J Solution Chem (2007) 36: 1479–1493DOI 10.1007/s10953-007-9192-8

O R I G I NA L PA P E R

The Oxidation of Iron(II) with Oxygen in NaCl Brines

J. Michael Trapp · Frank J. Millero

Received: 12 December 2006 / Accepted: 2 March 2007 / Published online: 20 September 2007© Springer Science+Business Media, LLC 2007

Abstract The oxidation of nanomolar levels of iron(II) with oxygen has been studied inNaCl solutions as a function of temperature (0 to 50 °C), ionic strength (0.7 to 5.6 mol·kg−1),pH (6 to 8) and concentration of added NaHCO3 (0 to 10 mmol·kg−1). The results have beenfitted to the overall rate equation:

d[Fe(II)]/dt = −kapp[Fe(II)][O2]

The values of kapp have been examined in terms of the Fe(II) complexes with OH− andCO2−

3 . The overall rate constants are given by:

kapp = αFe2+kFe + αFe(OH)+kFe(OH)+ + αFe(OH)2kFe(OH)2 + αFe(CO3)2kFe(CO3)2

where αi is the molar fraction and ki is the rate constant of species i. The individual rate con-stants for the species of Fe(II) interacting with OH− and CO2−

3 have been fitted by equationsof the form:

lnkFe2+ = 21.0 + 0.4I 0.5 − 5562/T

lnkFeOH = 17.1 + 1.5I 0.5 − 2608/T

lnkFe(OH)2 = −6.3 − 0.6I 0.5 + 6211/T

lnkFe(CO3)2 = 31.4 + 5.6I 0.5 − 6698/T

These individual rate constants can be used to estimate the rates of oxidation of Fe(II) over alarge range of temperatures (0 to 50 °C) in NaCl brines (I = 0 to 6 mol·kg−1) with differentlevels of OH− and CO2−

3 .

Keywords Fe(II) · Oxidation · NaCl · Brines

J.M. Trapp · F.J. Millero (�)Rosenstiel School of Marine and Atmospheric Science, University of Miami, Miami, Florida 33149,USAe-mail: fmillero@rsmas.miami.edu

1480 J Solution Chem (2007) 36: 1479–1493

1 Introduction

Recently, a number of workers [1–19] have been interested in the rates of oxidation of Fe(II)with O2 and H2O2. Stumm and Lee [1] first measured the oxidation of ferrous to ferric ironby oxygen in dilute solutions. Their results showed a second degree pH dependence frompH = 6.5 to 7.5. They fit their results to:

d[Fe(II)]/dt = −k[OH−]2[O2][Fe(II)] (1)

where the brackets denote concentrations and k is the rate constant. Lowson [2] later foundthat in the pH range less than two, the rates were independent of pH and followed the rateequation:

d[Fe(II)]/dt = −k[O2][Fe(II)] (2)

The effect of ionic interactions on the reaction rate was examined by Sung and Mor-gan [3]. They showed that at constant ionic strength, when ClO−

4 is replaced with Cl− andSO2−

4 , that the rates decreased from the initial value. Millero [4] first suggested that the for-mation of Fe(OH)+ and Fe(OH)2 complexes caused the changes in the rates of oxidation ofFe(II) over the pH range of the measurements of Stumm and Lee [1], and Roekens and VanGrieken [5]. Millero [4] also suggested that the effect of Cl− and SO2−

4 on the rates deter-mined by Sung and Morgan [3] could be attributed to the formation of FeCl+ and FeSO4

complexes. The rate constants verses pH resembled a titration curve of Fe(II) as a functionof OH− [4]. The effect of pH on the oxidation of iron(II) was attributed to the hydrolysisspecies having different rates of oxidation [4, 6, 7]. Thus, the reaction of iron(II) with oxy-gen can be considered to be composed of several parallel reactions involving the variousiron(II) species reacting at individual second-order rates ki (where the subscripted i denotesthe species) [4]:

Fe2+ + O2 → Products kFe (3)

Fe(OH)+ + O2 → Products kFe(OH)+ (4)

Fe(OH)2 + O2 → Products kFe(OH)2 (5)

Fe(OH)−3 + O2 → Products kFe(OH)3− (6)

