ORI GIN AL PA PER

Fe(III) Reduction in the Presence of Catechol in Seawater

J. Magdalena Santana-Casiano • M. Gonzalez-Davila • A. G. Gonzalez •

F. J. Millero

Received: 15 July 2009 / Accepted: 22 December 2009 / Published online: 21 January 2010� Springer Science+Business Media B.V. 2010

Abstract Fe(II)-Fe(III) redox behavior has been studied in the presence of catechol under

different pH, ionic media, and organic compound concentrations. Catechol undergoes

oxidation in oxic conditions producing semiquinone and quinone and reduces Fe(III) in

natural solutions including seawater (SW). It is a pH-dependent process. Under darkness,

the amount of Fe(II) generated is smaller and is related to less oxidation of catechol. The

Fe(II) regeneration is higher at lower pH values both in SW with log k = 1.86 (M-1 s-1)

at pH 7.3 and 0.26 (M-1 s-1) at pH 8.0, and in NaCl solutions with log k of 1.54 (M-1 s-1)

at pH 7.3 and 0.57 (M-1 s-1) at pH 8.0. At higher pH values, rate constants are higher in

NaCl solutions than in SW. This is due to the complexation of Mg(II) present in the media

with the semiquinone that inhibits the formation of a second Fe(II) through the reaction of

this intermediate with other center Fe(Cat)?.

Keywords Iron � Catechol � Oxidation � Reduction � Kinetics � Seawater

J. M. Santana-Casiano (&) � M. Gonzalez-Davila � A. G. GonzalezDepartamento de Quımica, Facultad de Ciencias del Mar, Universidad de Las Palmas de Gran Canaria,35017 Las Palmas de Gran Canaria, Spaine-mail: jmsantana@dqui.ulpgc.es

M. Gonzalez-Davilae-mail: mgonzalez@dqui.ulpgc.es

A. G. Gonzaleze-mail: aridaneglez@gmail.com

F. J. MilleroRosenstiel School of Marine and Atmospheric Science, University of Miami, 4600 RickenbackerCauseway, Miami, FL 33149, USAe-mail: fmillero@rsmas.miami.edu

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Aquat Geochem (2010) 16:467–482DOI 10.1007/s10498-009-9088-x

1 Introduction

Iron bioavailability has been shown to affect the primary productivity in several oceanic

waters (Boye et al. 2005; de Baar et al. 2005), but also in coastal areas (Bruland et al. 2001;

Hutchins and Bruland 1998). The extent to which iron limits primary production depends on

both the abundance of iron and its speciation. The solubility of iron in open ocean waters is

extremely low (Liu and Millero 2002; Waite 2001), and particulate and colloidal iron is

believed to be unavailable to phytoplankton (Rich and Morel 1990; Wells et al. 1983).

However, colloidal Fe can be converted into more reactive species by both photooxidation

on the Fe redox cycle (Borer et al. 2005, 2009 Miller and Kester 1994; Rijkenberg et al.

2004, 2006, 2008; Wells and Mayer 1991), and organic complexation of Fe(III) (Johnson

et al. 1997; Kuma et al. 1996) resulting in a higher bioavailability to phytoplankton.

The dissolved iron fraction has been proposed as the species used by the eukaryotic

phytoplankton (Anderson and Morel 1982) with Fe(II) assumed as the fraction more

suitable for the biological uptake (Anderson and Morel 1982; Maldonado and Price 2001;

Takeda and Kamatani 1989). Although Fe(II) is the species more soluble in seawater, it is

rapidly oxidized by O2 and H2O2 (Gonzalez-Davila et al. 2005; Millero et al. 1987; King

et al. 1995; Santana-Casiano et al. 2005), resulting in a low concentration of Fe(II) in the

ocean. As a result, some organisms like bacteria have developed the ability to produce

proteins for the binding of iron or small molecules called siderophores that are specific for

sequestering iron. Siderophores are ligands with a high affinity and specificity for iron, and

under iron-limiting conditions are excreted by cyano- and heterotrophic bacteria (Reid and

Butler 1991; Wilhelm and Trick 1994; Winkelmann 1991). In regard to phytoplankton, the

acquisition of iron from iron-siderophores complexes is controversial as discussed in the

literature. Eukaryotic phytoplankton generally does not produce siderophores. However,

under iron-limiting conditions, some species are able to use iron bound to strong organic

chelators via a cell surface reductase mechanism (Borer et al. 2005; Hutchins et al. 1999;

Jones et al. 1987; Maldonado and Price 2001). Other studies have indicated that strong

iron-siderophore complexes are not available to eukaryotic phytoplankton (Wells 1999;

Wells et al. 1994). In spite of the importance of organic compounds in the Fe assimilation

by phytoplankton, there is not a clear picture about the organic compounds effect on the

thermodynamic and kinetic behavior of iron in seawater and on the Fe cycle in the ocean.

