II. Reduction – gaining electrons (valence becomes more negative)
    A reductant is a compound that donates electrons

III. Oxidation – losing electrons (valence becomes more positive)
    An oxidant is a compound that accepts electrons
    “LEO the lion says GER”

III. Redox potential
    A. Arbitrary scale – how oxidized or reduced things are compared to a standard

         Standard reaction : 

    B. Eh – electron flow relative to the hydrogen standard (electron potential)
        As Eh decreases, the solution is more reduced (has more electrons to give) and as Eh increases, the solution is more
            oxidized (will accept electrons).

            average in lakes is ~300-500 mV

    C. pE – the electron activity of a solution
        - pE=-log[e-]
        - measured at pH=7; =16.9 Eh (so is related to electron potential)
        Large and positive in strongly oxidizing solutions (low electron activity)

    D. examples:  half-reactions

        1. Iron in oxygenated waters
        2. Sulfate reduction

        Redox reactions control the form of elements and the distribution of different forms of elements in lakes

V. Iron
    A. Forms of oxygen
        1. Fe²⁺ (ferrous) < ~300 mv Eh
        2. Fe³⁺ (ferric) > ~300 mv Eh
    B. Iron in oxygenated waters
    C. Iron in anoxic waters
    D. Significance –
        1. Fe³⁺ binds with P, element abundances are interrelated
        2. as a limiting nutrient

VI. The Redox ‘Battlefield’ - Most reactions are mediated by bacteria
    A. Players in the ‘battle’
        - One side – PS organisms take light energy and make reduced compounds (local chemical disequilibrium)
        - Other side – using these reduced compounds – heterotrophs extracting potential energy and attempting to reestablish
            equilibrium

    B. Who’s winning?
        - Example: in a lake at pH7, 25 degrees C, and oxygenated, predict that
        - C is in CO₂, HCO₃^-, CO₃^2-
        - N is in NO₃^- (no NH₄⁺)
        - S is in SO₄^2- (no H₂S)
        - oxidizers are winning the war
            if oxygen is available it will be used as a terminal electron acceptor because it has a large free energy difference

VII. Reactions
    A. How does this relate to energy?  The free energy depends on the energy difference between the reactants and the products.

    B. Gibbs free energy of reaction - Determines which reactions are most energetically favorable

        deltaG^o = standard free energy of formation
        R=gas constant (8.314 J K^-1 mol^-1)
        K_eq=equilibrium constant
        T=temperature (K)

       n=number of electrons involved in the reaction
        F=Faraday=23 kcal/mol
        deltaE=difference in redox potential in reaction

deltaG°’ = Sum (deltaG°’products) - Sum (deltaG°’reactants )

    C. Biology behind this
        The minimum energy necessary to be useful to organisms is deltaG^o=-7kcal/mol – that is what is necessary or ATP formation

    D. Redox reactions
        1. Aerobic Respiration (oxidation of organic matter)
            a) reductant CHO, oxidant O₂
            b) Reaction:

            c)   delta G^o = -686 kcal/mol

        2. Nitrogen
            a) NITRATE REDUCTION
                (i) reductant CHO, oxidant NO₃^-
                (ii) Reaction:  DISSIMILATORY NITRATE REDUCTION

                (iii)  delta G^o'  = -649 kcal/mol
                (iv) This is denitrification, reducing nitrate, NO₃^-, to N₂ gas
            b) Assimilatory nitrate reduction
                (i) Uptake of nitrate by an organism
                (ii) Reaction: 
            c) Nitrification – production of nitrate
                (i) Reactions:
                    a) delta G^o  = -65.7 kcal/mol
                    b) deltaG^o  = -17.5 kcal/mol
                (ii) So  delta G^o 1 = -65.7 kcal/mol and  delta G^o 2 = -17.5 kcal/mol
                (iii) This produces the source material for denitrification
        3. Iron
           IRON OXIDATION
            a) reductant Fe²⁺, oxidant O₂
            b) Reaction:

            c) delta G^o ’ = -10.6 kcal/mol

           IRON REDUCTION
            a) Reaction: 
            b) Can occur with some oxygen present, but not energetically favorable
            c)  delta G^o ’ = -300 kcal/mol

        4. Sulfur
           SULFATE REDUCTION
            a) reductant CHO, oxidant SO₄^2-
            b) Reaction: 
            c)  delta G^o ’ = -190 kcal/mol

       SULFIDE OXIDATION
            a) reductant HS^-, oxidant O₂
            b) Reaction: 
            c)  delta G^o ’ = -190 kcal/mol

        5. Methane
           METHANOGENESIS
            a) reductant H₂, oxidant CO₂
            b) Reaction: 
            c)  delta G^o ’ = -8.3 kcal/mol

           METHANE OXIDATION
            a) reductant CH₄, oxidant O₂
            b) Reaction: 
            c)  delta G^o ’ = -193.5 kcal/mol
            d) Whether methane can be oxidized with either SO₄^2- or NO₃^- is disputed

VIII. Oxidation of organic matter
    A. Sediments

 

    B. Other examples
            1. Sediment profile:
                   What happened from time 1 to time 2?
 
 

        2.

        3.

       4. You can also have the HS- level out before the sulfate concentration reaches low levels and stops decreasing.  Why?

            -Not because of presence of O₂ or NO₃^- because then SO₄^2- reduction wouldn't proceed
            -Often due to precipitation of pyrite: (see above)
                 for this to occur, there must be a source of Fe²⁺

        5.

 
 
 
 
 

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