The overall rate constant, k, would be the sum of the products of the species individualrates, ki , and their molar fractions, αi :

kapp = αFe2+kFe + αFe(OH)+kFe(OH)+ + αFe(OH)2kFe(OH)2 + αFe(OH)3−kFe(OH)3− (7)

Millero and coworkers [4, 8–12] completed a detailed kinetic study to quantify the roleof ionic strength, media composition, temperature and pH on the oxidation of iron(II). Theeffect of bromine, nitrate, sulfate, carbonate and borate were studied by Millero and Iza-guirre [8, 9]. The results in 0.7 mol·kg−1 NaCl showed a strong decrease in k when borateand sulfate were added to the NaCl solution and small changes due to the addition of nitrateand bromine. They also showed that rates were increased with added carbonate. This was inagreement with the unpublished earlier work of Ghosh [13]. This suggests that the formationof carbonate ion pairs is also important [8, 14, 15],

Fe2+ + HCO−3 → Fe(HCO3)

+ (8)

J Solution Chem (2007) 36: 1479–1493 1481

Fe2+ + CO2−3 → Fe(CO3) (9)

Fe2+ + 2CO2−3 → Fe(CO3)

2−2 (10)

and that the reactions of these complexes with O2 should be considered in the overall reac-tion rates:

Fe(CO3) + O2 → Products kFe(CO3)+ (11)

Fe(CO3)2−2 + O2 → Products kFe(CO3)2 (12)

Fe(HCO3)+ + O2 → Products kFe(HCO3) (13)

The overall reaction rate of the oxidation, ki , should be the sum for all of the species insolution times the mole fraction, αi , of individual complexes:

k = −d[Fe(II)]/dt =∑

kiαi (14)

King [14] developed a mixed specific interaction ion-paring model for the oxidation of Fe(II)with oxygen as a function of media and pH. His model focused on the importance of thecarbonate species [FeCO3, Fe(CO3)

2−2 and Fe(CO3)OH−] and showed that at high levels

of carbonate these species can dominate the oxidation in natural waters. Santana-Casiano[16, 17] used this work as a reference point to evaluate Fe(II) oxidation with O2 in theabsence and presence of organic ligands and to construct a wider ranging kinetic modelfor Fe(II) oxidation in natural waters [17, 18]. Santana-Casiano et al. [18] evaluated thereaction rates of iron as a function of pH (6.5 to 8.2), NaHCO3 (0.1 to 9 mmol·kg−1) andtemperature (3 to 35 °C) in seawater. They used these results to determine the values of ki ofthe kinetically most active species using a numerical model. These studies show that at thelevels of carbonate in natural seawater, the hydroxyl species of iron dominate the rates of theoxidation. This study also demonstrates that the rates of oxidation of Fe(II) at nano-molarlevels were different than earlier studies at micromolar levels. These differences have beenattributed to a change in the reaction mechanism [18, 19]. At low levels, iron was not ableto compete for the intermediates formed in the Harber-Weiss [20] mechanism

Fe(II) + O2 → Fe(III) + O−2 · (15)

Fe(II) + O−2 ·+2H+ → Fe(III) + H2O2 (16)

Fe(II) + H2O2 → Fe(III) + OH ·+OH− (17)

Fe(II) + OH·→ Fe(III) + OH− (18)

This mechanism suggests a 4:1 stoichiometry of iron to oxygen as shown in early studiesmade at micromolar levels of Fe(II) [8, 10, 11]. The studies at nanomolar levels of Fe(II)[14, 15, 18, 19] show the rates differ from this stoichiometry because other minor con-stituents in seawater can react with the intermediates (O−

2 · and OH·). The superoxide, O−2 ·,

for example can be scavenged by nanomolar concentrations of copper(II) in solution [18]and also the reduction of the Fe(III) to Fe(II) may provide competition to the oxidation. Thesuperoxide, O−

2 ·, and hydroxide radical, OH·, are also effective at oxidizing other reducedcompounds in the systems, such as Cl− and HCO−

3 .In this paper, the oxidation of nanomolar levels of iron (II) has been studied as a function

of pH, HCO−3 and temperature in NaCl solutions from 0.7 to 5.6 mol·kg−1. These results

have been used to examine the rate constants for the oxidation of hydroxide and carbonate

1482 J Solution Chem (2007) 36: 1479–1493

complexes of Fe(II). These derived rate constants enable the expansion of previous oxidationmodels over a wide range of temperature and ionic strength in NaCl brines, and thus willallow one to model and estimate the oxidation of Fe(II) in most natural waters.