The presence of organic compounds can influence the redox speciation of Fe in seawater

(Craig et al. 2009; Santana-Casiano et al. 2000, 2005; Theis and Singer 1974) and can also

induce photoreductive dissolution of Fe from colloidal material (Sulzberger and Laubscher

1995; Waite and Morel 1984). It is known that besides increasing the solubility of iron,

siderophores also accelerate iron oxide dissolution (Hersman et al. 1995; Kraemer 2004).

The intermediate product of many biological reactions is the superoxide anion, O2•-.

The O2/O2•- redox pair with a potential of Ew

0 = -0.160 V will reduce aqueous Fe(III) to

Fe(II) that has a potential of Ew0 = ? 0.77 V. The Fe(II) can then react with H2O2 to

produce hydroxyl radicals, OH•, in aerobic systems (Haber and Weiss 1934). These

reactive oxygen species can damage cells through reactions with lipids and other bio-

molecules (Harrington and Crumbliss 2009). The chelation of Fe(III) by siderophores

presents much lower redox potential, effectively controlling the redox behavior due to the

strong selectivity of the siderophore donor groups for Fe(III) over Fe(II). The result of this

redox controlling process is preventing damage to the organisms and the production of

reactive oxygen species. The redox potential for Fe(III)/Fe(II) couple can change from

-0.08 V for rhizoferrin to -0.99 V for the complex of enterobactin with Fe(III)

(Harrington and Crumbliss 2009). Moreover, the reduction of Fe(III) can facilitate the

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solubilization of the metal since Fe(II) has higher solubility, lower Bronster acidity than

Fe(III) and more rapid ligand exchange rates (Dhungana et al. 2007). Redox potential of

iron-siderophores complexes changes with pH. Siderophores with acidic moieties as

binding groups compete with protons and, as pH is lowered, protons became a more

effective competitor with Fe(III). The result is a gradual shift of the redox potential to more

positive values with variations of solution pH.

The catechol, ortho-dihydroxybenzene, is an organic functional compound typically

present in most of the siderophores. The catechol is a quinone. The term ‘‘quinone’’ refers

to organic structures in three oxidation states linked by one-electron redox reaction. It can

form stable complexes with various di- and tri-valent metal ions, the complexes with

trivalent ions being the most stable. Catechol complexes forming part of the siderophores

have the most negative redox potential, followed by hydroxamic acids complexes and a-

hydroxycarboxylic acid complexes (Harrington and Crumbliss 2009). Catechol can also

undergo redox reactions, cycling between catechol, semiquinone radicals, and o-benzo-

quinone. The catechol with three protonation levels denoted as CatH2, CatH-, and

Cat2-(pKa1 = 9.45, pKa2 = 13.74), is the fully reduced form. The semiquinone radical

(with two protonation levels, SqH•, Sq•-, pKa1 = 5.0) is an intermediate oxidation state,

and the benzoquinone, denoted by Bq, is the fully oxidized form (Uchimiya and Stone

2006). Oxidation of catechol involves reduction of the metal center of Fe(III)-catechol

species. At pH of seawater, catechol occurs in the non-dissociated form and the catechol

radical most probably in the dissociated form.

CatH2 , Sq�� þ 2Hþ þ e� ð1Þ

Sq�� , Bqþ e� ð2ÞEquations (1) and (2) can also be expressed using the schematic structure of the

compounds.

