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

Standard Potentials

Half-Reaction E0 (V)Li+(aq) + e- → Li(s) -3.05K+

(aq) + e- → K(s) -2.93Ba2+

(aq) + 2 e- → Ba(s) -2.90Sr2+

(aq) + 2 e- → Sr(s) -2.89Ca2+

(aq) + 2 e- → Ca(s) -2.87Na+

(aq) + e- → Na(s) -2.71Mg2+

(aq) + 2 e- → Mg(s) -2.37Be2+

(aq) + 2 e- → Be(s) -1.85Al3+

(aq) + 3 e- → Al(s) -1.66Mn2+

(aq) + 2 e- → Mn(s) -1.182 H2O + 2 e- → H2(g) + 2 OH-

(aq) -0.83Zn2+

(aq) + 2 e- → Zn(s) -0.76Cr3+

(aq) + 3 e- → Cr(s) -0.74Fe2+

(aq) + 2 e- → Fe(s) -0.44Cd2+

(aq) + 2 e- → Cd(s) -0.40PbSO4(s) + 2 e- → Pb(s) + SO4

2-(aq) -0.31

Co2+(aq) + 2 e- → Co(s) -0.28

Ni2+(aq) + 2 e- → Ni(s) -0.25

Sn2+(aq) + 2 e- → Sn(s) -0.14

Pb2+(aq) + 2 e- → Pb(s) -0.13

2 H+(aq) + 2 e- → H2(g) 0

Half-Reaction E0 (V)2 H+

(aq) + 2 e- → H2(g) 0Sn4+

(aq) + 2 e- → Sn2+(aq) 0.13

Cu2+(aq) + e- → Cu+

(aq) 0.13SO4

2-(aq) + 4 H+

(aq) + 2 e- → SO2(g) + 2 H2O 0.20AgCl(s) + e- → Ag(s) + Cl-(aq) 0.22Cu2+

(aq) + 2 e- → Cu(s) 0.34O2(g) + 2 H2 + 4 e- → 4 OH-

(aq) 0.40I2(s) + 2 e- → 2 I-

(aq) 0.53MnO4

-(aq) + 2 H2O + 3 e- → MnO2(s) + 4 OH-

(aq) 0.59O2(g) + 2 H+

(aq) + 2 e- → H2O2(aq) 0.68Fe3+

(aq) + e- → Fe2+(aq) 0.77

Ag+(aq) + e- → Ag(s) 0.80

Hg22+

(aq) + 2 e- → 2 Hg(l) 0.852 Hg2+

(aq) + 2 e- → Hg22+

(aq) 0.92NO3

-(aq) + 4 H+

(aq) + 3 e- → NO(g) + 2 H2O 0.96Br2(l) + 2 e- → 2 Br-

(aq) 1.07O2(g) + 4 H+

(aq) + 4 e- → 2 H2O 1.23MnO2(s) + 4 H+

(aq) + 2 e- → Mn2+(aq) + 2 H2O 1.23

Cr2O72-

(aq) + 14 H+(aq) + 6 e- → 2 Cr3+

(aq) + 7 H2O 1.33Cl2(g) + 2 e- → 2 Cl-(aq) 1.36Au3+

(aq) + 3 e- → Au(s) 1.50MnO4

-(aq) + 8 H+

(aq) + 5 e- → Mn2+(aq) + 4 H2O 1.51

Ce4+(aq) + e- → Ce3+

(aq) 1.61PbO2(s) + 4H+

(aq) + SO42-

(aq) + 2e- → PbSO4(s) + 2H2O 1.70H2O2(aq) + 2 H+

(aq) + 2 e- → 2 H2O 1.77Co3+

(aq) + e- → Co2+(aq) 1.82

O3(g) + 2 H+(aq) + 2 e- → O2(g) + H2O 2.07

F2(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

Pt wireelectrode H2 gas (1 atm)

Salt bridge

[H+] = 1Fe2+ andFe3+

Fe3+ + e- ↔ Fe2+

←: Pt wire removes electrons from half cell A

→: Pt wire provides electrons to the solution

H+ + e- ↔ ½ H2(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

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., SO4

2- - 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 positionsof the vertical bands chiefly reflect the standard E°

12

Measured vs. Computed Eh

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.

