Kinetics of Ferrous Iron Oxygenation and Redox Chemistry of Iron in Natural Waters

Author

Yu-Jean Liang, University of Rhode Island

Date of Award

1982

Degree Type

Dissertation

Degree Name

Doctor of Philosophy in Oceanography

Department

Oceanography

First Advisor

Dana R. Kester

Abstract

The kinetics and redox chemistry of iron in the marine environment was studied in three investigations: (1) the analytical techniques for the measurement of iron, (2) the kinetics of Fe(II) oxygenation, and (3) the chemistry of iron in an anoxic marine environment.

The bathophenanthroline method proved to be inadequate for the measurement of Fe(II) in the presence of Fe(III) due to Fe(III) reduction in the absence of a reducing reagent. A ferric sulfate complex method was developed to measure Fe(III) in the presence of Fe(II) at 303 run. This method obeys Beer-Lambert law for (Fe(III)) up to 10 μM with values of 2,028 ± 53 and 1,497 ± 32 (per molar-cm) for NaHCO₃ and FBW, respectively. The ferrozine method was also found to obey Beer-Lambert law for Fe(II) up to 31 μM with € value of about 29, l00 (per molar-cm).

This study provided the first systematic investigation of Fe(II) oxygenation in seawater media. Results showed that both in fresh and in marine waters the reaction rate law follows d(Fe(II))/dt = -k x P₀₂ x (OH-)² x (Fe(II)). The Fe(II) was oxidized completely during this reaction. The product, FeOOH•H₂O (≤10^-5M), did not affect the reaction rate. Based on the dependence of the reaction rate on temperature, pH, and P₀₂, the rate found at various conditions can be converted to a standard condition (STD) of 25 C, pH 8.0, and Po2 = 0.2095 atm according to k's TD = k' found x 10²(8.0 - pH) x 10^{(25.0 - T)/q} x (0.2095/P₀₂) with q = 12.52 and 15.25 for NaHCO₃ and filtered Narragansett Bay water, respectively (precision of this conversion: ± 9% for NaHCO₃ and ± 13% for seawater, one sigma). The above equation gives k' NaHCO₃ ,sTD = 32 ± 3 min-land k'_FBW,STD = 0.21 ± 0.03 min^-1. Marine dissolved organic matter was found not to retard the rate of the reaction. The major anions of seawater (Cl- and SO₄^2-), however, were found to retard the reaction and Cl- is more effective in reducing the rate than S042-. The mixing ratio experiments, using Cl- and SO₄^2-, and the ionic strength (I) effect on the reaction rate showed that the slower rate of Fe(II) oxygenation in seawater is a consequence of the effect of Cl- and SO₄^2- and the effect of higher ionic strength of seawater. Consequently, the reaction rate of Fe(II) oxygenation is controlled by the combination of the solution temperature, P₀₂, pH, solution composition, and the I of the medium.

Various parameters (T, S⁰/₀₀, pH, DO, total sulfide, (Fe(II)), total (Fe(II)), and the rate of Fe(II) oxygenation, k’) were measured in an anoxic marine environment to reveal the redox conditions and the redox chemistry of iron in that area. Results indicated that during the study period the water column in that area followed a three-layered pattern typical of fjord. The bottom layer (7-11 m) of the water contained very high (S^-2) (3.1 mM). Iron maxima were found in the middle-layer (3-6 m) at a depth of 4.5 m (2.3 and 5.9 μM). These are formed by the settling of Fe(III) from the oxic layer and the consequent reduction of Fe(III) in the mid-depth. Most of the iron was Fe(II), including the surface layer (0-2.5 m) which indicates that the flux of Fe(II) into the oxygenated waters was greater than the rate of Fe(II) oxygenation under natural conditions. The kinetic results revealed that the laboratory kinetic studies are representative of the processes that occur in the natural environment. By calculating the ionic product of (Fe(II)) and (S^2-) and comparing with other studies, it was found that in the anoxic environment the concentrations of Fe(II) and S^2- are controlled by the formation and the solubility of FeS.

Recommended Citation

Liang, Yu-Jean, "Kinetics of Ferrous Iron Oxygenation and Redox Chemistry of Iron in Natural Waters" (1982). Open Access Dissertations. Paper 1333.
https://digitalcommons.uri.edu/oa_diss/1333