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

Transcription

1 J Solution Chem (2007) 36: DOI /s ORIGINAL PAPER 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 in NaCl 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 NaHCO 3 (0 to 10 mmol kg 1 ). The results have been fitted to the overall rate equation: d[fe(ii)]/dt = k app [Fe(II)][O 2 ] The values of k app have been examined in terms of the Fe(II) complexes with OH and. The overall rate constants are given by: CO 2 3 k app = α Fe2+ k Fe + α Fe(OH)+ k Fe(OH)+ + α Fe(OH)2 k Fe(OH)2 + α Fe(CO3)2 k Fe(CO3)2 where α i is the molar fraction and k i is the rate constant of species i. The individual rate constants for the species of Fe(II) interacting with OH and CO 2 3 have been fitted by equations of the form: ln k Fe2+ = I /T ln k FeOH = I /T ln k Fe(OH)2 = I /T ln k Fe(CO3)2 = I /T These individual rate constants can be used to estimate the rates of oxidation of Fe(II) over a large range of temperatures (0 to 50 C) in NaCl brines (I = 0to6mol kg 1 ) with different levels of OH and CO 2 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, USA fmillero@rsmas.miami.edu

2 1480 J Solution Chem (2007) 36: Introduction Recently, a number of workers [1 19] have been interested in the rates of oxidation of Fe(II) with O 2 and H 2 O 2. Stumm and Lee [1] first measured the oxidation of ferrous to ferric iron by oxygen in dilute solutions. Their results showed a second degree ph dependence from ph = 6.5 to 7.5. They fit their results to: d[fe(ii)]/dt = k[oh ] 2 [O 2 ][Fe(II)] (1) where the brackets denote concentrations and k is the rate constant. Lowson [2] later found that in the ph range less than two, the rates were independent of ph and followed the rate equation: d[fe(ii)]/dt = k[o 2 ][Fe(II)] (2) The effect of ionic interactions on the reaction rate was examined by Sung and Morgan [3]. They showed that at constant ionic strength, when ClO 4 is replaced with Cl and SO 2 4, that the rates decreased from the initial value. Millero [4] first suggested that the formation of Fe(OH) + and Fe(OH) 2 complexes caused the changes in the rates of oxidation of Fe(II) over the ph range of the measurements of Stumm and Lee [1], and Roekens and Van Grieken [5]. Millero [4] also suggested that the effect of Cl and SO 2 4 on the rates determinedbysungandmorgan[3] could be attributed to the formation of FeCl + and FeSO 4 complexes. The rate constants verses ph resembled a titration curve of Fe(II) as a function of OH [4]. The effect of ph on the oxidation of iron(ii) was attributed to the hydrolysis species having different rates of oxidation [4, 6, 7]. Thus, the reaction of iron(ii) with oxygen can be considered to be composed of several parallel reactions involving the various iron(ii) species reacting at individual second-order rates k i (where the subscripted i denotes the species) [4]: Fe 2+ + O 2 Products k Fe (3) Fe(OH) + + O 2 Products k Fe(OH)+ (4) Fe(OH) 2 + O 2 Products k Fe(OH)2 (5) Fe(OH) 3 + O 2 Products k Fe(OH)3 (6) The overall rate constant, k, would be the sum of the products of the species individual rates, k i, and their molar fractions, α i : k app = α Fe2+ k Fe + α Fe(OH)+ k Fe(OH)+ + α Fe(OH)2 k Fe(OH)2 + α Fe(OH)3 k Fe(OH)3 (7) Millero and coworkers [4, 8 12] completed a detailed kinetic study to quantify the role of ionic strength, media composition, temperature and ph on the oxidation of iron(ii). The effect of bromine, nitrate, sulfate, carbonate and borate were studied by Millero and Izaguirre [8, 9]. The results in 0.7 mol kg 1 NaCl showed a strong decrease in k when borate and sulfate were added to the NaCl solution and small changes due to the addition of nitrate and bromine. They also showed that rates were increased with added carbonate. This was in agreement with the unpublished earlier work of Ghosh [13]. This suggests that the formation of carbonate ion pairs is also important [8, 14, 15], Fe 2+ + HCO 3 Fe(HCO 3) + (8)

