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Oxidation and Reduction

Oxidation and Reduction. Lecture 9. Redox in Aqueous Solutions. Redox reactions occur over a wide range of conditions: from groundwaters to magma. They are approached differently. We begin with aqueous solutions. Electrochemical Cells. A simple redox reaction would be:

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Oxidation and Reduction

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  1. Oxidation and Reduction Lecture 9

  2. Redox in Aqueous Solutions Redox reactions occur over a wide range of conditions: from groundwaters to magma. They are approached differently. We begin with aqueous solutions.

  3. Electrochemical Cells • A simple redox reaction would be: • We want to know ∆G of the reaction. Measuring energy of it in electrochemical cell might be good approach. • However, such a cell can only measure exchange of electrons (e.g., between Zn and Cu) • We really want to know are energies for individual redox reactions such as:

  4. Hydrogen Scale Potential • We assign a potential of 0 for the reaction: ½H2(g) = Haq+ + e- • in practice one side has Pt electrode in H2 gas, the other acid with aH+ = 1. • Then for the reaction • The potential is assigned to • Potentials measured in this way are called hydrogen scale potentials, written EH and have units of volts.

  5. Table 3.3 EH˚ and pe˚ for half-cell reactions

  6. EH and ∆G • Electrochemical energy is a form of free energy. EH is related to ∆Gr by: ∆Gr = -zFEH • where F is the Faraday constant (96,485 coulombs) and converts volts to joules. • and ∆G˚ = -zFE˚ • Values of E˚ available in compilations (e.g., Table 3.3) • Since • then • This is known as the Nernst Equation.

  7. pe • Consider again the reaction: • The equilibrium constant expression for this reaction is ? • In log form: • We define pe as: • So

  8. Standard State pe and Relation to EH • Continuing with the reaction • In an aqueous solution, the standard state activities are? • Therefore pe˚ = log K • More generally, • So for this reaction: • pe is related to EH as:

  9. What pe is really telling us • We have defined pe as the negative log of the activity of the electron. So a high pe means a low activity and concentration of electrons in our system. A low concentration of electrons implies an oxidized system; a high concentration (and low pe) implies a reduced system. • Same is true of EH. • So these are parameters that tell us about the redox state of our system (just as pH tells us about acidity).

  10. Speaking of pe and pH… • A commonly used diagram to illustrate chemical variation in aqueous solutions is the pe-pH diagram (or EH-pH) • Water only stable over limited range, so we start by setting boundaries. ½O2(g) + 2e- + 2H+ = H2O • In the standard state: pe = 20.78-pH • The is a line with intercept of 20.78 and slope of -1. • Similarly: H+ + e- = ½H2(g) • and pe = -pH

  11. pe-pH Diagrams • To construct the diagrams • Write a reaction relating species of interest. • Redox reactions should contain e- • pH dependent reactions should contain H+ • Write the equilibrium constant expression. • Get in log form, solve for pe with equation of the form pe = a + bpH • Find or calculate value of log K.

  12. Drawing stability boundaries • Now consider: • For equal activities of the two species, • pe = log K • (horizontal line with intercept = K) • Next Fe3+–Fe(OH)2+: • Fe3+ + H2O = Fe(OH)2+ + H+ Use H+ rather than OH-!

  13. Fe2+–Fe(OH)2+ • Our reaction is: • Equilibrium constant expression is: • In the form we want: • We can write it as as the sum of two reactions, • we sum • to yield • The log equilibrium constant of the net reaction is the sum of the equilibrium constants of the two.

  14. Line 5 has a slope of -1 and an intercept of log K. We can also use pe-pH diagrams to illustrate stability of solid phases in presence of solution. In this case, we must choose concentration.

  15. More about pe-pH diagrams • pe-pH diagrams are a kind of stability or predominance diagram. • They differ from phase diagrams because lines indicate not phase boundaries, but equal concentrations. • There is only 1 phase in this this diagram – an aqueous solution. • Regions are regions of predominance. • The aqueous species continue to exist beyond their fields, but their concentrations drop off exponentially.

  16. Environmental Interpretation of pe-pH

  17. Redox in Magmas

  18. Oxygen Fugacity • Igneous geochemists use oxygen fugacity ƒO2 to represent the redox state of the system. Hence, the oxidation of ferrous to ferric iron would be written as: 2FeO + O2(g) = Fe2O3 • For example, oxidation of magnetite to hematite: 2Fe3O4+ ½O2(g) = 3Fe2O3 • (Actually, there isn’t much O2 gas in magmas. Reaction more likely mediated by water and hydrogen).

  19. Redox in Magmatic Systems • For magnetite-hematite • 2Fe3O4+ ½O2(g) = 3Fe2O3 • assuming the two are pure solids • At a temperature such as 1000K

  20. Oxygen Fugacity Buffers • The log ƒO2– T diagram is a phase diagram illustrating boundaries of phase stability. The two phases coexist only at the line. • Reactions such as magnetite-hematite (or iron-wüstite or fayalite-magnetite-quartz) are buffers. • For example, if we bleed O2 into a magma containing magnetite, the ƒO2 cannot rise above the line until all magnetite is converted to hematite (assuming equilibrium!)

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