated by GEPASI Experiments

30

40000

y p

H2 O2 = 10 mM

H2 O2 : 50 µM

50

20000 30000 Time (s)

b) Photolysis of H2O2 ([H2O2]0 =100 mM). Experimental results and simulation performed by Gepasi with I0 = 4.52 10-6 E s-1.

60

[H2 O2 ], µM

10000

c) Photolysis of H2O2. (In diluted medium [H2O2]0 =50 µM). Experimental results and simulation performed by Gepasi with I0 = 4.52 10-6 E s-1).

6 4 2

o C

Simulated by Gepasi Experiments

0

0

5000

10000 15000 Time (s)

20000

d) Photolysis of H2O2.([H2O2]0 =10 mM). Experimental results and simulation performed by Gepasi with I0 = 4.52 10-6 E s-1).

Figure C.1: Determination of Io from photolysis experiments of H2O2 Figure C.1(a) shows that the general equation of photolysis of H2O2 can be used to describe reaction kinetics for all the initial concentrations used in this study, 50 µM to 100 mM, with percentages of photons emitted at 254 nm reaching the inner wall of the reactor between 0% and 99.3%. The application of the general equation leads to a mean value of I0 equal to 4.52 10-6 E s-1 (SD: 0126 10-6 E s-1) (Table C.1). Curves simulated by the software GEPASI correctly describe the experimental rates of decomposition of H 2O2 (Figures C.1 (b, c, d)). Determination of photon flux of UV light by chemical actinometry of Fe 2(C2O4)3 Figure C.2 shows some experimental results obtained from ferrioxalate actinometry. Follow the formation of Fe(II) during the UV irradiation of potassium ferrioxalate solutions (Fe2(C 2O4)3 : 6 mM, pH acid). These experiments led to a mean value of I 0 equal to 5.14 μE s-1 (SD: 0.098 μE s-1) by taking a value of 1.25 for the quantum yield of photoreduction of Fe2(C2O4)3 (Ferrioxalate) (Kuhn et al., 2004). [Fe2+]t – [Fe2+]0 = Ferrioxalate

322


I0 t = kappt V

T. G. Le et al. 0.005

[Fe(II)] (mM)

0.004 0.003 0.002 y = 3.33E-06x R2 = 9.99E-01

0.001 0.000 0

500

1000

1500

Time (s)

Figure C.2: Determination of photon flux by chemical actinometry with Fe2(C2O4)3

Table C.2: Determination of photon flux of UV light (Lourmat Vibert, 6 watt) by chemical actinometry with ferrioxalate. No

[Ferrioxalate] (mM)

pH

Reaction time (s)

kapp (s-1)

1

5.956

1.235

1230 1200

I0

R2

(ΦFerrioxalate = 1.25)

3.212 10-6

0.997

5.01

3.247 10

-6

0.999

5.06

3.330 10

-6

0.999

5.19

-6

0.999 0.999

Mean value

5.23 5.21 5.14

Standard deviation

0.098

2 3 4 5

6.043 6.036 6.116 6.094

1.180 1.195 1.200 1.210

1170 1200 1200

3.351 10 3.344 10-6

(E s-1)

The results showed that the photonic fluxes obtained by actinometry with potassium ferrioxalate (I 0= 5.14 10-6 E.s-1) is approximately 11.7 to 16.5 % higher than the one determined by actinometry with hydrogen peroxide (I0 = 4.52 10-6 E.s-1).


323

T. G. Le et al.

Table C.3: Mean values of photonic fluxes determined in the present work. Quantum yield of photolysis (Φ) at 254 nm and  < 400

Actinometry

I0 (10-6 E.s-1)

nm

Do at

Do at

253.7 nm

313 nm

(t = 0 s)

(t = 0 s)

H2O2 (100 m M)

1.0

4.64 ± 0.096

6.486

0.3209

H2O2 (50 mM)

1.0

4.47 ± 0.094

3.243

0.1604

H2O2 (10 mM)

1.0

4.46 ± 0.086

0.648

0.0321

H2O2 (0.05 mM)

1.0

4.41 ± 0.063

0.0032

0.0002

31.567

25.978

Mean

4.52 ± 0.126

Fe2(C2O4)3 (6 mM)