2 Experimental

The oxidization experiments were carried out in a 250 cm3 glass thermostatted beaker thatwas temperature controlled by a NesLab RTE-221 circulating water-bath (±0.02 °C). Thetop to the beaker was constructed with four openings for the working and reference elec-trodes, a glass fritted air bubbler and a pipette to remove sample aliquots.

All reagents used were of reagent grade. The solutions were made with Milli-Q H2O(18 M�). The stock solutions of Fe(II) at 2.5 × 10−3 mol·kg−1 were made from ferrousammonium sulfate, acidified with omni grade HCl to a pH = 2 to prevent the oxidation ofthe iron(II). Ferriozine (C20H12N4NaO6S2·H2O) stock and pH = 5.5 buffers were preparedusing the methods of Zhang et al. [21]. The molality of the NaCl solutions was checked usingdensity measurements (Anton Parr) and the equations of LoSurdo et al. [22]. The solutionswere saturated with O2 by bubbling air through them for one hour that had previously beenbubbled through MnO−

4 and Milli-Q H2O. Then the samples were adjusted to the desiredpH with HCl and maintained throughout the experiment to ±0.02 pH units.

Samples were removed from the reaction vessel by a 10 cm3 repeat pipetter, into a 25 cm3

volumetric flask containing 2 cm3 of buffer and 50 × 10−6 cm3 of Ferrozine stock thatwas brought to volume using milli-Q ion-exchange water and allowed to equilibrate for30 min before measuring. Iron(II) measurements were determined using a World PrecisionInstruments 5 m waveguide capillary flowing cell and Ocean Optics S2000 Spectropho-tometer. Measurements of the purple Ferrozine-iron(II) complex were made at 562 nm. Lin-ear calibrations of the cell and spectrophotometer were made using Fe(II) solutions from10 nmol·kg−1 to 1 µmol·kg−1 concentrations of iron (II).

The pH was measured using a glass electrode and a calomel electrode filled with thecorresponding NaCl concentration to avoid junction potentials. The electrode system wascalibrated by titration with HCl in all the NaCl solutions at each temperature and ionicstrength. The [OH−] concentration was determined from the pH using the values of Kw, forwater in NaCl using a Pitzer ionic interaction model [23, 24]. The molar fractions, αi , ofthe Fe(II) complexes at a given temperature and ionic strength were also calculated usingthe Pitzer model [23, 24]. Oxygen concentrations in NaCl as a function of temperature andionic strength were calculated from equations of Millero et al. [25].

3 Results and Discussion

The rates of iron(II) oxidation with oxygen have been expressed as an apparent oxidationrate, kapp, independent of the mechanism describing the process, given by:

d[Fe(II)]/dt = −kapp[Fe(II)][O2] (19)

When the reaction occurs in an excess of oxygen, it becomes a pseudo-first-order reactionexpressed as:

d[Fe(II)]/dt = −k1[Fe(II)] (20)

J Solution Chem (2007) 36: 1479–1493 1483

Table 1 The pseudo-first-orderrate constant for the oxidation ofnmol·kg−1 levels of Fe(II) withO2 at total carbon =2.0 mmol·kg−1 in NaCl solutions