The complexation of catechol with iron has been previously studied (Avdeef et al. 1978;

Hider et al. 1981). Depending on the pH, a number of mono-, di-, or tri-coordinated

complexes can be formed (Schweigert et al. 2001). Fe(II) complexes form preferentially at

acidic pH, whereas Fe(III) complexes occurs at pH [ 7. The type of iron-catechol complex

present depends on the iron-catechol ratios. Mono-coordinate complexes form preferen-

tially at higher ratios. Catechol is also oxidized enzymatically or in the presence of oxygen

and heavy metals. In these cases, one electron is transferred to molecular oxygen and a

superoxide is formed. In the presence of heavy metals, the superoxide is further reduced to

hydrogen peroxide and hydroxyl radicals. Catechol complexes with iron also prevent the

metal from undergoing redox reactions, as in the case of di-coordinated iron(III) com-

plexes. In contrast, mono-coordinate complexes allow iron to be involved in electron

transfer reactions (Avdeef et al. 1978). The complexes formed by Fe(II) and catecholates

compounds have strong reductant properties linked to the low standard one-electron

reduction potential of the associated Fe(III)/Fe(II) redox couple (Kim et al. 2009).

In this work, we study the effects produced by the catechol in the Fe(II)-Fe(III) behavior

under different pH, ionic media and catechol concentrations.

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2 Experimental

2.1 Chemicals

Fe(II) stock solutions were prepared using ferrous ammonium sulfate hexahydrate (Sigma),

acidified with Suprapur HCl to pH = 2.0. In most of the studies, the initial concentration of

Fe(II) was kept at 200 nmol l-1 in the reaction vessel. Pure catechol (Sigma) was added

directly to the sample to get final concentrations in the reaction vessel from 0.2 to

1 lmol l-1.

All the chemicals used for the Fe(II) determination and the preparation of the different

solutions were trace analytical grade. The seawater used in this study was collected in the

ESTOC station (European station for time series in the ocean at the Canary Islands), and it

was filtered previously by 0.45 lm and then by 0.1 lm working under axenic conditions.

2.2 Redox Experiments

The reactions were studied in a 200-ml glass thermostated vessel. The pH of the solution

was determined on the free scale, pHF, following Millero (1986). The pH was adjusted to

the desired value with additions of small amounts of 1 M HCl. The pH for the study was

recorded during the reaction. It was automatically controlled by a 719 titrinoTM (Methrom)

to keep it constant. The pH value was monitored for any change after the addition of the

Fe(II) and along the reaction time. The changes in pH were always less than 0.02, with the

highest effects occurring at low pH where the buffer capacity of the carbonate is lowest.

The samples were aerated with pure air prior and during the experiment.

2.3 Fe(II) and Catechol Analysis

A 5-m-long waveguide capillary flow cell (LWCFC) from World Precision Instruments

connected to the UV detector S4000 (Ocean OpticsTM) was used to make the measure-

ments of both Fe(II) and the organic compound.

2.3.1 Fe(II) Regeneration Studies

A 200 nM concentration of Fe(II) was added to 100 ml of solution (0.7 M NaCl or sea-

water) in a 200-ml thermostated vessel controlled to 0.02�C. The initial solution was aerated

with air for 2 h under saturation conditions in order to fully oxidize the Fe(II) solution to

form Fe(III). After that, 10-3 mol l-1 ferrozine (FZ) was added (no Fe(II) was determined

in the solution under such conditions). Then, a blank was done in the solution containing the

generated Fe(III) and the FZ. After this, the catechol was added to the solution, and this was

considered the time zero for the reaction. Some experiments were also done by adding

directly Fe(III) instead of Fe(II), and no differences were observed. When both Fe(III) and

catechol are present in the solution, the Fe(II) was regenerated and the Fe(II)-ferrrozine

complex was continuously followed by UV–VIS spectroscopy at 562 nm.

2.3.2 Catechol

Oxidation of catechol was studied under the experimental conditions of this work by

measuring the evolution of its spectra and recording the absorbance at 407 and 489 nm. A

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calibration curve was done for the oxidized catechol. Catechol solution of 1 lM was

oxidized during 1 month under aerated conditions, and a calibration curve was done using

the formed quinone in the concentration range of 0.2–1 lM. The absorbance was measured

at k = 407 nm, resulting a molar absorptivity of e407 = 1,182 ± 60 M-1 s-1.

2.3.3 Fe(II)-FZ Method Validation

In order to validate the FZ method for the Fe(II) analysis, the catechol spectrum was

compared with both the catechol spectrum in the presence of FZ and that of catechol in the

presence of Fe(II) and FZ, observing that the Fe(II)-FZ complex can be determined without

any interference in the presence of the quinone (Fig. 1).