1.25

5.14 ± 0.098

This difference which has also been observed by Nicole et al. (1990) can be attributed to the fact that low pressure mercury lamps also emit radiation at wavelengths higher than 300 nm. Photonic flux measured by actinometry with Fe2(C2O4)3 takes into account a fraction of the radiation emitted at wavelengths higher than 300 nm whereas only photons emitted at wavelengths lower than 300 nm were totally absorbed by H 2O2 (Fig.C.3). 16

t o

14

DO (l = 3.45 cm)

12 10 8

N

H2 O2

o D

6 4 2

o C

y p

Fe 2 (C2 O4 )3

0

200

250

300

350

400

450

500

Wavelength (nm)

Figure C.3: Calculated spectra of solutions of H2O2 ( 50 mM, pH < 8) and of Fe2(C2O4)3 (6 mM). The radiations at wavelengths higher than 254 nm can also contribute to the photoreduction of ferric ion in reactors equipped with low pressure mercury lamps. For the experiments conducted with H 2O2, calculations with Gepasi indicate that the contributions of UV irradiation emitted at  > 254 nm to the degradation of H 2O2 are nearly equal to 4.1 % for [H2O2]0 = 100 mM, and 0.34 % for [H2O2]0 = 50 µM. These calculations have been made by taking the following values: I 0 at 253.7 nm = 4.52 10-6 E s-1, I0 at  =313 nm = 6.2 10-7 E s-1, l = 3.45 cm.

324


T. G. Le et al.

D. Photoreduction of ferric species: The Fe(III)/Tert-butanol/N2 system Table D.1: Effect of tert-butanol concentration on the apparent quantum yields of photoreduction of Fe(III). ([Fe(III)]T,0 = 3 mM , pH  1.05 - 1.08 ; D0  33 and 1.0 for l = 3.45 at 254 nm and 313 nm, respectively). FeIII Tert-B (mM) 3.000 0

pH

1.05

Final Φapp Φapp kOH* [FeII]final reaction [FeII] final kapp (Io H2O2 (Io FeOx = s-1 time (mM) (10-7 M.s-1) = -6 5.14.10 ) {B} (s) 4.52 10-6) 600