NaCl Temp pH [O2] log10 k1

mol·kg−1 °C µmol·kg−1 min−1

0.25 25 8.00 241 −0.24

0.70 5 6.13 312 −2.89

0.70 5 6.34 312 −2.74

0.70 5 6.59 312 −2.62

0.70 5 6.86 312 −2.51

0.70 5 6.86 312 −2.51

0.70 5 7.04 312 −2.35

0.70 5 7.05 312 −2.39

0.70 5 7.29 312 −2.37

0.70 5 7.40 312 −2.22

0.70 5 7.40 312 −2.36

0.70 5 7.55 312 −1.89

0.70 5 7.62 312 −1.80

0.70 5 7.64 312 −1.77

0.70 5 7.83 312 −1.60

0.70 15 7.08 315 −2.00

0.70 25 5.95 212 −2.42

0.70 25 6.07 212 −2.49

0.70 25 6.25 212 −2.49

0.70 25 6.29 212 −2.37

0.70 25 6.50 212 −2.26

0.70 25 7.06 212 −2.02

0.70 25 7.34 212 −1.46

0.70 25 7.42 212 −1.53

0.70 25 7.42 212 −1.52

0.70 25 7.75 212 −1.05

0.70 25 7.92 212 −0.84

0.70 25 8.00 212 −0.52

0.70 25 8.29 212 −0.08

0.70 25 8.52 212 0.18

0.70 40 5.84 171 −2.03

0.70 40 6.05 171 −2.05

0.70 40 6.33 171 −2.01

0.70 40 6.61 171 −1.77

0.70 40 6.86 171 −1.46

0.70 40 6.89 171 −1.34

0.70 40 7.20 171 −0.88

0.70 40 7.41 171 −0.58

0.70 40 7.63 171 −0.19

0.70 40 7.93 171 −0.16

3.17 25 5.60 117 −2.70

3.17 25 5.97 117 −2.44

3.17 25 6.12 117 −2.61

1484 J Solution Chem (2007) 36: 1479–1493

Table 1 (Continued)NaCl Temp pH [O2] log10 k1

mol·kg−1 °C µmol·kg−1 min−1

3.17 25 6.24 117 −2.39

3.17 25 6.38 117 −2.46

3.17 25 6.69 117 −2.34

3.17 25 7.00 117 −2.08

3.17 25 7.00 117 −2.07

3.17 25 7.38 117 −1.73

3.17 25 7.78 117 −1.31

3.17 5 7.42 154 −2.15

3.17 40 7.42 101 −0.82

4.32 25 7.99 88 −0.94

4.40 25 8.00 91 −0.70

5.56 25 5.85 70 −2.66

5.56 25 6.25 70 −2.66

5.56 25 6.25 70 −2.63

5.56 25 6.48 70 −2.49

5.56 25 6.74 70 −2.42

5.56 25 7.01 70 −2.34

5.56 25 7.31 70 −2.24

5.56 25 7.40 70 −2.15

5.56 25 7.42 70 −1.86

5.56 25 7.56 70 −2.00

5.56 25 7.75 70 −1.54

5.56 25 8.07 70 −1.07

5.56 25 8.35 70 −0.56

where k1 = kapp[O2] and includes the effect of temperature and media on both kapp and thesolubility of oxygen. In order to study the effect of the concentration of carbonate, pH, ionicstrength and temperature on the pseudo-first-order rate constant, a series of experimentswere conducted and the results are given in Tables 1 and 2. The first series of measurementswere made at a constant concentration of carbonate by varying the pH (5 to 9), ionic strength(0.7 to 6 mol·kg−1) and temperature (5–40 °C). These results are given in Table 1. Thevalues of log10 k1 as a function of pH (I = 0.7 mol·kg−1) at different temperatures are shownin Fig. 1. They all show a second-degree function of pH as found in earlier studies. Thevalues of log10 k1 are a linear function of ionic strength, I 0.5, over the entire temperaturerange (see Fig. 2). The values of log10 k1 are linear functions of temperature (1/T ) over theentire concentration range (see Fig. 3). The effect of the concentration of HCO−

3 on the rateconstants are given in Table 2 and shown in Fig. 3. The values of log10 k1 are linear functionsof [HCO−

3 ]. The measurements given in Tables 1 and 2 have been used to fit the rates to theempirical functions (σ = 0.2 and 0.3, respectively)

log10 k1 = −0.33I 0.5 − 3178/T − 3.61pH + 0.32(pH)2 + 1.14 log10[HCO−3 ] (21)

log10 kapp = 0.02I 0.5 − 3732/T + 3.74pH − 0.20(pH)2 + 0.71 log10[HCO−3 ] (22)

J Solution Chem (2007) 36: 1479–1493 1485

Table 2 The pseudo-first-orderrate constant for the oxidation ofnmol·kg−1 levels of Fe(II) withO2 as a function of pH,temperature and total carbonatein NaCl solutions