3 Results

3.1 Catechol Oxidation in NaCl and Seawater

The oxidation of catechol with oxygen was studied at pH = 8.0 in 0.7 M NaCl solutions

and seawater. In the presence of oxygen, the catechol was oxidized to benzoquinone

through the semiquinone radical. Figure 2 shows the spectra evolution at different times

for both solutions (2, 7, 20, 60, and 113 min for seawater and 100 min for NaCl solu-

tion). After 1 h, a different spectra evolution for each solution was observed. In seawater,

a peak appears initially at 407 nm due to the presence of benzoquinone, but with time, a

second peak appears at 489 nm corresponding with the semiquinone radical. In NaCl,

this second peak was not clearly developed and the absorbance after 100 min was lower

than for seawater after 20 min. The differences observed in the catechol spectra between

both solutions indicated interactions between ions presented in the seawater and not in

the NaCl solution with the catechol. They were investigated and presented in the fol-

lowing sections.

Fig. 1 Visible spectra of catechol, catechol in the presence of ferrozine and with added 200 nM Fe(II)

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3.2 Oxidation of Catechol in Seawater. pH Effect

The oxidation of catechol was studied in seawater at different pH between 7.3 and 8.0

(Fig. 3). It was observed that as pH increases, the absorbance is higher for all the spectra.

Also the semiquinone peak at 489 nm becomes the most important aspect as it is shown

later.

3.3 Fe(II) Regeneration in NaCl–NaHCO3. pH Effect

In order to determine whether Fe(II) can be regenerated from Fe(III) in the presence of

catechol, the formation of Fe(II) in 0.7 m NaCl–2 mM NaHCO3 was followed adding FZ

Fig. 2 Catechol visible spectra (1 lM) in 0.7 M NaCl–2 mM NaHCO3 after 100 min oxidation and inseawater at different oxidation times. The structure of the semiquinone radical and quinone is included at thecorresponding wavelengths of 407 and 489 nm

Fig. 3 Catechol visible spectra (1 lM) in seawater after 2 h oxidation time at different pH values

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to the Fe(III) catechol solution at different pH from 7.0 to 7.9 (Fig. 4). Fe(II) was effi-

ciently formed being a pH-dependent process.

The reduction of Fe(III) followed a first-order kinetic with rate constants decreasing as

pH increases from log k0 = -4.13 at pH = 7.0 to log k0 = -5.42 at pH 7.9 as shown in

Table 1. The dependence of log k0 for the Fe(III) reduction as a function of pH is described

by a second-order polynomial equation

log k0 ¼ �43:37þ 12:88 pH� 0:96 pH2 �0:08ð Þ ð3Þ

3.4 Fe(II) Regeneration in NaCl–NaHCO3. Catechol Concentration Effect

The effect of catechol concentration on the Fe(III) reduction was studied in the range of

0.2–1 lM. The log k0 changes from -4.45 at 1 lM to -4.95 at 0.2 lM (Table 2; Fig. 5).

When catechol is in excess at concentrations higher than 0.6 lM, a first-order oxidation

rate dependence is obtained. Using initial Fe(III) reduction velocities, an order one with

respect to the catechol concentration for the regeneration of Fe(II) was found.The rate

equation for the reduction of Fe(III) with CaH2 is given by

d Fe IIIð Þ½ �=dt ¼ � k Fe IIIð Þ½ � CatH2½ � ð4Þ

The rate constant was obtained measuring the appearance of Fe(II). The disappearance

of [Fe(III)] = [Fe(III)]0 - [Fe(II)] was determined as a function of time under pseudo

first-order conditions (CatH2 in excess at a constant pH), the subscript zero denotes the

Fig. 4 Fe(II) regeneration as a function of time from Fe(III)-catechol complex in 0.7 M NaCl–2 mMNaHCO3 at different pH values

Table 1 Values of log k0 (in s-1)for the Fe(III) reduction at dif-ferent pH

[Fe(II)]o = 200 nM,[catechol]o = 1 lM

pH k0 (s-1) log k0

7.0 7.353 9 10-5 ± 1.56 9 10-7 -4.134

7.3 3.544 9 10-5 ± 3.73 9 10-8 -4.451

7.6 1.721 9 10-5 ± 6.28 9 10-9 -4.764

7.9 3.746 9 10-6 ± 3.47 9 10-9 -5.426

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initial concentration of Fe(III). The first-order rate constant for the disappearance of Fe(III)

under theses conditions is given by

d Fe IIIð Þ½ �=dt ¼ � k0 Fe IIIð Þ½ � ð5Þ

where k0 = k [CatH2].