3 J Solution Chem (2007) 36: Fe 2+ + CO 2 3 Fe(CO 3 ) (9) Fe CO 2 3 Fe(CO 3 ) 2 2 (10) and that the reactions of these complexes with O 2 should be considered in the overall reaction rates: Fe(CO 3 ) + O 2 Products k Fe(CO3)+ (11) Fe(CO 3 ) O 2 Products k Fe(CO3)2 (12) Fe(HCO 3 ) + + O 2 Products k Fe(HCO3) (13) The overall reaction rate of the oxidation, k i, should be the sum for all of the species in solution times the mole fraction, α i, of individual complexes: k = d[fe(ii)]/dt = k i α i (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 the carbonate species [FeCO 3,Fe(CO 3 ) 2 2 and Fe(CO 3 )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 O 2 in the absence and presence of organic ligands and to construct a wider ranging kinetic model for Fe(II) oxidation in natural waters [17, 18]. Santana-Casiano et al. [18] evaluated the reaction rates of iron as a function of ph (6.5 to 8.2), NaHCO 3 (0.1 to 9 mmol kg 1 ) and temperature (3 to 35 C) in seawater. They used these results to determine the values of k i of the kinetically most active species using a numerical model. These studies show that at the levels of carbonate in natural seawater, the hydroxyl species of iron dominate the rates of the oxidation. This study also demonstrates that the rates of oxidation of Fe(II) at nano-molar levels were different than earlier studies at micromolar levels. These differences have been attributed to a change in the reaction mechanism [18, 19]. At low levels, iron was not able to compete for the intermediates formed in the Harber-Weiss [20] mechanism Fe(II) + O 2 Fe(III) + O 2 (15) Fe(II) + O 2 +2H+ Fe(III) + H 2 O 2 (16) Fe(II) + H 2 O 2 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 studies made 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 constituents 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. The superoxide, O 2, and hydroxide radical, OH, are also effective at oxidizing other reduced compounds 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

4 1482 J Solution Chem (2007) 36: complexes of Fe(II). These derived rate constants enable the expansion of previous oxidation models over a wide range of temperature and ionic strength in NaCl brines, and thus will allow 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 cm 3 glass thermostatted beaker that was temperature controlled by a NesLab RTE-221 circulating water-bath (±0.02 C). The top to the beaker was constructed with four openings for the working and reference electrodes, 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 H 2 O (18 M ). The stock solutions of Fe(II) at mol kg 1 were made from ferrous ammonium sulfate, acidified with omni grade HCl to a ph = 2 to prevent the oxidation of the iron(ii). Ferriozine (C 20 H 12 N 4 NaO 6 S 2 H 2 O) stock and ph = 5.5 buffers were prepared using the methods of Zhang et al. [21]. The molality of the NaCl solutions was checked using density measurements (Anton Parr) and the equations of LoSurdo et al. [22]. The solutions were saturated with O 2 by bubbling air through them for one hour that had previously been bubbled through MnO 4 and Milli-Q H 2O. Then the samples were adjusted to the desired ph with HCl and maintained throughout the experiment to ±0.02 ph units. Samples were removed from the reaction vessel by a 10 cm 3 repeat pipetter, into a 25 cm 3 volumetric flask containing 2 cm 3 of buffer and cm 3 of Ferrozine stock that was brought to volume using milli-q ion-exchange water and allowed to equilibrate for 30 min before measuring. Iron(II) measurements were determined using a World Precision Instruments 5 m waveguide capillary flowing cell and Ocean Optics S2000 Spectrophotometer. Measurements of the purple Ferrozine-iron(II) complex were made at 562 nm. Linear calibrations of the cell and spectrophotometer were made using Fe(II) solutions from 10 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 the corresponding NaCl concentration to avoid junction potentials. The electrode system was calibrated by titration with HCl in all the NaCl solutions at each temperature and ionic strength. The [OH ] concentration was determined from the ph using the values of K w,for water in NaCl using a Pitzer ionic interaction model [23, 24]. The molar fractions, α i,of the Fe(II) complexes at a given temperature and ionic strength were also calculated using the Pitzer model [23, 24]. Oxygen concentrations in NaCl as a function of temperature and ionic 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 oxidation rate, k app, independent of the mechanism describing the process, given by: d[fe(ii)]/dt = k app [Fe(II)][O 2 ] (19) When the reaction occurs in an excess of oxygen, it becomes a pseudo-first-order reaction expressed as: d[fe(ii)]/dt = k 1 [Fe(II)] (20)

