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Oxidation-Reduction (Redox) Reactions 1. Measuring voltage Standard potentials (E°) have been determined for how much voltage (potential) a reaction is.

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Presentation on theme: "Oxidation-Reduction (Redox) Reactions 1. Measuring voltage Standard potentials (E°) have been determined for how much voltage (potential) a reaction is."— Presentation transcript:

1 Oxidation-Reduction (Redox) Reactions 1

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 2

3 Standard Potentials Half-ReactionE 0 (V) Li + (aq) + e - → Li (s) K + (aq) + e - → K (s) Ba 2+ (aq) + 2 e - → Ba (s) Sr 2+ (aq) + 2 e - → Sr (s) Ca 2+ (aq) + 2 e - → Ca (s) Na + (aq) + e - → Na (s) Mg 2+ (aq) + 2 e - → Mg (s) Be 2+ (aq) + 2 e - → Be (s) Al 3+ (aq) + 3 e - → Al (s) Mn 2+ (aq) + 2 e - → Mn (s) H 2 O + 2 e - → H 2(g) + 2 OH - (aq) Zn 2+ (aq) + 2 e - → Zn (s) Cr 3+ (aq) + 3 e - → Cr (s) Fe 2+ (aq) + 2 e - → Fe (s) Cd 2+ (aq) + 2 e - → Cd (s) PbSO 4(s) + 2 e - → Pb (s) + SO 4 2- (aq) Co 2+ (aq) + 2 e - → Co (s) Ni 2+ (aq) + 2 e - → Ni (s) Sn 2+ (aq) + 2 e - → Sn (s) Pb 2+ (aq) + 2 e - → Pb (s) H + (aq) + 2 e - → H 2(g) 0 Half-ReactionE 0 (V) 2 H + (aq) + 2 e - → H 2(g) 0 Sn 4+ (aq) + 2 e - → Sn 2+ (aq) 0.13 Cu 2+ (aq) + e - → Cu + (aq) 0.13 SO 4 2- (aq) + 4 H + (aq) + 2 e - → SO 2(g) + 2 H 2 O0.20 AgCl (s) + e - → Ag (s) + Cl - (aq) 0.22 Cu 2+ (aq) + 2 e - → Cu (s) 0.34 O 2(g) + 2 H e - → 4 OH - (aq) 0.40 I 2(s) + 2 e - → 2 I - (aq) 0.53 MnO 4 - (aq) + 2 H 2 O + 3 e - → MnO 2(s) + 4 OH - (aq) 0.59 O 2(g) + 2 H + (aq) + 2 e - → H 2 O 2(aq) 0.68 Fe 3+ (aq) + e - → Fe 2+ (aq) 0.77 Ag + (aq) + e - → Ag (s) 0.80 Hg 2 2+ (aq) + 2 e - → 2 Hg (l) Hg 2+ (aq) + 2 e - → Hg 2 2+ (aq) 0.92 NO 3 - (aq) + 4 H + (aq) + 3 e - → NO (g) + 2 H 2 O0.96 Br 2(l) + 2 e - → 2 Br - (aq) 1.07 O 2(g) + 4 H + (aq) + 4 e - → 2 H 2 O1.23 MnO 2(s) + 4 H + (aq) + 2 e - → Mn 2+ (aq) + 2 H 2 O1.23 Cr 2 O 7 2- (aq) + 14 H + (aq) + 6 e - → 2 Cr 3+ (aq) + 7 H 2 O1.33 Cl 2(g) + 2 e - → 2 Cl - (aq) 1.36 Au 3+ (aq) + 3 e - → Au (s) 1.50 MnO 4 - (aq) + 8 H + (aq) + 5 e - → Mn 2+ (aq) + 4 H 2 O1.51 Ce 4+ (aq) + e - → Ce 3+ (aq) 1.61 PbO 2(s) + 4H + (aq) + SO 4 2- (aq) + 2e - → PbSO 4(s) + 2H 2 O1.70 H 2 O 2(aq) + 2 H + (aq) + 2 e - → 2 H 2 O1.77 Co 3+ (aq) + e - → Co 2+ (aq) 1.82 O 3(g) + 2 H + (aq) + 2 e - → O 2(g) + H 2 O2.07 F 2(g) + 2 e > F - (aq) 2.87 The greater the E°, the more easily the substance reduced Strong Reducing Agents Strong Oxidizing Agents Written as reductions