3.12 10-5 0.520 -3

1.363

0.0197

8.42 103

0.0588

0.0517

3.12 10

5

0.0222

kOH*[Tert-B] s-1 {A} 0.00 3.40 10

A/B

0 6

10.9

3.000 5

1.05

8400

1.16 10

3.001 5

1.08

7380

9.84 10-4 1.321

0.057

0.0501

2.66 105

3.40 106

12.0

3.000 5

1.05

7300

9.91 10-4 1.344

0.058

0.0510

2.67 105

3.40 106

12.7

-3

5

6

23.7

3.010

10 1.05

7380

1.06 10

1.426

0.0615

0.0541

2.87 10

3.008

10 1.09

6900

9.59 10-4 1.377

0.0594

0.0522

2.59 105

6.80 106

26.3

7440

1.03 10

-3

0.0519

2.77 10

5

6.80 10

6

24.5

-3

5

1.70 10

7

61.2 60.0

3.004

10 1.10

1.368

0.059

6.80 10

3.021 25

1.05

7200

1.03 10

1.416

0.0611

0.0537

2.78 10

3.020 25

1.08

7200

1.05 10-3 1.444

0.0623

0.0548

2.83 105

1.70 107

-4

5

7

235.5

2.921

53 0.87

3840

5.67 10

1.463

0.0631

0.0555

1.53 10

3.60 10

2.925

53 1.02

3960

5.58 10-4 1.395

0.0602

0.0529

1.51 105

3.60 107

239.4

2.955

53 0.75

7200

1.04 10-3 1.435

0.0619

0.0544

2.81 105

3.60 107

128.1

-3

5

7

126.7

3.004

53 0.77

7200

1.05 10

1.449

0.0625

0.0550

2.85 10

3.000

53 1.10

3300

4.66 10-4 1.398

0.0603

0.0530

1.26 105

3.60 107

286.7

6900

-4

0.0542

2.69 10

5

3.60 10

7

134.2

5

3.60 10

7

126.6

3.004

53 1.08

9.95 10

-3

1.428

0.0616

3.60 10

3.001

53 1.05

7440

1.05 10

1.405

0.0606

0.0533

2.85 10

3.003

53 1.10

7440

1.07 10-3 1.430

0.0617

0.0543

2.90 105

3.60 107

124.3

-3

0.062

0.0545

2.82 10

5

7

241.3

3.003 100

1.10

7200

1.04 10

3.020 100

1.05

7200

1.03 10-3 1.423

0.0614

0.0540

2.79 105

6.80 107

243.4

200 1.09

6900

1.00 10-3 1.437

0.062

0.0545

2.70 105

1.36 108

503.5

-4

5

8

524.2

1.36 108

479.2

3.000

1.437

3.008

200 1.10

6600

9.61 10

1.442

0.0622

0.0547

2.59 10

3.008

200 1.08

7260

1.05 10-3 1.435

0.0619

0.0544

2.84 105

6.80 10

1.36 10


325

T. G. Le et al.

E. Effect of pH on the rate of photoreduction of Fe(III) complexes Table E.1: Effect of pH on the apparent quantum yields of photoreduction of Fe(III). ([Fe(III)] T,0 = 1 or 3 mM, pH : 0.7 – 3.0; D0  34 and 7.0 (for l = 3.45 cm) at 254 nm and 313 nm, respectively). Do Do Φapp Φapp FeIII Tert-B % 254nm 313 nm Reaction time kapp 3+ No pH. % Fe 2+ -7 -1 (Io H2O2 = (Io FeOx = mM mM FeOH (l =3.45 (l=3.45 (s) (10 M.s ) 4.64.10-6) 5.14.10-6) cm) cm) N1 2.921 53 0.87 97.86 2.11 32.86 0.82 3840 1.437 0.0620 0.0560