5 J Solution Chem (2007) 36: Table 1 The pseudo-first-order rate constant for the oxidation of nmol kg 1 levels of Fe(II) with O 2 at total carbon = 2.0 mmol kg 1 in NaCl solutions NaCl Temp ph [O 2 ] log 10 k 1 mol kg 1 C µmol kg 1 min

6 1484 J Solution Chem (2007) 36: Table 1 (Continued) NaCl Temp ph [O 2 ] log 10 k 1 mol kg 1 C µmol kg 1 min where k 1 = k app [O 2 ] and includes the effect of temperature and media on both k app and the solubility of oxygen. In order to study the effect of the concentration of carbonate, ph, ionic strength and temperature on the pseudo-first-order rate constant, a series of experiments were conducted and the results are given in Tables 1 and 2. The first series of measurements were 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. The values of log 10 k 1 as a function of ph (I = 0.7mol kg 1 ) at different temperatures are shown in Fig. 1. They all show a second-degree function of ph as found in earlier studies. The values of log 10 k 1 are a linear function of ionic strength, I 0.5, over the entire temperature range (see Fig. 2). The values of log 10 k 1 are linear functions of temperature (1/T) over the entire concentration range (see Fig. 3). The effect of the concentration of HCO 3 on the rate constants are given in Table 2 and shown in Fig. 3. The values of log 10 k 1 are linear functions of [HCO 3 ]. The measurements given in Tables 1 and 2 have been used to fit the rates to the empirical functions (σ = 0.2 and 0.3, respectively) log 10 k 1 = 0.33I /T 3.61pH (pH) log 10 [HCO 3 ] (21) log 10 k app = 0.02I /T pH 0.20(pH) log 10 [HCO 3 ] (22)

7 J Solution Chem (2007) 36: Table 2 The pseudo-first-order rate constant for the oxidation of nmol kg 1 levels of Fe(II) with O 2 as a function of ph, temperature and total carbonate in NaCl solutions NaCl Temp Total ph [O 2 ] log 10 k 1 mol kg 1 C carbonate µmol kg 1 min 1 mmol kg

8 1486 J Solution Chem (2007) 36: Table 2 (Continued) NaCl Temp Total ph [O 2 ] log 10 k 1 mol kg 1 C carbonate µmol kg 1 min 1 mmol kg Fig. 1 The effect of ph on the pseudo-first-order rate constant for the oxidation of Fe(II) with O 2 as a function of temperature in 0.7 mol kg 1 NaCl Fig. 2 The effect of ionic strength on the pseudo-first-order rate constant for the oxidation of Fe(II) with O 2 at 25 C and ph = 7.42 as a function the square root of ionic strength and the concentration of bicarbonate