NaCl Temp Total pH [O2] log10 k1

mol·kg−1 °C carbonate µmol·kg−1 min−1

mmol·kg−1

0.70 5 2.0 6.86 312 −2.51

0.70 5 3.0 6.86 312 −2.48

0.70 5 5.0 6.86 312 −2.36

0.70 5 7.5 6.86 312 −2.21

0.70 5 10.0 6.86 312 −2.15

0.70 5 2.0 7.40 312 −2.22

0.70 5 2.0 7.40 312 −2.36

0.70 5 3.0 7.42 312 −2.07

0.70 5 5.0 7.42 312 −1.73

0.70 5 7.5 7.42 312 −1.63

0.70 15 2.0 7.08 315 −2.00

0.70 15 3.0 7.08 315 −1.82

0.70 15 5.0 7.08 315 −1.54

0.70 15 7.5 7.08 315 −1.46

0.70 25 2.0 7.42 212 −1.53

0.70 25 2.0 7.42 212 −1.52

0.70 25 3.0 7.42 212 −1.13

0.70 25 5.0 7.42 212 −0.90

0.70 25 7.5 7.42 212 −0.64

0.70 25 10.0 7.42 212 −0.55

0.70 40 2.0 6.86 171 −1.46

0.70 40 3.0 6.86 171 −1.15

0.70 40 5.0 6.86 171 −1.10

0.70 40 5.0 6.86 171 −1.19

0.70 40 7.5 6.86 171 −0.71

0.70 40 7.5 6.86 171 −0.83

0.70 40 10.0 6.86 171 −0.84

0.70 40 2.0 7.41 171 −0.58

0.70 40 3.0 7.42 171 −0.29

0.70 40 5.0 7.42 171 0.05

0.70 40 7.5 7.42 171 0.33

0.70 40 10.0 7.42 171 0.35

3.17 5 2.0 7.42 154 −2.15

3.17 5 3.0 7.42 154 −1.90

3.17 5 5.0 7.42 154 −1.74

3.17 5 7.5 7.42 154 −1.67

3.17 5 10.0 7.42 154 −1.42

3.17 25 3.0 7.42 107 −1.60

3.17 25 5.0 7.42 107 −1.33

3.17 25 7.5 7.42 107 −1.13

3.17 25 10 7.42 107 −0.76

3.17 40 2.0 7.42 101 −0.82

1486 J Solution Chem (2007) 36: 1479–1493

Table 2 (Continued)NaCl Temp Total pH [O2] log10 k1

mol·kg−1 °C carbonate µmol·kg−1 min−1

mmol·kg−1

3.17 40 3.0 7.42 101 −0.64

3.17 40 7.5 7.42 101 0.18

3.17 40 10.0 7.42 101 0.46

5.56 25 2.0 7.42 70 −1.86

5.56 25 3.0 7.42 70 −1.75

5.56 25 7.5 7.42 70 −1.41

5.56 25 10.0 7.42 70 −1.34

5.56 40 3.0 7.27 64 −1.42

5.56 40 5.0 7.27 64 −1.28

5.56 40 7.5 7.27 64 −1.05

5.56 40 10.0 7.27 64 −0.88

Fig. 1 The effect of pH on thepseudo-first-order rate constantfor the oxidation of Fe(II) withO2 as a function of temperaturein 0.7 mol·kg−1 NaCl

Fig. 2 The effect of ionicstrength on the pseudo-first-orderrate constant for the oxidation ofFe(II) with O2 at 25 °C and pH =7.42 as a function the square rootof ionic strength and theconcentration of bicarbonate

J Solution Chem (2007) 36: 1479–1493 1487

Fig. 3 The effect of theconcentration of bicarbonate onthe pseudo-first-order rate for theoxidation of Fe(II) with O2 in0.7 mol·kg−1 NaCl as a functiontemperature

As shown earlier (Eq. 7), the apparent rate constant is the sum of the mole fractions of allthe iron(II) species multiplied by the individual rate constants. In order to evaluate these indi-vidual ki , a corresponding matrix of the values of αi (αi = [FeXi]/[[Fe(II)]) was calculatedfrom the formation constants of the individual complexes using the Pitzer model of ionicinteraction [24] for every reaction condition. The equations for the formation of the majorFe(II) species as a function of ionic strength and temperature are given in the Appendix.The individual rate constants for the individual complexes, ki , were first determined for thecarbonate species using the measurements given in Table 2 at a fixed pH using

kapp = A + αFe(CO3)2kFe(CO3)2 (23)

where the constant A is give by

A = αFekFe + αFe(OH)kFe(OH) + αFe(OH)2kFe(OH)2 (24)

