Chemical Forms of Iron (Fe) in Seawater (Pourbaix Diagram)

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• Conditions under which iron(III) hydroxide is reduced to form divalent iron ions
  • Elution of divalent iron ions (Fe²⁺) from iron hydroxide (Fe(OH)₃)
    

    　In oxygenated seawater, iron ions (Fe²⁺) are quickly oxidized and precipitated as iron hydroxide (Fe(OH)₃) particles. Iron hydroxide deposited on the seafloor is reduced and leached as Fe²⁺ ions when placed under anoxic conditions in the sediment. The Fe²⁺ released from the sediment into the water directly above the sediment is oxidized relatively quickly, but if it encounters high molecular weight organic matter such as humic acid before being oxidized, it forms a complex and can remain dissolved as organic complex iron (II). The iron thus supplied to the water directly above and dissolved as organic complex iron is considered important for the marine iron cycle. In this course, we seek to determine the conditions under which iron hydroxide dissolves in sediments and is eluted as Fe²⁺.

    
    

  • 　The half-reaction equation for the leaching of Fe²⁺ from iron hydroxide and the standard Gibbs energy of formation for each component (per mol) are shown below.
    

    
    

                   　      Fe₍OH)₃   +   3 H⁺   +    e-   = Fe²⁺   +     3 H₂O   

    G_F⁰ (kJ/mol)      －696.9        0          　　      －78.9       －237.18

    
    

    The total difference in standard Gibbs energy of formation ⊿G_f⁰ between the product and the original form is,
    

    ⊿G_f⁰  = (3(－273.18)+(－78.9))－(－696.9) = －93.54×10³

     

    We calculate the standard electrode potential E₀ for this half-reaction equation.
    

    E₀ =－⊿G_F⁰ / (n・F)  = －(－93.54×10³) / 96485 / 1 = 0.969 (V)

     

    　Using Nernst equation, given a certain Fe²⁺ concentration, pH (H⁺ concentration), and temperature (T = 273+25 K), find the redox potential E for this half-reaction.　

    E = E₀　－RT/(F)・Ln {  [Fe²⁺] / [H⁺]³ }

      =  0.969 － 0.024387・Ln [Fe²⁺] + 0.024387・3・Ln [H⁺]

      =  0.969 － 0.024387・Ln [Fe²⁺] + 0.073161・(1/Log e)・(－pH)

      =  0.969 － 0.024387・Ln [Fe²⁺] － 0.168459pH

     

    　When a colloid of Fe(OH)₃ is suspended in seawater with a pH of 8.1 and a redox potential of 0.45 (V), the concentration of Fe²⁺ in dissolution equilibrium with that Fe(OH)₃ must be very low, 0.0003 pmol/L (picomole: 10^-12 mole). In other words, Fe(OH)₃ particles are almost insoluble in seawater.

    　The particles adsorb onto the organic particles as they settle and are deposited on the seafloor surface. In the fluffy layer on the sediment surface (0 to 1 cm above the sediment surface), there is a struggle between oxygen consumption by the decomposition of organic particles and oxygen supply from the water directly above, so that barely any oxygen remains. In such layers, the redox potential approaches zero. At E = 0.04 (V) and pH = 7.5, the dissolved equilibrium concentration of Fe²⁺ is 0.17 μmol/L. Although a concentration of Fe²⁺ of 0.17 μmol/L may seem low, the dissolved iron concentration (including colloidal iron and organic complex iron ) in seawater is about 10 nmol/L, so this is an order of magnitude higher. Once Fe²⁺ dissolved from Fe(OH)₃ forms complexes with humic organic matter in the sediment surface layer, dissolution of Fe(OH)₃ continues. In fact, measurements of dissolved iron concentrations in pore water in the 0 to 1 cm layer at the sediment surface have been found to reach several tens of μmol/L.
    

    (Most of the high concentrations of dissolved iron accumulated in the sediment surface layer are thought to be organic complex iron. The formation pathway is thought to be both the leached Fe²⁺ from Fe(OH)₃ particles forming complexes with humic organic matter and the leaching of organic complex iron from the organic particles into seawater as a result of decomposition).
    

• Pourbaix diagram of iron (Fe)
  • Again, we note the half-reaction equation and Nernst equation for the dissolution of iron hydroxide.

     

    Dissolution of iron hydroxide ( half-reaction equation)：　Fe(OH)₃   +   3 H⁺   +    e-   = Fe²⁺   +     3 H₂O      

    Nernst equation

    　E = E⁰　－RT/(F)・Ln {  [Fe²⁺] /（ [H⁺]³ ・[Fe(OH)₃]）

      　=  0.969 － 0.024387・Ln｛ [Fe²⁺] / [Fe(OH)₃]｝ － 0.168459pH

     

    (As a physical chemistry promise, in a solution reaction of a solid (Fe(OH)₃), the solid concentration = 1, given that there is a sufficient amount of that solid.)
    

    
    

    The condition for separating the Fe²⁺ and Fe(OH)₃ abundance ratios [Fe²⁺] / [Fe(OH)₃] = 1,
    

    ([Fe(OH)₃] = 1 (mol/L), so when [Fe²⁺] = 1 (mol/L))

     E = 0.969 - 0.168459pH.
    

     

    Taking pH on the horizontal axis and E on the vertical axis, this relational equation is depicted in the figure below.
    
    

    
    

    　The area above the black bold line in the boundary condition (Fe(OH)₃ ⇔ Fe²⁺) means that solid Fe(OH)₃ is present, and Fe²⁺ has reached dissolution equilibrium at less than 1 mol/L.

    　So what would the line be for [Fe²⁺] < 1 mol/L?
    

    　The line at the dissolution equilibrium concentration [Fe²⁺] = 1 µmol/L is added to the figure below. You can see a large shift.

    
    

• Iron and sulfur reactions in the ocean (summary)
  • 　If anoxic conditions exist near the sediment surface, the redox potential drops to near 0 (V). This causes a rapid dissolution of iron hydroxide (Fe(OH)₃) precipitated on the sediment surface, and the concentration of Fe²⁺ reaches a dissolution equilibrium of several tens of µmol/L. This is a thousand times higher than the dissolved iron concentration in the seawater immediately above. In deeper sediments, hydrogen sulfide is produced by the action of sulfate-reducing bacteria when the concentration drops to -0.1 (V).
    

    

  • 　Where hydrogen sulfide ions occur, the presence of previously leached Fe²⁺ causes a reaction between the two, resulting in the black precipitation of iron sulfide FeS. FeS (which eventually becomes FeS₂) does not dissolve easily, so it will be buried deep beneath the seafloor, creating iron deposits of pyrite (FeS₂) in a few tens of millions of years . Thus, in sedimentary layers that have progressed to sulfate reduction, the concentration of dissolved Fe²⁺ ions in the pore water will be significantly lower. HS^- ions that occur relatively deep in the sediment and escape reaction with Fe²⁺ are transported by molecular diffusion to the upper part of the sediment. HS^- ions are oxidized when they encounter oxygen, so HS^- ions are affected only up to a redox potential near zero (anoxic layer). If the water directly above the marine sediments is anoxic for some reason, HS^- ions may be affected up to near the surface of the sediments. Of course, if the bottom waters of the ocean are anoxic and the redox potential of seawater falls below -0.1 (V), sulfate reduction occurs in the seawater and HS^- ions are generated. This can happen when bottom water stagnates in a closed system inner bay.
    

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