9 J Solution Chem (2007) 36: Fig. 3 The effect of the concentration of bicarbonate on the pseudo-first-order rate for the oxidation of Fe(II) with O 2 in 0.7 mol kg 1 NaCl as a function temperature As shown earlier (Eq. 7), the apparent rate constant is the sum of the mole fractions of all the iron(ii) species multiplied by the individual rate constants. In order to evaluate these individual k i, a corresponding matrix of the values of α i (α i =[FeX i ]/[[Fe(II)]) was calculated from the formation constants of the individual complexes using the Pitzer model of ionic interaction [24] for every reaction condition. The equations for the formation of the major Fe(II) species as a function of ionic strength and temperature are given in the Appendix. The individual rate constants for the individual complexes, k i, were first determined for the carbonate species using the measurements given in Table 2 at a fixed ph using where the constant A is give by k app = A + α Fe(CO3)2 k Fe(CO3)2 (23) A = α Fe k Fe + α Fe(OH) k Fe(OH) + α Fe(OH)2 k Fe(OH)2 (24) The Fe(CO 3 ) 2 2 species is dominate, even though the α Fe(CO3)2 are always under 1% of the total iron. This is due to the large value of k Fe(CO3)2. The measurements made at a fixed [HCO 3 ] in Table 1 were then used to estimate the individual rate constants k Fe, k Fe(OH) and k Fe(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 Fe 2+ and FeOH + species are important. The product of the α i and k i for Fe(OH) 2 determined at high ph can be subtracted from the k app value allowing the k Fe2+ and k Fe(OH)+ to be obtained. The estimates of the log 10 k i determined in this manner from all the measurements are tabulated in Tables 3 and 4. This study found that the concentrations or rates of most of. The effect of temperature and ionic strength on the values of k i 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 the equations the minor Fe(II) species were either too low to compete with OH and CO 2 3 ln k Fe = I /T σ = 0.3 (25) ln k FeOH = I /T σ = 0.5 (26) ln k Fe(OH)2 = I /T σ = 0.6 (27) ln k Fe(CO3)2 = I /T σ = 0.7 (28)

10 1488 J Solution Chem (2007) 36: Table 3 Rate constants for the oxidation of Fe(CO 3 ) 2 2 with O 2 as a function of ionic strength and temperature NaCl Temp ph log 10 k Fe(CO3)2 mol kg 1 C Table 4 Rate constants for the oxidation of Fe 2+,FeOH + and Fe(OH) 2 with O 2 as a function of ionic strength and temperature I t log 10 k Fe(OH)2 log 10 k FeOH log 10 k Fe2+ mol kg 1 C Fig. 4 The rate constant for the oxidation of Fe(CO 3 ) 2 2 with O 2 as a function of the square root of ionic strength and temperature Figure 8 shows a comparison of measured (solid points) and model (lines) calculated concentrations of Fe(II) as a function of time in 3.2 mol kg 1 NaCl (2 mmol kg 1 carbonate) at 25 C for studies at ph from 6.2 to 8.3. The model calculated results are in good agreement with the experimental measurements.

11 J Solution Chem (2007) 36: Fig. 5 The rate constant for the oxidation of Fe(CO 3 ) 2 2 with O 2 as a function of temperature and the ionic strength Fig. 6 The rate constants for the oxidation of Fe 2+,FeOH + and Fe(OH) 2 with O 2 as a function of the square root of ionic strength Fig. 7 The effect of temperature on the rate constants for the oxidation of Fe 2+,FeOH + and Fe(OH) 2 with O 2 in 0.7 mol kg 1 NaCl as a function of temperature

