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Oxidation-Reduction (Redox) Reactions

Oxidation-Reduction (Redox) Reactions. Measuring voltage. Standard potentials (E°) have been determined for how much voltage ( potential ) a reaction is capable of producing or consuming at standard conditions Nernst Equation. Standard Potentials . Written as reductions.

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Oxidation-Reduction (Redox) Reactions

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  1. Oxidation-Reduction (Redox) Reactions

  2. Measuring voltage • Standard potentials (E°) have been determined for how much voltage (potential) a reaction is capable of producing or consuming at standard conditions • Nernst Equation

  3. Standard Potentials Written as reductions Strong Reducing Agents The greater the E°, the more easily the substance reduced Strong Oxidizing Agents

  4. Redox Cell Salt bridge Pt wire electrode H2 gas (1 atm) Fe2+ and Fe3+ [H+] = 1 • H+ + e- ↔ ½ H2(g) • Fe3+ + e- ↔ Fe2+ ←: Pt wire removes electrons from half cell A →: Pt wire provides electrons to the solution

  5. Redox Cell using Platinum • Voltage meter registers difference in potential (E) between the 2 electrodes • Potential of SHE = 0, so E = potential of electrode in half-cell A • Defined as Eh; measured in volts • Eh is positive when [e-] in solution A less than [e-] in SHE • Eh is negative when [e-] in solution A greater than [e-] in SHE

  6. Eh as Master Variable • From electrochemistry: GR = -nF Eh • n = number of electrons • F = Faraday constant = 9.642 x 104 J / V∙mole • By convention, sign of Eh set for half-reaction written with e- on left side of equation • Can calculate E° = -GR° / nF (from Gf° values) • Determine GR° from the way the reaction is written (products – reactants)

  7. Eh as Master Variable • From electrochemistry: GR = -nF Eh • Re-write Nernst Equation: • At 25°C • Oxidized species on side where e- are

  8. Calculating Eh: Example • SO42- + Fe2+ + 8H+ + 8e- FeS + 4H2O

  9. Eh and redox pairs • Redox pair = 2 species of an element with different valences • e.g., SO42- - H2S; Fe3+ - Fe2+ • For every redox pair in a solution, an Eh can be defined • What if a solution has more than one redox pair? • An Eh can be calculated for each pair • All Eh’s will be equal if system at chemical equilibrium • But not so in nature, so different Eh values • Therefore, there is no unique Eh of a solution

  10. Measuring Eh • Eh is typically measured using a platinum (Pt) electrode + reference electrode • The reference electrode is a standard by which the Pt electrode measurement is made against • Ag:AgCl commonly used • Only responds to certain redox pairs • Doesn’t respond to solids • Best response to dissolved metals (e.g. Fe) • Better in reducing waters

  11. Computed vs. Measured Field Eh - Internal equilibrium not achieved - Computed Eh values do not agree with measured - Note vertical bands - Horizontal positions of the vertical bands chiefly reflect the standard E°

  12. Measured vs. Computed Eh - Samples with >1 redox pair - Points connected by vertical line derived from single sample - No internal redox equilibrium Lindberg, R.D. and D.D. Runnells (1984). Ground water redox reactions: an analysis of equilibrium state applied to Eh measurements and geochemical modeling. Science 225(4665):925-927.

  13. Measuring Eh • The Eh value is usually not very accurate in natural waters because of a lack of redoxequilibrium • One half of redox pair often below detection • It does usually give a good general idea of how oxidizing or reducing an environment is • Best to use Eh as a semi-quantitative measurement, giving you a relative idea of the redox potential of the water

  14. Eh – pH Diagrams • A different type of stability diagrams, but using Eh as variable instead of activity • Lines indicate equilibrium • Domains define areas of stability for minerals and aqueous species

  15. Water Stability Limits (H and O) in terms of pH and Eh • H2O(l) 2H+ + ½O2 + 2e- • From thermodynamic data, get: • ΔGR° = 2Gf°(H+) + ½Gf°(O2) + 2Gf°(e-) - Gf°(H2O) • ΔGR° = - Gf°(H2O) = 237.13 kJ/mole • ΔGR° = -nF E° • E° = ΔGR° / nF = 237.13 / [(2)(96.5)] = 1.23 V

  16. Water Stability Limits (H and O) • Eh = 1.23 + 0.0148 log[O2] – 0.059 pH • Establishes relationship among Eh, pH, and fO2 • f = fugacity; basically activity of a gas

