REDOX REACTIONS

I. Valence – atoms have different valences (degree of negative or positive charge) determined by whether they have
    enough protons to balance their electrons.  Elements and molecules have tendencies to associate or dissociate
    with electrons (change valence)

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 : dissociation of Hydrogen into proton and electron

    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-, CO32-
        - N is in NO3- (no NH4+)
        - S is in SO42- (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

        deltaGo = standard free energy of formation
        R=gas constant (8.314 J K-1 mol-1)
        Keq=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 deltaGo=-7kcal/mol – that is what is necessary or ATP formation

redox reactions 
 

    D. Redox reactions
        1. Aerobic Respiration (oxidation of organic matter)
            a) reductant CHO, oxidant O2
            b) Reaction:
reaction for aerobic respiration
            c)   delta Go = -686 kcal/mol

        2. Nitrogen
            a) NITRATE REDUCTION
                (i) reductant CHO, oxidant NO3-
                (ii) Reaction:  DISSIMILATORY NITRATE REDUCTION
dissimilatory nitrate reduction
                (iii)  delta Go'  = -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: assimilatory nitrate reduction
            c) Nitrification – production of nitrate
                (i) Reactions:
                    a) nitrification of ammonium to nitritedelta Go  = -65.7 kcal/mol
                    b) conversion of nitrite to nitratedeltaGo  = -17.5 kcal/mol
                (ii) So  delta Go 1 = -65.7 kcal/mol and  delta Go 2 = -17.5 kcal/mol
                (iii) This produces the source material for denitrification
        3. Iron
           IRON OXIDATION
            a) reductant Fe2+, oxidant O2
            b) Reaction:
iron oxidation
formation of iron hydroxides

            c) delta Go ’ = -10.6 kcal/mol

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

        4. Sulfur
           SULFATE REDUCTION
            a) reductant CHO, oxidant SO42-
            b) Reaction: sulfate reduction
            c)  delta Go ’ = -190 kcal/mol

       SULFIDE OXIDATION
            a) reductant HS-, oxidant O2
            b) Reaction: sulfide oxidation
            c)  delta Go ’ = -190 kcal/mol

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

           METHANE OXIDATION
            a) reductant CH4, oxidant O2
            b) Reaction: methane oxidation
            c)  delta Go ’ = -193.5 kcal/mol
            d) Whether methane can be oxidized with either SO42- or NO3- is disputed

VIII. Oxidation of organic matter
    A. Sediments
sequence of redox reactions during oxidation of organic matter in the sediments
 

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

example of sulfate reduction
        2.
changes in methane with depth in sediment

        3.

changes in concentrations of HS- with sulfate reduction; effects on alkalinity

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

            -Not because of presence of O2 or NO3- because then SO42- reduction wouldn't proceed
            -Often due to precipitation of pyrite: (see above)
                 for this to occur, there must be a source of Fe2+

        5.
other redox reactions that could increase alkalinity
 
 
 
 
 

 Return to Limnology Lecture homepage

 Return to K.L. Schulz's homepage