326

N2

2.955 53

0.75

98.37 1.61

33.36

0.73

7200

1.409

0.0608

0.0549

N3

3.004 53

0.77

98.30 1.69

33.90

0.76

7200

1.426

0.0615

0.0555

N4

2.925 53

1.02

96.98 2.96

32.69

0.99

3960

1.372

0.0592

0.0534

N5

3.000 53

1.10

96.38 3.53

33.39

1.13

3300

1.375

0.0593

0.0535

N6

3.004 53

1.08

96.54 3.38

33.47

1.10

6900

1.405

0.0606

0.0547

N7

3.001 53

1.05

96.77 3.16

33.49

1.05

7440

1.381

0.0538

N8

3.003 53

1.10

96.38 3.53

33.42

1.13

7440

1.407

0.0607

0.0548

N9

3.000 200

1.09

96.46 3.46

33.41

1.11

6900

1.412

0.0609

0.0550

N10 3.008 200

1.10

96.38 3.53

33.47

1.13

6600

1.419

0.0612

0.0552

N11 3.008 200

1.08

96.54 3.38

33.51

1.10

7260

y p

0.0596

1.412

0.0609

0.0549

N12 3.020 100

1.08

96.54 3.38

33.65

1.10

7200

1.412

0.0609

0.0550

N13 3.003 100

1.10

96.38 3.53

33.42

1.13

7200

1.400

0.0604

0.0545

N14 3.000 53

2.08

70.89 24.82 26.87

5.42

4860

1.729

0.0746

0.0674

N15 3.000 53

2.05

72.45 23.67 27.28

5.19

4420

1.741

0.0751

0.0678

N16 3.004 53

1.99

75.38 21.45 28.09

3480

1.597

0.0689

0.0622

N17 2.769 53

2.35

54.86 35.77 20.78

7.03

3600

1.989

0.0858

0.0774

M1 0.997 53

0.68

98.62 1.37

11.28

0.23

4020

1.428

0.0616

0.0556

M2 0.988 53

0.78

98.27 1.72

11.15

0.25

3060

1.419

0.0612

0.0553

M3 0.997 53

0.75

o D

4.75

98.38 1.61

11.26

0.25

3600

1.412

0.0609

0.0550

M4 0.980 53

0.76

98.34 1.65

11.06

0.24

3600

1.421

0.0613

0.0554

M5 1.018 53

0.87

97.87 2.11

11.45

0.29

4200

1.419

0.0612

0.0553

M6 0.958 53

1.10

96.43 3.54

10.67

0.36

3720

1.423

0.0614

0.0554

M7 0.987 53

1.10

96.43 3.54

10.99

0.37

3600

1.416

0.0611

0.0552

M8 1.012 53

1.12

96.27 3.70

11.26

0.39

3600

1.421

0.0613

0.0553

M9 0.975 53

1.05

96.81 3.16

10.89

0.34

3600

1.412

0.0609

0.0550

M10 1.010 53

1.07

96.66 3.31

11.26

0.36

5520

1.423

0.0614

0.0554

M11 1.018 53

1.08

96.59 3.38

11.35

0.37

5400

1.384

0.0597

0.0539

M12 1.027 53

1.10

96.43 3.54

11.44

0.39

3600

1.412

0.0609

0.0550

M13 1.036 53

1.15

96.01 3.95

11.50

0.42

3600

1.432

0.0618

0.0558

M14 1.010 53

1.11

96.35 3.61

11.24

0.39

4200

1.432

0.0618

0.0558


t o

N

o C

T. G. Le et al. M15 0.964 53

1.58

89.77 9.94

10.21

0.78

3966

1.500

0.0647

0.0584

M16 1.016 53

1.83

82.88 16.32 10.19

1.26

4800

1.569

0.0677

0.0611

M17 1.023 53

1.86

81.84 17.27 10.17

1.33

4200

1.588

0.0685

0.0618

M18 1.020 53

1.88

81.11 17.92 10.08

1.37

4800

1.558

0.0672

0.0607

M19 0.965 53

1.81

83.55 15.71 9.73

1.15

3600

1.636

0.0706

0.0638

M20 0.965 53

1.93

79.20 19.63 9.38

1.41

3300

1.660

0.0716

0.0646

M21 0.980 53

1.95

78.39 20.35 9.46

1.48

3600

1.664

0.0718

0.0648

M22 1.010 53

1.91

79.98 18.93 9.89

1.43

3570

1.579

0.0681

0.0615

M23 0.965 53

2.21

65.77 31.06 8.29

2.15

3600

1.796

0.0775

0.0700

M24 0.982 53

2.47

50.07 43.04 7.08

2.98

3780

2.065

0.0891

0.0804

M25 1.036 53

2.75

32.79 53.70 5.81

3.88

3300

2.181

0.0941

0.0850

M26 0.953 53

2.94

22.71 57.62 4.38

3.81

3000

2.385

0.1029

0.0929

M27 0.947 53

2.68

36.93 51.48 5.68

3.40

3420

2.200

0.0949

0.0856

M28 1.010 53

2.68

36.93 51.48 6.06

3.63

3660

2.225

0.0960

0.0867

M29 1.025 53

2.70

35.73 52.15 6.04

3.73

3630

2.241

0.0967

0.0873

M30 1.016 53

2.80

29.96 55.05 5.42

3.89

3600

2.269

0.0979

0.0884

M31 1.031 53

2.76

32.21 53.99 5.72

3.88

3300

2.299

0.0992

0.0895

Table E.2: Experimental conditions, rate constants, Fe(III) distribution and initial absorbances of Fe(III) solutions at 254 and 313 nm and quantum yields of photoreduction of Fe(III) obtained in the presence of chloride ions (100 or 200 mM), [FeIII] = 1 mM, [Tert- Butanol] = 106 mM 254 nm 313 nm Φapp Φappt (l=3.45 cm) (l=3.45 cm) kapp Fe Cl (Io H2O2 (Io FeOX % % % % pH (10-7 Do of 3+ 2+ + 2+ (mM) (mM) = = Fe FeOH FeCl FeCl Do of Do of Fe Do of Fe 2 M.s-1) Free -6 -6 4.52.10 ) 5.14.10 ) Fe free complexes complexes Fe III