  17. Water Stability Limits (H and O) • What are the stability limits of liquid water on Earth? • 2H2O(l) 2H2(g) + O2(g) • ΔGR° = 2 x 237.13 kJ/mole; K = 10-83.1 • At equilibrium, [O2][H2]2 = 10-83.1 • Total pressure of all gases occurring naturally at Earth’s surface must be ≤ 1 atm • If > 1, bubbles form in water exposed to the atmosphere and gases escape • So, fO2 and fH2 must each be ≤ 1 atm for liquid H2O to be stable • So if fH2 is at its maximum (1 atm), [O2] = 10-83.1

  18. Water Stability Limits (H and O) • So, fO2 can vary between 1 – 10-83.1 in equilibrium with H2O(l) at Earth’s surface • Eh = 1.23 + 0.0148 log[O2] – 0.059 pH • For O2 = 1 atm, Eh = 1.23 – 0.059 pH • For O2 = 10-83.1, Eh = 1.23 + 0.0148(-83.1) – 0.059 pH • Eh = 1.23 -1.23 - 0.059 pH • Eh = -0.059 pH

  19. Eh-pH Diagrams • Eh = 1.23 – 0.059 pH (fO2 = 1 atm) • Eh = -0.059 pH(fO2 = 10-83.1atm) • (y = mx + b) • These 2 equations plot as parallel straight lines on an Eh vs. pH plot (same slope) • And for any value of fO2, we would get additional parallel straight lines • Eh = 1.23 + 0.0148 log[O2] – 0.059 pH

  20. O2 and H2 are present in entire H2O stability range Oxidizing and reducing with respect to SHE Oxidizing environments may contain only small amounts of O2

  21. Oxygen • Most common and strongest oxidizing agent at the Earth’s surface is dissolved O2 • Consider pH = 7, Eh = +0.6 V • In groundwater environments, this is very oxidizing • Eh = 1.23 + 0.0148 log[O2] – 0.059 pH • [O2] = 10-14.6atm

  22. Oxidizing environment, but death to fish

  23. Eh-pH Diagrams • Positive Eh = oxidizing environments; tend to function as electron acceptors • Negative Eh = reducing environments; tend to function as electron donors

  24. Stability of Iron Compounds as a function of Eh and pH • Iron (Fe) is a common element on Earth, and is found in many forms and several valence states • Two main valence states are +2 (ferrous) and +3 (ferric); also 0 for native Fe • Solid phases: oxides, oxyhydroxides, sulfides, carbonates, silicates, native • Dissolved: usually Fe2+, Fe3+ in acidic, oxidizing waters • Common nuisance contaminant in groundwater • Important in biochemical processes; essential nutrient

  25. Plotting Fe reactions on Eh-pH Diagram • Select compounds and reactions of interest • Consider solubilities of iron oxide • Hematite (Fe2O3) • 2Fe2+ + 3H2O  Fe2O3 + 6H+ + 2e- • (note: by convention, e- always on right side of reactions) • GR = +126.99 kJ/mole • E = +0.66 V • Eh = 0.66 – 0.178 pH – 0.0592 log [Fe2+] • This produces a family of parallel lines (when [Fe2+] is defined expressing solubility of hematite in Eh-pH plane

  26. Solubility increases with decreasing pH and Eh; i.e., hematite dissolved under these conditions • [Fe2+] = 10-8 • [Fe2+] = 10-4 • [Fe2+] = 10-6

  27. Plotting Fe reactions on an Eh-pH Diagram • Next, magnetite (Fe3O4) and Fe2+ • 3Fe2+ + 4H2O  Fe3O4 + 8H+ + 2e- • GR = +169.82 kJ/mole • E = +0.88 V • Eh = 0.88 – 0.237 pH – 0.089 log [Fe2+]

  28. 10-4 • [Fe2+] = 10-6 10-8

  29. Equilibria between Fe2+ and 2 minerals How do we determine where each mineral dominates?

  30. Plotting Fe reactions on Eh-pH Diagram • Need to consider equilibrium between magnetite and hematite • 2 Fe3O4+ H2O 3Fe2O3 + 2H+ + 2e- • GR = +41.33 kJ/mole • E = +0.21 V • Eh = 0.21 – 0.0592 pH • [Fe2+] not a variable, don’t have to define its activity