3.5 Fe(II) Regeneration in Seawater

In the next step, the formation of Fe(II) was studied in seawater with catechol at different

pH. The studies were done using a pH range from 7.3 to 8.2. From Fig. 6, where the

spectra are shown, only at a pH lower than 7.8 is the complex Fe(II)-FZ obtained. As the

pH increased, the peak corresponding to the semiquinone became clearly developed, and at

pH higher than 8, no signal for the Fe(II)-FZ was presented. In Fig. 6, the spectra corre-

spond with that obtained after more than 1 h in the reaction vessel.

3.6 Catechol Spectra at Different Ionic Media

To account for the behavior observed in the seawater media, the Fe(II) regeneration was

studied in NaCl solution with the different major ions at the seawater concentration

usually found at pH = 8.0. First, the spectrum evolution of catechol without Fe(III) in the

reaction vessel was studied. The results are shown in Fig. 7. There was no difference with

respect to what was observed in NaCl solution with the presence of K?, F-, and SO42-.

Table 2 Values of log k0 (in s-1)for the Fe(III) reduction at dif-ferent catechol concentration

[Fe(II)]o = 200 nM, pH = 7.3

Catechol (lM) k0 (s-1) log k0

0.2 5.995 9 10-6 ± 2.715 9 10-8 -5.222

0.6 1.545 9 10-5 ± 3.516 9 10-8 -4.811

0.8 2.831 9 10-5 ± 3.231 9 10-10 -4.548

1 3.504 9 10-5 ± 3.729 9 10-8 -4.450

Fig. 5 Fe(II) regeneration as a function of time expressed as disappearance of Fe(III) in 0.7 M NaCl–2 mMNaHCO3 at different catechol concentrations

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However, Ca2? ions produced an increase in the signal of the spectrum, while Mg2?

strongly increased the semiquinone signal. This ion accounts for the results obtained in the

seawater solution.

3.7 Fe(II) Regeneration at Different Ionic Media

The same type of studies as mentioned earlier were done in the presence of Fe(III) at

pH = 8.0. The NaCl solutions with K?, F-, SO42-, and Ca2? ions present all the Fe(II)-FZ

signal while in the NaCl–Mg2? solution, a peak is observed at around 500 nm similar to

that found at the same wavelength in seawater (Fig. 8). The log k0 value obtained in each

Fig. 6 Visible spectra of Fe(II)-ferrozine complex in seawater in the presence of catechol (1 lM) oxidizedafter 2 h at different pH in the range 7.3 and 8.2

Fig. 7 Catechol visible spectra (1 lM) in NaCl solution after 2 h oxidation time with the different majorions at the seawater concentration (ionic strength of 0.7 M) and pH 8.0. The spectra in pure seawater is alsoincluded

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case is presented in Table 3. The interaction of the semiquinone with Mg2? ions that block

the semiquinone to quinone electron transfer and consequently the reduction of Fe(III) can

account for the observed behavior. This process should be pH dependent as we found that

Fe(II) was successfully generated at lower pH. In order to elucidate this question, both the

oxidation of catechol (Fig. 9) and the Fe(II) regeneration (Fig. 10) in the presence of

catechol were studied in a solution of NaCl–Mg2? at pH = 7.3 and compared with the

results obtained at pH = 8.0. Figure 9 showed that at pH 7.3, in a NaCl–Mg2? media, with

only catechol in the solution, the semiquinone peak was not observed in contrast to what

happen at pH = 8.0. When Fe(III) was present, the Fe(II)-FZ was detected at pH 7.3 and

Fe(II) was quantitatively formed (Fig. 10), indicating that the Mg-semiquinone complex is

only formed in alkaline solutions at pH over 7.8 and that the complex formation with Mg2?

blocks the Fe(III) reduction. For the NaCl–Mg2? solution, a more complex mechanism at

higher pH probably implying the semiquinone complex must be taking place. More studies

are planned in order to account for this process.