-

1.018 100

1.2

2.1

0.091

0.080

59.8

1.8

4.9

33.5

2.17

7.27

0.16

3.73

1.001 100 1.61

2.34

0.101

0.089

55.7

5.1

5.2

33.9

2.21

7.06

0.35

3.62

1.012 100 1.91

2.5

0.108

0.095

52.1

9.6

5.1

32.6

2.27

6.73

0.64

3.45

1.02

2.2

2.83

0.123

0.107

46.8

17

4.7

29.7

2.34

6.20

1.12

3.18

1.016 100 2.56

3.25

0.141

0.123

36.3

30.4

3.7

22.3

2.45

4.93

2.05

2.53

1.015 100

2.8

3.73

0.162

0.142

27.3

39.8

2.8

17.6

2.46

3.84

2.72

1.97

1.023 100 2.85

3.85

0.167

0.146

25.4

41.5

2.6

16.4

2.46

3.58

2.85

1.84

1.001 200 1.15

1.84

0.0794

0.070

40.3

1.2

13.1

45.2

1.00

8.46

0.10

4.43

1.015 200

1.6

2.12

0.0915

0.080

37.1

3.3

14

45.4

1.00

8.36

0.20

4.37

1.006 200 2.04

2.41

0.1040

0.091

34.1

8.5

13.6

43.1

1.04

8.10

0.35

4.24

1.010 200 2.52

2.98

0.1286

0.113

27.9

21.3

11.4

35.8

1.28

7.18

1.01

3.76

1.016 200

3.32

0.1432

0.126

21.9

32

9

28.3

1.50

6.12

1.65

3.20

100

2.8


327

T. G. Le et al.

Table E.3: Experimental conditions, rate constants, Fe(III) distribution and initial absorbances of Fe(III) solutions at 254 and 313 nm and quantum yields of photoreduction of Fe(III) obtained in the presence of sulfate ions (5 or 33 mM), [FeIII] = 1 mM; [Tert-Butanol] = 53 mM. FeIII SO42(mM) pH (mM)

Fe

3+

2+

+

FeOH FeSO4 Fe(SO4)2

-

254 nm 313 nm Φappt kapp Φappt (Io FeOX (l=3.45 cm) (l=3.45 cm) -7 (10 (Io H2O2 = = -6 Do of Do of Fe Do of Do of Fe -1 4.52.10 ) Ms ) 5.14.10-6) Fe free complexes Fe free complexes

1.007 33.33 1.15 28.20 0.998 33.33 1.64 11.00

0.00

65.90

5.30

1.070

0.046

0.041

0.421

8.091

0.021

6.694

1.10

73.20

14.70

0.706

0.030

0.027

0.178

8.269

0.016

6.980

1.005 33.33 2.05 7.90 1.012 33.33 2.56 6.50

2.00

70.70

19.30

0.604

0.026

0.023

0.014

8.392

0.006

7.249

5.30

66.40

21.30

0.653

0.028

0.025

0.058

8.280

0.096

7.237

1.012 33.33 2.90 5.80

10.20

61.30

20.60

0.816

0.035

0.031

0.085

8.191

0.162

7.182

0.998

5

0.94 73.60

1.40

24.70

0.00

1.380

0.059

0.053

2.636

6.369

0.117

5.177

1.002

5

1.13 63.80

2.00

33.70

0.00

1.350

0.058

0.051

2.590

6.428

0.117

5.225

1.007

5

1.16 62.20

2.20

35.10

0.00

1.360

0.059

0.052

2.568

6.428

0.118

5.225

0.997

5

1.55 44.20

4.10

50.30

1.40

1.240

0.053

0.047

1.147

7.444

0.161

6.087

1.005

5

2.05 30.00

9.10

58.00

2.50

1.190

0.051

0.045

0.691

7.706

0.232

6.341

1.003

5

2.48 22.80 18.50

54.00

2.80

1.230

0.053

0.047

0.588

7.558

0.439

6.258

1.012

5

2.80 17.00 29.10

44.60

2.50

1.270

0.055

0.048

0.663

7.116

0.733

5.909

t o

o C

y p

Table E.4: Experimental conditions, Fe(III) distribution and rate constants of photoreduction of Fe(III) obtained in the presence of chloride ions (0 – 200 mM) at pH  1.1.

328

o D

N

FeOH2+ %

FeCl2+ %

FeCl 2+ %

Free Fe %

Fe(III) Complexes %

kapp (10-7 M s-1)

97.6

2.4

0

0

100

0

1.45

106 106

95.0

2.4

0

2.6

97.4

2.6

1.55

85.7

2.1

0

11.7

87.8

11.7

1.79

106

75.9

1.9

1.5

20.7

77.8

22.2

1.15

106

59.9

1.8

4.8

33.5

61.7

38.3

1.96 2.10

0.2

1.08

106

49.3

1.2

8.7

40.6

50.5

49.3

2.05

0.3

1.10

106

41.2

0.2

12.8

44.9

41.4

57.7

1.84

[FeIII] (mM)

NaCl (mM)

I (M)

pH

1.001

0

0.2

1.08

1.003

5

0.2

1.05

1.006

25

0.2

1.11

1.032

50

0.2

1.10

1.018

100

0.2

1.007

150

1.001

200

[TertButanol]

