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5 Redox Processes.doc

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University of Toronto St. George
Earth Sciences
John Ferris

5 Redox Processes 5.0 Introduction The importance of redox (i.e., oxidation-reduction) processes in biogeochemistry cannot be over emphasized. These are reactions that involve the transfer of electrons between one or more chemical species. Life would not exist without oxidation-reduction reactions, and the intense biological and geochemical cycling of redox sensitive elements that essentially define the uniqueness of the Earth system would not occur. Everything we appreciate about our planet depends on oxidation-reduction reactions. Many elements occur in nature in more than one oxidation state. The major elements with such redox behavior include H, O,C,S,N, and Fe. In some systems, Mn is also a major redox element. As a general rule, most reactions that involve electrons also involve protons. Oxidation usually releases protons or acidity. This is a basic cause of acid mine drainage. Conversely, reduction generally consumes protons, and the pH rises. 5.1 Oxidation-Reduction Reactions In redox reactions, electrons are transferred from an electron-donor or reductant to an electron acceptor or oxidant. The surrender of electrons from a compound is referred to as oxidation, and the acceptance of electrons by a different compound is reduction. In a complete redox reaction, the reductant is transformed into a conjugate oxidized species and the oxidant is transformed into a conjugate reduced species. Consider a general oxidation-reduction reaction in which reduced species red is 1 oxidized when it donates one electron per molecule to oxidized species ox . The 2verall reaction can be taken to be the result of two half-cell reactions as follows Oxidation: red1=conjugateox + e1 − (1) Reduction: ox2+ e = conjugatered 2 (2) Overall redox reaction: red1+ ox 2conjugateox +conj1gatered 2 (3) The strength of an oxidant or reductant is determined electrochemically by their relative capacities to accept or donate electrons, respectively from a reductant. As such, a strong oxidant is a potent electron acceptor, whereas a weak oxidant has a poor capacity to accept electrons. Similarly, a strong reductant donates electrons readily, whereas electrons are not given up easily by a weak reductant. These sweeping generalizations emphasize the importance of adopting a reference scale to measure the actual strength of 1 oxidants and reductants. The conventional approach is to use the H /H ha2f-cell reaction H (Eqn. 4) in which the corresponding equilibrium constant K is assigned a standard value of unity. + - 2H + 2e = H 2 (4) K H = {H }2 =1 {H }{e }− 2 (5) In accordance with IUPAC guidelines, all half cell reactions are written as n+ - reduction reactions with oxidized species A and n electrons eon the left and reduced species B on the right n+ - A + ne = B (6) The corresponding mass action relationship for the half-cell reaction in which oxidized species A is reduced to species B is A/ B {B} K = n+ − n (7) {A }{e } A/B + where K is the equilibrium constant for the reduction reaction. Combining the H /H 2 half-cell reaction with Eqn. 6 gives the overall reaction An+ + n H = B + n H + 2 2 2 (8) with a mass action expression of R {B}{H }+ n/ 2 K A/ B A/ B K = n+ n/2 = H = K (9) {A }pH 2 K This relationship shows that the standard H /H h2lf-cell reaction makes the equilibrium constant for the overall oxidation-reduction reaction K equivalent to the equilibrium n+ A/B constant of the A /B half-cell K . n+ As the mass action expresn+on for the A /B half-cell (Eqn. 7) indicates that the greater the capacity of oxidant A to serve as an electron acceptor, the greater the value of K . Taking common logarithms of Eqn. 7 and rearranging gives 1  {B}  −log{e }= pe = logK A/ B−log  (10) n  {A }+  1 International Union of Pure and Applied Chemistry 2 where, as the negative common logarithm of electron activity, pe is by analogy similar to the definition of pH = -log{H }. When {B} = {A }+ {B} =1 (11) {A } such that 1 A/ B 0 pe = logK = pe (12) n This gives from Eqns. 3 and 5 0 1 {B} pe = pe − log n+ (13) n {A } Considering that − ΔG 0 lnK = 2.3logK A/ = (14) RT then from Eqn. 5 0 − ΔG 0 pe = (15) 2.303nRT and pe = − ΔG (16) 2.303nRT Combining Eqns. 13, 15 and 16 gives the change in free energy for the half-cell reaction illustrated by Eqn. 1. ΔG = ΔG + 2.303RT log {B} (17) {A } In most aqueous systems, electron transfer in a complete oxidation-reduction reaction takes place on a molecular level, and there is no practical way of measuring the actual magnitude of electron activity; however, oxidation and reduction steps can be carried out at separate electrodes if the electrons yielded by the oxidation half-cell are conveyed to the electron acceptor at another electrode through an electrical circuit. The 3 free energy difference associated with the transfer of electrons between the two electrodes is given by the fundamental electrochemical relationship ΔG = −nFEm (18) and at standard state ΔG = −nFEm 0 (19) Where F is the Faraday constant and Em is the potential difference in volts between the two electrodes of the electrical cell Because the absolute value of Em depends fundamentally on two different half- cell reactions at separate electrodes, it is advantageous to adopt a standard reference electrode against which to measure the tendency of redox active substances to donate or accept electrons. This is done by according the standard hydrogen (half-cell) electrode (SHE) with a potential of zero with unit electron, H and H ac2ivities. Relative to the SHE, Em values are expressed more explicitly in terms of Eh, the potential difference of the electrical half-cell relative to the SHE, which gives ΔG = −nFEh (20) and additionally ΔG = −nFEh 0 (21) Another important outcome from Eqn. 21 is that the standard potential of a half-cell 0 0 reaction Eh can be calculated directly from ΔG , or indirectly from standard free energy of formation values (Table 5.1). Substitution of Eqn. 20 and Eqn. 21 into Eqn. 17 and rearranging yields the Nernst Equation 2.303RT {A }+  Eh = Eh +0 log  (22) nF  {B}  This can be generalized for any half cell reaction 2.303RT  activityproduct of oxidizedspecies  Eh = Eh +0 log  (23) nF  activityproduct of reducedspecies  4 Table 5.1: Examples of common half-cell reactions with corresponding Eh , ΔG ,0 and log K values. Conjugate 0 ΔG 0 Oxidant Reductant Eh (V) (kJ/mol) log K O 2 H 2 1.231 -118.7 20.8 - NO 3 N 2 1.248 -120.5 21.1 NO - NO - 0.840 -81.1 14.2 3 2 NO 3 NH 4 0.882 -85.1 14.9 3+ 2+ Fe Fe 0.770 -74.3 13.0 Fe(OH) Fe 2+ 1.065 -102.8 18.0 3 SO 42- H 2 0.311 -30.0 5.3 2- - SO 4 HS 0.251 -24.3 4.3 CO 2 CH 4 0.170 -16.4 2.9 CO 2 CH O2 -0.071 6.9 -1.2 + H H 2 0.000 0.0 0.0 5.2 Redox Stability Diagrams As illustrated by Eqn. 8, redox reactions often involve consumption or production of protons. This means that Eh can be expressed as a function of pH on the basis of the Nernst Equation (Eqn. 15). Theoretical expressions developed from this approach provide a convenient way to display stability relationships between oxidized and reduced species in redox reactions; however, an important point to remember is t
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