  31. Equilibria between Fe2+ and 2 minerals

  32. Equilibria between Fe2+ and 2 minerals

  33. Plotting Fe reactions on Eh-pH Diagram • Iron can also be Fe3+ in solution • Consider relationship between Fe2+ and Fe3+ • Fe2+ Fe3+ + e- • Eh = 0.77 V; independent of pH • Constant Eh, horizontal line

  34. Plotting Fe reactions on Eh-pH Diagram • Iron can also be Fe3+ in solution • Fe2O3 + 6H+ 2Fe3+ + 3H2O • log [Fe3+] + 3 pH = -1.88 • Independent of Eh because no change in valence state (Fe in hematite is Fe3+ as well) • Constant pH, vertical line • Fe3O4 + 8H+ 3Fe3+ + 4H2O + e- • Eh = -0.55 + 0.473 pH

  35. Plotting Fe reactions on Eh-pH Diagram • Now let’s consider an iron carbonate mineral, siderite (FeCO3) • Fe is in the Fe2+ state (reduced); more common in subsurface • 3FeCO3 + H2O  Fe3O4 + 3CO2 + 2H+ + 2e- • Eh = 0.265 – 0.0592 pH + 0.0887 log [CO2] • At atmospheric PCO2 (3 x 10-4): • Eh = -0.048 – 0.0592 pH • Siderite-magnetite line plots below H2O stability limit • Thus siderite can’t precipitate unless PCO2 > atmospheric

  36. Plotting Fe reactions on Eh-pH Diagram • FeCO3 + 2H+ Fe2+ + CO2 + H2O • K = ([CO2] [Fe2+]) / [H+]2 • 2pH = 6.958 – log [CO2] - log[Fe2+] • Note: it is independent of Eh (no e- transfer), so if we set [CO2] and [Fe2+], it’s a vertical line • For [CO2] = 10-2 and [Fe2+] = 10-6 mol/L, pH = 7.48

  37. Evolution of Water Chemistry

  38. Source of dissolved species • Primarily from chemical weathering • Primary minerals + acid  secondary minerals + dissolved ions • The essential ingredients needed to produce chemical weathering are water and acid • Water sources: start with precipitation

  39. Chemical composition of precipitation (snow and rain) • Low TDS: ≤ 15 mg/L (water in contact with “rocks” for short period) • Acidic pH 5-6 naturally, in industrial area pH 3-4 (acid rain) • Dissolved ion composition variable, dependent on regional dust composition • e.g., in coastal areas Na+ and Cl- dominate (marine aerosols) • Regional limestones: Ca2+ and HCO3- dominate • Others: SO42- or NO3- can dominate • Also has dissolved gases: CO2 and O2 most important

  40. Soils • In most areas, soils are the first geologic unit to come into contact with precipitation • Soils have the highest rate of chemical weathering • Soil CO2 increases due to decay of organic matter • When water reaches water table, TDS has usually increased by more than 10x • Complex interactions involving geologic materials (rocks or sediments), water, plants, animals, microorganisms, gases • Role of biology is key: produce acids (CO2 and organic), decay of organics, affect soil structure, bioturbation

  41. Soil horizons

  42. Soil reactions • Throughout soil column: • CO2 produced by decay of organics and plant respiration • O2 consumed by decay of organics and redox reactions (Fe and S minerals) • N cycling • Soils continually produce acid (carbonic and organic)

  43. Soils and acidity • Soil CO2 is 10 – 100 X greater than in atmosphere, thus 10 – 100 X greater acidity • CO2 + H2O  H2CO3 H+ + HCO3- • Carbonic acid does most weathering • Organic acids: accounts for some weathering; also complexation with inorganic ions • Can affect solute transport mechanisms

  44. Plants/Animals • Plants take up and release inorganic elements as nutrients • Seasonal affects • On a seasonal basis, element uptake does not equal its release • But on an annual basis, uptake approximately equals release • Over decadal-century time frame, uptake approximately equals release (steady state) • No steady state if crops are harvested; this is why fertilizers must be added

  45. Generalized nutrient requirements of plants (molar) • 800 CO2 • 6 NH4+ • 4 Ca2+ • 1 Mg2+ • 2 K+ • 1 Al(OH)2+ • 1 Fe2+ • 2 NO3- • 1 H2PO4- • 1 SO42- • H2O • Micronutrients: B, Cu, Mn, Mo, Zn, Cl- • Na+ only major ion not involved in biological activity

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