3.8 Fe(II) Regeneration in Seawater. pH Effect

The Fe(II) regeneration in seawater in the presence of catechol was studied at different

pH (7.3–8.0) (Fig. 11). For the reduction of Fe(III) in seawater and after considering the

Fig. 8 Visible spectra of Fe(II)-ferrozine complex in NaCl solution and 1 lM catechol after 2 h oxidationtime with the different major ions at the seawater concentration and pH 8.0. The spectra in pure seawater isalso included

Table 3 Values of log k0 (in s-1)for the Fe(III) reduction in NaCl,pH = 8.0, with different majorions. [Fe(II)]o = 200 nM

Ionic media k0(s-1) log k0

NaCl 6.0 9 10-6 -5.220

NaCl–HCO3- 3.746 9 10-6 -5.426

NaCl–F- 2.880 9 10-6 -5.541

Na-K? 2.449 9 10-6 -5.611

NaCl–Ca2? 2.106 9 10-6 -5.676

NaCl–SO42- 1.740 9 10-6 -5.833

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effect of the semiquinone peak formation in the Fe(II)-FZ signal, the Fe(III) reduction

rate was computed, following a first-order reaction dependence (Table 4). It is observed

that the Fe(II) formation decreases as pH increases, being affected at higher pH by the

blocking effect of the Mg-semiquinone complex.

3.9 Experiments in the Dark

At pH 7.5 in the dark, the oxidation of catechol forms the quinone. The spectra do not

present differences with respect to that observed under light, with no signal due to the

semiquinone intermediate (data not shown). In the presence of Fe(III) and in the dark, the

Fig. 9 Catechol visible spectra (1 lM) in 0.55 M NaCl–0.05 M MgCl2 solution after 2 h oxidation time atpH 7.5 and 8.0

Fig. 10 Visible spectra of Fe(II)-ferrozine complex in 0.55 M NaCl–0.05 M MgCl2 solution after 2 hoxidation time at pH 7.5 and 8.0

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reduction of the Fe(III)-Cat complex is lower than in the presence of light. The reduction of

Fe(III) followed a first-order kinetics under both conditions with rate constants of

5.13 9 10-5 s-1 (with light) and 4.42 9 10-6 s-1 (dark). After 2 h, 29 nM of Fe(II) were

formed under the light condition at pH 7.5 compared with only 6 nM of Fe(II) under

darkness. At pH 8.0, the blocking effect of Mg(II) ion in the oxidation of catechol in

seawater is also observed at studies in the dark. In the presence of Fe(III), both studies

presented the formation of semiquinone-Mg complexes decreasing the amount of Fe(II)

formation, without a clear difference between the two conditions.

4 Discussion

The catechol is spontaneously oxidized to benzoquinone when the oxygen is present in the

solution. This process is more rapid in a seawater solution than in a NaCl solution as shown

in Fig. 2. In the presence of oxygen, the oxidation of catechol is explained through a

spontaneous two electron transfer between the catechol and the o-benzoquinone through

the semiquinone radical (Eqs. 1, 2).

CatH2 þ Bq! 2Sq�� þ 2Hþ ð6ÞA third reaction (Eq. 6), known as comproportionation, can also be possible. This

reaction takes place between the catechol and the benzoquinone and produces semiquinone

Fig. 11 Fe(II) regeneration as a function of time expressed as disappearance of Fe(III) in seawater atdifferent catechol concentrations

Table 4 Values of log k0 (in s-1)and log k (in M-1 s-1) for theFe(III) reduction in seawater

[Fe(II)]o = 200 nM

pH log k0 log k

7.3 -4.138 1.86

7.5 -4.290 1.71

7.8 -4.697 1.30

8.0 -5.742 0.26

8.2 -8.154 -2.15

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radicals. The comproportionation is largely independent of pH below the pKSq� of the

semiquinone radical (pKSq� ¼ 5) but grows in importance with increasing pH once the pH

is greater than pKSq� (Uchimiya and Stone 2006).

The oxidation of catechol is pH dependent, being faster at higher pH values (Fig. 3).