4 4 Pt wire electrode H 2 gas (1 atm) Salt bridge [H + ] = 1 Fe 2+ and Fe 3+ Fe 3+ + e - ↔ Fe 2+ ←: Pt wire removes electrons from half cell A →: Pt wire provides electrons to the solution H + + e - ↔ ½ H 2(g) Redox Cell

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 5

6 Eh as Master Variable From electrochemistry:  G R = -nF Eh –n = number of electrons –F = Faraday constant = x 10 4 J / V∙mole –By convention, sign of Eh set for half-reaction written with e - on left side of equation –Can calculate E° = -  G R ° / nF (from  G f ° values) Determine  G R ° from the way the reaction is written (products – reactants) 6

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

8 Calculating Eh: Example SO Fe H + + 8e -  FeS + 4H 2 O 8

9 Eh and redox pairs Redox pair = 2 species of an element with different valences –e.g., SO H 2 S; Fe 3+ - Fe 2+ 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 9

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 10

11 Computed vs. Measured Field Eh 11 - 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 12 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): Samples with >1 redox pair - Points connected by vertical line derived from single sample - No internal redox equilibrium

13 Measuring Eh The Eh value is usually not very accurate in natural waters because of a lack of redox equilibrium –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 13

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 14

15 Water Stability Limits (H and O) in terms of pH and Eh H 2 O (l)  2H + + ½O 2 + 2e - From thermodynamic data, get: – –ΔG R ° = 2G f °(H + ) + ½G f °(O2) + 2G f °(e - ) - G f °(H 2 O) –ΔG R ° = - G f °(H 2 O) = kJ/mole –ΔG R ° = -nF E° –E° = ΔG R ° / nF = / [(2)(96.5)] = 1.23 V 15

16 Water Stability Limits (H and O) Eh = log[O 2 ] – pH Establishes relationship among Eh, pH, and f O2 –f = fugacity; basically activity of a gas 16

17 Water Stability Limits (H and O) What are the stability limits of liquid water on Earth? –2H 2 O (l)  2H 2(g) + O 2(g) –ΔG R ° = 2 x kJ/mole; K = –At equilibrium, [O 2 ][H 2 ] 2 = 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, f O2 and f H2 must each be ≤ 1 atm for liquid H 2 O to be stable –So if f H2 is at its maximum (1 atm), [O 2 ] =

18 Water Stability Limits (H and O) So, f O2 can vary between 1 – in equilibrium with H 2 O (l) at Earth’s surface Eh = log[O 2 ] – pH –For O 2 = 1 atm, Eh = 1.23 – pH –For O 2 = , Eh = (-83.1) – pH Eh = pH Eh = pH 18

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

20 20 Oxidizing and reducing with respect to SHE O 2 and H 2 are present in entire H 2 O stability range Oxidizing environments may contain only small amounts of O 2

21 Oxygen Most common and strongest oxidizing agent at the Earth’s surface is dissolved O 2 Consider pH = 7, Eh = +0.6 V –In groundwater environments, this is very oxidizing –Eh = log[O 2 ] – pH –[O 2 ] = atm 21

22 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 23

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25 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 Fe 2+, Fe 3+ in acidic, oxidizing waters –Common nuisance contaminant in groundwater –Important in biochemical processes; essential nutrient 25

26 Plotting Fe reactions on Eh-pH Diagram Select compounds and reactions of interest Consider solubilities of iron oxide Hematite (Fe 2 O 3 ) 2Fe H 2 O  Fe 2 O 3 + 6H + + 2e - –(note: by convention, e - always on right side of reactions) –  G R  = kJ/mole –E  = V –Eh = 0.66 – pH – log [Fe 2+ ] –This produces a family of parallel lines (when [Fe 2+ ] is defined expressing solubility of hematite in Eh-pH plane 26

27 27 Solubility increases with decreasing pH and Eh; i.e., hematite dissolved under these conditions [Fe 2+ ] = [Fe 2+ ] = [Fe 2+ ] = 10 -8

28 Plotting Fe reactions on an Eh-pH Diagram Next, magnetite (Fe 3 O 4 ) and Fe 2+ 3Fe H 2 O  Fe 3 O 4 + 8H + + 2e - –  G R  = kJ/mole –E  = V –Eh = 0.88 – pH – log [Fe 2+ ] 28