12 1490 J Solution Chem (2007) 36: Fig. 8 Comparison of measured rates of oxidation of Fe(II) with the model calculations (the lines) as a function of ph in 5.6 mol kg 1 NaCl solution at 25 C and [HCO 3 ]= 2 mmol kg 1 Table 5 Comparison of the rate constants for the oxidation of Fe 2+,FeOH +, Fe(OH) 2,Fe(CO 3 ) 2 2 with O 2 obtained in this study and literature results at 0.7 mol kg 1 and 25 C Author log 10 k i Fe 2+ Fe(OH) + Fe(OH) 2 Fe(CO 3 ) 2 2 This study King [14] 4.3 a 2.6 a Santana-Casiano et al. [17] Gonzalez-Davila et al. [26] a Values in pure water A comparison between this study and a previous work [14, 17, 26]atI = 0.7mol kg 1 is shown in Table 5. Our work is in good agreement with the kinetic studies in seawater under similar reaction conditions. Santana-Casiano et al. [16] used a different speciation model than was used in this study. They consider the formation of Fe(II) chloro complexes that are not considered in our Pitzer model [23, 24]. Our results for the rate constants of free Fe 2+ and the complexes with OH and CO 2 3 adequately represent the experimental data over a wide range of conditions in NaCl solutions. Figures 9 and 10 show the percent contribution of the most reactive species determined in this study. The Fe(CO 3 ) 2 2 and Fe(OH) 2 species remain the dominate ones in agreement with earlier studies. The FeOH + species becomes important at lower ph and can be shown to be analogous to the Fe(CO 3 )OH. In acidic solutions, the free species Fe 2+ can contribute to the overall rate. The earlier low ph results of Stumm and Lee [1]giveavalueoflog 10 k Fe2+ = 3.4 in pure water that can be compared to our extrapolated pure water value of log 10 k Fe2+ = 0.4. The large differences are related to the large extrapolation of our measurements from ph of 6 to 4 where Fe 2+ is dominant. In this study, we find a value of k Fe2+ = 12 in 0.7 mol kg 1 NaCl whereas Santana-Casiano et al. [17] found Fe 2+ to be less reactive, k Fe2+ = Part of this difference is due to the speciation models used (our model does not consider the formation of chloro complexes of iron(ii)).

13 J Solution Chem (2007) 36: Fig. 9 Percent contribution to the overall rate of oxidation of Fe(II) as a function of carbonate concentration at 5 C in a 3.2 mol kg 1 NaCl solution at ph = 7.42 Fig. 10 Percent contribution of the overall rate of oxidation of Fe(II) as a function of ph in a 3.2 mol kg 1 NaCl solution with a 2 mmol kg 1 carbonate buffer concentration at 25 C 4 Conclusions The oxidation of nanomolar levels of Fe(II) with O 2 have been determined as a function of ph, carbonate concentration, ionic strength and temperature in NaCl solutions. The results have been used to determine the rate constants for the oxidation of Fe 2+, Fe(OH) +, Fe(OH) 2 and Fe(CO 3 ) 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 0 to 50 C and I = 0to6mol kg 1. Acknowledgements The authors wish to acknowledge the support of the Oceanographic Section of the National Science Foundation and the National Oceanic and Atmospheric Administration for supporting this research.