The studies done with catechol and the oxidized Fe(II) show that the catechol can reduce

quantitatively Fe(III) in natural waters in the pH range 7.0–8.2. This is explained through

the equations

CatH2 þ 2Fe(OH)3 þ 4Hþ ¼ Bqþ 2Fe(II)þ 6H2O ð7Þ

CatH2 ¼ Bqþ 2Hþ þ 2e� ð8Þ

Fe(OH)3 þ 3Hþ þ e� ¼ Fe(II)þ 3H2O ð9ÞThe regeneration process of Fe(II) from Fe(III) with catechol in NaCl media and sea-

water is found to be different. In NaCl media, the following reactions can explain the

observed behavior and stoichiometry

Fe(III)þ CatH2 $ Fe(Cat)þ þ 2Hþ pKc ¼ 20 ð10Þ

Fe(Cat)þ ! Fe(II)þ Sq�� k ðslowÞ ð11Þ

Fe(Cat)þ þ Sq�� ! Fe(II)þ Bq k2ðfastÞ ð12ÞFe(III) reacts with catechol to form a complex that produces Fe(II) and the semiquinone

radical. This radical reacts with other center Fe(Cat)? and results in a second Fe(II)

molecule and the benzoquinone.

However, cations are known to participate in bidentate coordination to o-phenol by

replacing the hydrogen atoms of hydroxyl groups, O–Me2?–O. By contrast, anion coor-

dination with both monatomic and polyatomic anions results in a complex between the

anion and the hydroxyl groups, OH–Am–OH. Both types of ions recognition events alter

the properties of catechol (Brooksby et al. 2008).

In seawater or NaCl solution with NaCl–Mg2? ions, reaction (12) is blocked by the

presence of Mg2? ions that form a strong complex with the semiquinone radical (Eq. 13).

Sq�� þMg(II)$ SqMg alkaline conditions ð13ÞThis process is pH dependent. The comparison of the spectra recorded during the

catechol oxidation in different ionic media including seawater (Fig. 7) reveals that the

presence of Mg(II) ions greatly influences the catechol oxidation in alkaline solutions

(Fig. 9) due to the formation of a complex between the Mg2? ion and the semiquinone

radical. Moreover, the Mg(II)-semiquinone complex affects the Fe(II) regeneration in

seawater, being a pH-dependent process. These results are in agreement with that found by

Nikolic et al. (1998) that showed Mg2? ion may form labile complexes with catechol in

weakly alkaline aqueous solutions. Taking into account the effect of magnesium, the rate

constant for the reduction of Fe(III) in seawater in the presence of 1 lM catechol was

calculated. It can be described by a rational function of pH valid from pH 7.3–8.2

(Fig. 11).

The studies show that Fe(II) regeneration is higher at low pH values both in SW, with

log k = 1.86 (M-1 s-1) at pH 7.3 and 0.26 (M-1 s-1) at pH 8.0, and in NaCl solutions,

with log k of 1.54 (M-1 s-1) at pH 7.3 and 0.57 (M-1 s-1) at pH 8.0. At higher pH values,

rate constants are higher in NaCl solutions than in SW due to the complexation of Mg(II)

with the semiquinone that inhibits the second reduction process. When the studies are

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carried out under darkness and at pH 8.0, this complexation process dominates over the

reduction of Fe(III), being the amount of Fe(II) formed similar to that under light condi-

tions. However, at pH 7.5, where the formation of the Mg-complex is not appreciable, the

reduction of catechol under light conditions favors the formation of Fe(II), and the overall

equilibrium is shifted toward the formation of iron(II). The decrease in pH potentially

increases the complex redox potential, resulting in reduction of the Fe(III) center and

higher Fe(II) formation by binding sites of catechol.

5 Conclusions

Catechol can acts as an important organic compound by reducing Fe(III) in natural solu-

tions including seawater. This is a pH-dependent process related to shifts in the redox

potential of Fe(III)/Fe(II) and effective complex stability. Under darkness, the amount of

Fe(II) generated is smaller related to less oxidation of the catechol. The comparison of the

spectra recorded in different ionic media including seawater revealed that the presence of

Mg(II) ions greatly influences the catechol oxidation in alkaline solutions due to the

formation of a complex between the Mg(II) ion and the semiquinone radical. The Mg(II)-

semiquinone complex affects the Fe(II) regeneration in seawater. The Fe(II) regeneration is

higher at lower pH values both in SW and in NaCl solutions. At a pH higher than 7.8, the

Mg(II) complexation decrease the oxidation rate in seawater and in NaCl–Mg solutions.

Acknowledgments This study was supported by the Project CTM2006-09857 of Ministerio de Ciencia yTecnologıa from Spain. F.J. Millero wishes to acknowledge the support of the Oceanographic Section of theNational Science Foundation and the National Oceanic and Atmospheric Administration for supporting hismarine physical chemistry studies.

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