- 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

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 = 10-83.1 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 fO2, we would get

additional parallel straight lines– Eh = 1.23 + 0.0148 log[O2] – 0.059 pH

20

Oxidizing and reducingwith respect to SHE

O2 and H2 are presentin entire H2O stabilityrange

Oxidizing environmentsmay contain only smallamounts 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] = 10-14.6 atm

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

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 Fe2+, Fe3+ in acidic, oxidizing

waters– Common nuisance contaminant in groundwater– Important in biochemical processes; essential nutrient

26

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

27

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++ Hematite

25°C

Walt Mon Feb 06 2006

Dia

gram

Fe+

+,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1;

Sup

pres

sed:

FeO

(c),

Goe

thite

, F

e(O

H) 3

(ppd

), F

e(O

H) 3

, F

e++

+,

Fe(

OH

) 4- , F

e(O

H) 2+

, F

eOH+

+,

Fe

2(OH

) 2++

++,

Fe 3

(OH

) 4(5

+),

Mag

netit

e, F

e(O

H) 2

(ppd

)

Solubility increases with decreasing pH and Eh;i.e., hematite dissolved under these conditions

[Fe2+] = 10-6

[Fe2+] = 10-4

[Fe2+] = 10-8

28

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+]

29

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++

Magnetite

25°C

Walt Mon Feb 06 2006

Dia

gram

Fe+

+,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1;

Sup

pres

sed:

FeO

(c),

Hem

atite

, G

oeth

ite,

Fe(

OH

) 3(p

pd),

Fe(

OH

) 3,

Fe+

++,

Fe(

OH

) 4- , F

e(O

H) 2+

,

FeO

H+

+,

Fe 2

(OH

) 2++

++,

Fe 3

(OH

) 4(5

+)

10-4

10-8

[Fe2+] = 10-6

30

Equilibria between Fe2+ and 2 minerals

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 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

32

Equilibria between Fe2+ and 2 minerals

33

Equilibria between Fe2+ and 2 minerals

34

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

35

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++

Fe+++

25°C

Walt Mon Feb 06 2006

Dia

gram

Fe+

++,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1;

Sup

pres

sed:

FeO

(c),

Goe

thite

, F

e(O

H) 3

(ppd

), F

e(O

H) 3

, F

e(O

H) 4- ,

Fe(

OH

) 2+,

FeO

H+

+,

Fe 2

(OH

) 2++

++,

Fe

3(OH

) 4(5+),

Mag

netit

e, F

e(O

H) 2

(ppd

), F

eOH

+,

Fe(

OH

) 2,

Fe(

OH

) 3-

36

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

37

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++

Fe+++

Hematite

Magnetite25°C

Walt Mon Feb 06 2006

Dia

gram

Fe+

+,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1;

Sup

pres

sed:

FeO

(c)

38

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

39

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

40

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++

Fe+++

Hematite

Magnetite

Siderite

25°C

Walt Mon Feb 06 2006

Dia

gram

Fe+

+,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1,

a [

HC

O3- ]

=

10–

2;

Sup

pres

sed:

FeO

(c),

FeC

H3C

OO

+,

Fe(

CH

3C

OO

) 3,

Fe(

CH

3C

OO

) 2+,

FeC

H3C

OO

++

41

0 2 4 6 8 10 12 14

–.5

0

.5

1

pH

Eh

(vo

lts)

Fe++

Fe+++

Hematite

Magnetite

Pyrite

FeO(c)

Troilite

25°C

Walt Fri May 05 2006

Dia

gram

Fe+

+,

T

=

25 °

C ,

P

=

1.01

3 ba

rs,

a [m

ain]

=

10

–6,

a [H

2O

] =

1,

a [

SO

4--]

=

10–

5 (

spec

iate

s)

42

Evolution of Water Chemistry

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

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: Ca2+ and HCO3

- dominate

– Others: SO42- or NO3

- can dominate

• Also has dissolved gases: CO2 and O2 most important

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 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

46

Soil horizonsO 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:– 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)

48

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

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

50

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