Fe3+ %

106


T. G. Le et al.

Table E.5: Experimental conditions, Fe(III) distribution and rate constants of production of Fe(II) obtained at pH  1.1. SO42(mM)

FeIII (mM)

pH

% Fe3+

% FeOH2+

% FeSO4+

% Fe(SO4)2-

0

1.003

1.07

97.65

2.35

0

0

0

100

1.463

0.0631

0.5

1.006

1.11

92.50

2.80

4.60

0.00

4.60

95.30

1.419

0.0612

1.0

1.012

1.11

88.40

2.70

8.90

0.00

8.90

91.10

1.412

0.0609

2.0

1.003

1.07

81.10

2.50

16.30

0.23

16.53

83.60

1.363

0.0588

3.0

1.007

1.14

74.80

2.30

22.70

0.28

22.98

77.10

1.335

5.0

0.998

1.13

64.80

2.00

32.70

0.56

33.26

66.80

1.337

0.0576 0.0577

10

1.09

48.60

1.50

48.50

1.40

49.90

50.10

1.268

0.0547

20

1.005 1.007

1.10

32.70

0

62.70

3.70

66.40

32.70

1.189

0.0513

33.3

1.012

1.15

17.20

0

74.70

7.50

82.20

17.20

1.062

0.0458

50

1.010

1.14

12.00

0

75.40

12.10

87.50

12.00

1.041

0.0449

% % kapp Fe Total Fe (10-7 M.s-1) complexes free

Φappt (Io H2O2 = 4.52.10-6)

Table E.6: Decomposition of H2O2 by Fe(III)/H2O2/UV in the presence of hydroxyl radical scavenger and in the absence of oxygen Anion

[Anion]0 (mM)

I (M)

[H2O2]0 (mM)

[FeIII]0 (mM)

pH

ClO4ClO4ClO4ClO4ClO4ClO4ClClClClClSO42SO42SO42-

100 100 100 100 100 100 100 100 100 100 100 33 33 33

0.2 0.2 0.2 0.2 0.1 0.1 0.2 0.2 0.1 0.2 0.1 0.2 0.2 0.1

1 1 1 1 1 1 1 1 1 1 1 1 1 1

1 1 1 1 1 1 1 1 1 1 1 1 1 1

1.19 1.21 1.17 2.11 2.06 2.7 1.23 1.26 2.09 2.12 2.73 1.24 2.23 3

kapp (H2O2 Tert-B degradation) (mM) (s-1)

106 106 106 106 106 106 106 106 106 106 106 106 106 106

3.23 10-07 3.22 10-07 3.20 10-07 3.12 10-07 3.18 10-07 4.80 10-07 7.74 10-07 7.55 10-07 1.31 10-06 1.37 10-06 1.65 10-06 2.63 10-07 2.12 10-07 2.46 10-07


329

T. G. Le et al. 0.0007

(II)

[Fe ], M

0.0006

pH 2.35 pH 2.05 pH 1.02

0.0005 0.0004 0.0003 0.0002 0.0001 0 0

1000

2000 Time (s)

3000

4000

Figure E.1: Photoreduction of Fe(III) in NaClO4/HClO4 solution: effect of pH; [Fe(III)]0 = 3 mM; [tertbutanol] = 53 mM; I = 0.2 M 0.6 [FeIII] = 1 mM;

Fe II (mM)

[Cl-] = 200 mM 0.4

t o

0.2

0.0 0

III

N

500

o D

o C

y p

pH 2.80 pH 2.52 pH 2.04 pH 1.60 pH 1.15

1000 Time (s)

1500

2000

[FeIII] = f(t)

-

[Fe ] = 1 mM ; [Cl ] = 200 mM; [Tert-Butanol] = 106 mM

Figure E.2: Effect of pH on the formation of ferrous ion in the presence of chloride ion (200 mM); Experimental conditions given in table E.2 0.5

[FeIII] =1 mM

FeII (mM)

0.4

[SO42-]= 5 mM

0.3 0.2 0.1

pH = 1.13; 1.25; 2.05; 2.48; 2.80

0 0

1000

2000

3000

Time(s)

[FeIII] = f(t) ;[FeIII] = 1 mM ; [SO42-] = 5 mM; [Tert-Butanol] = 53 mM Figure E.3: Effect of pH on the formation of ferrous ion in the presence of sulfate ion (5 mM); Experimental conditions given in table E.3 330