The Fe(CO3)2−2 species is dominate, even though the αFe(CO3)2 are always under 1% of the

total iron. This is due to the large value of kFe(CO3)2. The measurements made at a fixed[HCO−

3 ] in Table 1 were then used to estimate the individual rate constants kFe, kFe(OH)

and kFe(OH)2. At a high pH, a slope of about two is observed which indicates that Fe(OH)2

controls the oxidation rate. In the lowest range of pH, the slope is almost zero and both Fe2+and FeOH+ species are important. The product of the αi and ki for Fe(OH)2 determinedat high pH can be subtracted from the kapp value allowing the kFe2+ and kFe(OH)+ to beobtained. The estimates of the log10 ki determined in this manner from all the measurementsare tabulated in Tables 3 and 4. This study found that the concentrations or rates of most ofthe minor Fe(II) species were either too low to compete with OH− and CO2−

3 . The effect oftemperature and ionic strength on the values of ki for the complexes are shown in Figs. 4,5, 6 and 7. All appear to be a linear function of 1/T and I 0.5, and have been fitted to theequations

lnkFe = 21.0 + 0.4I 0.5 − 5562/T σ = 0.3 (25)

lnkFeOH = 17.7 + 1.5I 0.5 − 2608/T σ = 0.5 (26)

lnkFe(OH)2 = −6.3 − 0.6I 0.5 + 6211/T σ = 0.6 (27)

lnkFe(CO3)2 = 31.4 + 5.6I 0.5 − 6698/T σ = 0.7 (28)

1488 J Solution Chem (2007) 36: 1479–1493

Table 3 Rate constants for theoxidation of Fe(CO3)2−

2 with O2as a function of ionic strengthand temperature

NaCl Temp pH log10 kFe(CO3)2

mol·kg−1 °C

0.70 5 6.86 5.34

0.70 5 7.40 5.08

0.70 15 7.08 5.61

0.70 25 7.42 5.72

0.70 40 6.86 6.54

0.70 40 7.42 6.80

2.08 25 7.42 6.70

3.17 5 7.42 7.53

3.17 25 7.42 7.49

3.17 40 7.42 8.75

5.56 5 7.40 10.00

5.56 25 7.42 9.71

5.56 40 7.27 10.11

Table 4 Rate constants for theoxidation of Fe2+, FeOH+ andFe(OH)2 with O2 as a functionof ionic strength and temperature

I t log10 kFe(OH)2 log10 kFeOH log10 kFe2+mol·kg−1 °C

0.70 5 8.45 4.18 0.66

0.70 40 7.42 4.48 1.66

0.70 25 7.57 3.98 1.08

3.17 25 8.94 4.83 1.23

5.56 25 10.40 5.61 1.49

Fig. 4 The rate constant for theoxidation of Fe(CO3)2−

2 with O2as a function of the square root ofionic strength and temperature

Figure 8 shows a comparison of measured (solid points) and model (lines) calculated con-centrations of Fe(II) as a function of time in 3.2 mol·kg−1 NaCl (2 mmol·kg−1 carbonate) at25 °C for studies at pH from 6.2 to 8.3. The model calculated results are in good agreementwith the experimental measurements.

J Solution Chem (2007) 36: 1479–1493 1489

Fig. 5 The rate constant for theoxidation of Fe(CO3)2−

2 with O2as a function of temperature andthe ionic strength

Fig. 6 The rate constants for theoxidation of Fe2+, FeOH+ andFe(OH)2 with O2 as a function ofthe square root of ionic strength

Fig. 7 The effect of temperatureon the rate constants for theoxidation of Fe2+, FeOH+ andFe(OH)2 with O2 in0.7 mol·kg−1 NaCl as a functionof temperature

1490 J Solution Chem (2007) 36: 1479–1493

Fig. 8 Comparison of measuredrates of oxidation of Fe(II) withthe model calculations (the lines)as a function of pH in5.6 mol·kg−1 NaCl solution at25 °C and [HCO−