14 1492 J Solution Chem (2007) 36: Appendix Stability constants for the formation of Fe(II) complexes [23, 24] fit by: log 10 K i = log 10 K i + AI + BI CI 2 + D/T + E ln T Species log 10 K i A B C D E Std. Dev. FeCO Fe(CO 3 ) Fe(OH) Fe(OH) References 1. Stumm, W., Lee, G.F.: Oxygenation of ferrous iron. Ind. Eng. Chem. 53, (1961) 2. Lowson, R.T.: Aqueous oxidation of pyrite by molecular oxygen. Chem. Rev. 82, (1982) 3. Sung, W., Morgan, J.J.: Kinetics and product of ferrous oxygenation in aqueous solutions. Environ. Sci. Technol. 14, (1980) 4. Millero, F.J.: The effect of ionic interactions on the oxidation of metals in natural waters. Geochim. Cosmochim. Acta 49, (1985) 5. Roekens, E.J., Van Grieken, R.E.: Kinetics of iron (II) oxidation in seawater of various ph. Mar. Chem. 13, (1983) 6. Wilkins, R.G.: The Study of Kinetics and Mechanism of Reactions of Transition Metal Complexes, p Allyn and Bacon, Boston (1974) 7. Baes, C.F., Mesmer, R.E.: Iron. In: Baes, C.F., Mesmer, R.E. (eds.) The Hydrolysis of Cations, p Wiley, New York (1976) 8. Millero, F.J., Izaguirre, M.: Effect of ionic strength and ionic interactions on the oxidation of Fe(II). J. Solution Chem. 18, (1989) 9. Millero, F.J., Izaguirre, M., Sharma, V.K.: The effect of ionic interactions on the rates of oxidation in natural waters. Mar. Chem. 22, (1990) 10. Millero, F.J., Sotolongo, S.: The oxidation of Fe(II) with H 2 O 2 in seawater. Geochim. Cosmochim. Acta 53, (1989) 11. Millero, F.J., Sotolongo, S., Izaguirre, M.: The kinetics of oxidation of Fe(II) in seawater. Geochim. Cosmochim. Acta 51, (1987) 12. Millero, F.J., Stade, D.J., Sotolongo, S., Vega, C.: The effect of ionic interactions on the oxidation of Fe(II) with H 2 O 2. J. Solution Chem. 20, (1991) 13. Ghosh, M.M.: Oxygenation of ferrous iron(ii) in highly buffered water. In: Rubin, A.J. (ed.) Aqueous Environment of Metals, pp Ann Arbor Science, Ann Arbor (1974) 14. King, D.W.: Role of carbonate speciation on the oxidation rate of Fe(II) in aquatic systems. Environ. Sci. Technol. 32, (1998) 15. King, D.W., Lounsbury, H.A., Millero, F.J.: Rates and mechanism of Fe(II) oxidation at nanomolar total iron concentrations. Environ. Sci. Technol. 29, (1995) 16. Santana-Casiano, J.M., Gonzalez-Davila, M., Rodriquez, M.J., Millero, F.J.: The effects of organic compounds in the oxidation kinetics of Fe(II). Mar. Chem. 70, (2000) 17. Santana-Casiano, J.M., Gonzalez-Davila, M., Millero, F.J.: The oxidation of Fe(II) in NaCl-HCO 3 and seawater solutions in the presence of phthalate and salicylate ions: a kinetic model. Mar. Chem. 85, (2004) 18. Santana-Casiano, J.M., Gonzalez-Davila, M., Rodriquez, M.J., Millero, F.J.: Oxidation of nanomolar levels of iron(ii) with oxygen in seawater. Environ. Sci. Technol. 39, (2005)

15 J Solution Chem (2007) 36: González-Davila, M., Santana-Casiano, J.M., Millero, F.J.: Competition between O 2 and H 2 O 2 in the oxidation of Fe(II) in natural waters. J. Solution Chem. 35, (2006) 20. Harber, F., Weiss, J.: The catalytic decomposition of hydrogen peroxide by iron salts. Proc. R. Soc. London (Ser. A) 149, (1934) 21. Zhang, J.Z., Kelble, C., Millero, F.J.: Gas-segmented continuous flow analysis of iron in water with a long liquid waveguide capillary flow cell. Anal. Chim. Acta 438, (2001) 22. Lo Surdo, A., Alzola, E.M., Millero, F.J.: The PVT properties of concentrated aqueous electrolytes. I. Densities and apparent molal volume of NaCl, Na 2 SO 4, MgCl 2, and MgSO 4 solutions from 0.1 mol kg 1 to saturation and from to K. J. Chem. Thermodyn. 13, (1982) 23. Millero, F.J., Yao, W., Aicher, J.: The speciation of Fe(II) and Fe(III) in seawater. Mar. Chem. 50, (1995) 24. Millero, F.J., Pierrot, D.: A chemical equilibrium model for natural waters. Aqua Geochem. 4, (1998) 25. Millero, F.J., Huang, F., Laferriere, A.L.: Solubility of oxygen in the major sea salts as a function of temperature and concentration. Mar. Chem. 78, (2002) 26. González-Davila, M., Santana-Casiano, J.M., Millero, F.J.: Oxidation of iron(ii) nanomolar with H 2 O 2 in seawater. Geochim. Cosmochim. Acta 68(Suppl.), A377 (2004), abstracts of the 13 th Annual V.M. Goldschmidt Conference, Copenhagen, Denmark, June 5 11, 2004