29 [Fe 2+ ] = 10 -6

30 Equilibria between Fe 2+ and 2 minerals 30 How do we determine where each mineral dominates?

31 Plotting Fe reactions on Eh-pH Diagram Need to consider equilibrium between magnetite and hematite 2 Fe 3 O 4 + H 2 O  3Fe 2 O 3 + 2H + + 2e - –  G R  = kJ/mole –E  = V –Eh = 0.21 – pH –[Fe 2+ ] not a variable, don’t have to define its activity 31

32 32 Equilibria between Fe 2+ and 2 minerals

33 33 Equilibria between Fe 2+ and 2 minerals

34 Plotting Fe reactions on Eh-pH Diagram Iron can also be Fe 3+ in solution Consider relationship between Fe 2+ and Fe 3+ Fe 2+  Fe 3+ + e - –Eh = 0.77 V; independent of pH Constant Eh, horizontal line 34

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36 Plotting Fe reactions on Eh-pH Diagram Iron can also be Fe 3+ in solution Fe 2 O 3 + 6H +  2Fe H 2 O –log [Fe 3+ ] + 3 pH = –Independent of Eh because no change in valence state (Fe in hematite is Fe 3+ as well) Constant pH, vertical line Fe 3 O 4 + 8H +  3Fe H 2 O + e - –Eh = pH 36

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38 Plotting Fe reactions on Eh-pH Diagram Now let’s consider an iron carbonate mineral, siderite (FeCO 3 ) Fe is in the Fe 2+ state (reduced); more common in subsurface 3FeCO 3 + H 2 O  Fe 3 O 4 + 3CO 2 + 2H + + 2e - –Eh = – pH log [CO 2 ] –At atmospheric P CO2 (3 x ): Eh = – pH Siderite-magnetite line plots below H 2 O stability limit Thus siderite can’t precipitate unless P CO2 > atmospheric 38

39 Plotting Fe reactions on Eh-pH Diagram FeCO 3 + 2H +  Fe 2+ + CO 2 + H 2 O –K = ([CO 2 ] [Fe 2+ ]) / [H + ] 2 –2pH = – log [CO 2 ] - log[Fe 2+ ] Note: it is independent of Eh (no e - transfer), so if we set [CO 2 ] and [Fe 2+ ], it’s a vertical line –For [CO 2 ] = and [Fe 2+ ] = mol/L, pH =

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42 Evolution of Water Chemistry 42

43 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 43

44 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: Ca 2+ and HCO 3 - dominate –Others: SO 4 2- or NO 3 - can dominate Also has dissolved gases: CO 2 and O 2 most important 44

45 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 CO 2 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 (CO 2 and organic), decay of organics, affect soil structure, bioturbation 45

46 Soil horizons 46 O horizon: surface layer predominately organic matter A horizon: highly weathered, high organic matter, Fe/Al leached; high N Zone of Leaching B horizon: accumulated clay, Fe/Al hydroxides, humus (stable organic matter; gaseous diffusion and aqueous transport between B and C Zone of Accumulation C horizon: altered parent material, solute and gases exchange with saturated zone; periodically saturated when water table high Partly decomposed and unaltered bedrock Saturated zone

47 Soil reactions Throughout soil column: –CO 2 produced by decay of organics and plant respiration –O 2 consumed by decay of organics and redox reactions (Fe and S minerals) –N cycling –Soils continually produce acid (carbonic and organic) 47

48 Soils and acidity Soil CO 2 is 10 – 100 X greater than in atmosphere, thus 10 – 100 X greater acidity –CO 2 + H 2 O  H 2 CO 3  H + + HCO 3 - –Carbonic acid does most weathering Organic acids: accounts for some weathering; also complexation with inorganic ions –Can affect solute transport mechanisms 48

49 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 49

50 Generalized nutrient requirements of plants (molar) 800 CO 2 6 NH Ca 2+ 1 Mg 2+ 2 K + 1 Al(OH) Fe 2+ 2 NO H 2 PO SO 4 2- H 2 O Micronutrients: B, Cu, Mn, Mo, Zn, Cl - Na + only major ion not involved in biological activity 50


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