3 ] =2 mmol·kg−1

Table 5 Comparison of the rate constants for the oxidation of Fe2+ , FeOH+, Fe(OH)2, Fe(CO3)2−2 with

O2 obtained in this study and literature results at 0.7 mol·kg−1 and 25 °C

Author log10 ki

Fe2+ Fe(OH)+ Fe(OH)2 Fe(CO3)2−2

This study 1.2 4.2 6.1 5.9

King [14] −4.3a 2.6a 7.0 5.1

Santana-Casiano et al. [17] −0.9 2.5 8.7 6.8

Gonzalez-Davila et al. [26] 0.5 2.4 8.0 8.7

aValues in pure water

A comparison between this study and a previous work [14, 17, 26] at I = 0.7 mol·kg−1 isshown in Table 5. Our work is in good agreement with the kinetic studies in seawater undersimilar reaction conditions. Santana-Casiano et al. [16] used a different speciation modelthan was used in this study. They consider the formation of Fe(II) chloro complexes that arenot considered in our Pitzer model [23, 24]. Our results for the rate constants of free Fe2+

and the complexes with OH− and CO2−3 adequately represent the experimental data over a

wide range of conditions in NaCl solutions. Figures 9 and 10 show the percent contributionof the most reactive species determined in this study. The Fe(CO3)

2−2 and Fe(OH)2 species

remain the dominate ones in agreement with earlier studies. The FeOH+ species becomesimportant at lower pH and can be shown to be analogous to the Fe(CO3)OH−. In acidicsolutions, the free species Fe2+ can contribute to the overall rate. The earlier low pH resultsof Stumm and Lee [1] give a value of log10 kFe2+ = −3.4 in pure water that can be comparedto our extrapolated pure water value of log10 kFe2+ = −0.4. The large differences are relatedto the large extrapolation of our measurements from pH of 6 to 4 where Fe2+ is dominant.In this study, we find a value of kFe2+ = 12 in 0.7 mol·kg−1 NaCl whereas Santana-Casianoet al. [17] found Fe2+ to be less reactive, kFe2+ = 0.12. Part of this difference is due to thespeciation models used (our model does not consider the formation of chloro complexes ofiron(II)).

J Solution Chem (2007) 36: 1479–1493 1491

Fig. 9 Percent contribution tothe overall rate of oxidation ofFe(II) as a function of carbonateconcentration at 5 °C in a3.2 mol·kg−1 NaCl solution atpH = 7.42

Fig. 10 Percent contribution ofthe overall rate of oxidation ofFe(II) as a function of pH in a3.2 mol·kg−1 NaCl solution witha 2 mmol·kg−1 carbonate bufferconcentration at 25 °C

4 Conclusions

The oxidation of nanomolar levels of Fe(II) with O2 have been determined as a function ofpH, carbonate concentration, ionic strength and temperature in NaCl solutions. The resultshave been used to determine the rate constants for the oxidation of Fe2+, Fe(OH)+, Fe(OH)2

and Fe(CO3)2−2 with oxygen over a wide range of ionic strength and temperature. These rate

constants can be used to model the rate of oxidation of Fe(II) in natural NaCl brines from 0to 50 °C and I = 0 to 6 mol·kg−1.

Acknowledgements The authors wish to acknowledge the support of the Oceanographic Section of theNational Science Foundation and the National Oceanic and Atmospheric Administration for supporting thisresearch.

1492 J Solution Chem (2007) 36: 1479–1493

Appendix

Stability constants for the formation of Fe(II) complexes [23, 24] fit by:

log10 K∗i = log10 Ki + AI + BI 0.5 + CI 2 + D/T + E lnT

Species log10 Ki A B C D E Std.

Dev.

FeCO3 26.2876 −2 −0.7 −0.031 −2029.8 −2.4676 0.05

Fe(CO3)2−2 6.7908 −2.287 0.029 −0.026 −88.68 0.1116 0.09

Fe(OH)+ −218.186 −1.1 −0.02 10120.8 33.1239 −218.186 0.03

Fe(OH)2 −267.22 −1.93 0.47 −0.03 12411.4 40.8906 0.04

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