Transcript Lecture 3-15-11

Oxidation-Reduction (Redox)
Reactions
1
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
Standard Potentials
Written as reductions
Strong Reducing Agents
Half-Reaction
Li+(aq) + e- → Li(s)
K+(aq) + e- → K(s)
Ba2+(aq) + 2 e- → Ba(s)
Sr2+(aq) + 2 e- → Sr(s)
Ca2+(aq) + 2 e- → Ca(s)
Na+(aq) + e- → Na(s)
Mg2+(aq) + 2 e- → Mg(s)
Be2+(aq) + 2 e- → Be(s)
Al3+(aq) + 3 e- → Al(s)
Mn2+(aq) + 2 e- → Mn(s)
2 H2O + 2 e- → H2(g) + 2 OH-(aq)
Zn2+(aq) + 2 e- → Zn(s)
Cr3+(aq) + 3 e- → Cr(s)
Fe2+(aq) + 2 e- → Fe(s)
Cd2+(aq) + 2 e- → Cd(s)
PbSO4(s) + 2 e- → Pb(s) + SO42-(aq)
Co2+(aq) + 2 e- → Co(s)
Ni2+(aq) + 2 e- → Ni(s)
Sn2+(aq) + 2 e- → Sn(s)
Pb2+(aq) + 2 e- → Pb(s)
2 H+(aq) + 2 e- → H2(g)
E0 (V)
-3.05
-2.93
-2.90
-2.89
-2.87
-2.71
-2.37
-1.85
-1.66
-1.18
-0.83
-0.76
-0.74
-0.44
-0.40
-0.31
-0.28
-0.25
-0.14
-0.13
0
The greater the E°, the more
easily the substance reduced
Half-Reaction
E0 (V)
2 H+(aq) + 2 e- → H2(g)
0
4+
2+
Sn (aq) + 2 e → Sn (aq)
0.13
2+
+
Cu (aq) + e → Cu (aq)
0.13
2+
SO4 (aq) + 4 H (aq) + 2 e → SO2(g) + 2 H2O
0.20
AgCl(s) + e → Ag(s) + Cl (aq)
0.22
Cu2+(aq) + 2 e- → Cu(s)
0.34
O2(g) + 2 H2 + 4 e → 4 OH (aq)
0.40
I2(s) + 2 e → 2 I (aq)
0.53
MnO4 (aq) + 2 H2O + 3 e → MnO2(s) + 4 OH (aq)
0.59
+
O2(g) + 2 H (aq) + 2 e → H2O2(aq)
0.68
Fe3+(aq) + e- → Fe2+(aq)
0.77
+
Ag (aq) + e → Ag(s)
0.80
2+
Hg2 (aq) + 2 e → 2 Hg(l)
0.85
2+
2+
2 Hg (aq) + 2 e → Hg2 (aq)
0.92
+
NO3 (aq) + 4 H (aq) + 3 e → NO(g) + 2 H2O
0.96
Br2(l) + 2 e- → 2 Br-(aq)
1.07
+
O2(g) + 4 H (aq) + 4 e → 2 H2O
1.23
+
2+
MnO2(s) + 4 H (aq) + 2 e → Mn (aq) + 2 H2O
1.23
2+
3+
Cr2O7 (aq) + 14 H (aq) + 6 e → 2 Cr (aq) + 7 H2O
1.33
Cl2(g) + 2 e → 2 Cl (aq)
1.36
Au3+(aq) + 3 e- → Au(s)
1.50
+
2+
MnO4 (aq) + 8 H (aq) + 5 e → Mn (aq) + 4 H2O
1.51
4+
3+
Ce (aq) + e → Ce (aq)
1.61
+
2PbO2(s) + 4H (aq) + SO4 (aq) + 2e → PbSO4(s) + 2H2O
1.70
+
H2O2(aq) + 2 H (aq) + 2 e → 2 H2O
1.77
Co3+(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
Strong Oxidizing Agents
Redox Cell
Pt wire
electrode
Salt bridge
H2 gas (1 atm)
Fe2+
and
Fe3+
Fe3+ + e- ↔ Fe2+
←: Pt wire removes electrons from half cell A
→: Pt wire provides electrons to the solution
[H+] = 1
H+ + e- ↔ ½ H2(g)
4
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
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)
6
Eh as Master Variable
• From electrochemistry: GR = -nF Eh
• Re-write Nernst Equation:
– At 25°C
– Oxidized species on side where e- are
7
Calculating Eh: Example
• SO42- + Fe2+ + 8H+ + 8e-  FeS + 4H2O
8
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
9
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
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°
11
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.
12
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
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
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
15
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
16
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
17
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
18
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
19
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
20
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
21
Oxidizing environment,
but death to fish
22
Eh-pH Diagrams
• Positive Eh = oxidizing environments; tend
to function as electron acceptors
• Negative Eh = reducing environments;
tend to function as electron donors
23
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
0
[Fe2+] = 10-4
–.5
0
2
4
–6
[Fe2+] = 10-8
.5
Fe
++
Hematite
[Fe2+] = 10-6
25°C
6
8
pH
10
12
Walt Mon Feb 06 2006
++++
Fe 2(OH) 2 ,
3
4
2
Fe (OH) (5 ), Magnetite, Fe(OH) (ppd)
+
Diagram Fe , T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H2 O] = 1; Suppressed: FeO(c), Goethite, Fe(OH)3 (ppd), Fe(OH)3 , Fe+++, Fe(OH)- , Fe(OH)+, FeOH++,
4
2
++
Eh (volts)
1
Solubility increases with decreasing pH and Eh;
i.e., hematite dissolved under these conditions
14
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
–6
10-4
0
[Fe2+] = 10-6
10-8
.5
Fe
++
Magnetite
0
–.5
25°C
2
4
6
8
pH
10
12
Walt Mon Feb 06 2006
FeOH ,
++
++++
,
Fe (OH)
2
2
3
4
Fe (OH) (5 )
+
Diagram Fe , T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H2 O] = 1; Suppressed: FeO(c), Hematite, Goethite, Fe(OH)3 (ppd), Fe(OH)3 , Fe+++, Fe(OH)- , Fe(OH)+,
4
2
++
Eh (volts)
1
14
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
25°C
0
–6
0
2
.5
Fe
++
–.5
2
4
6
8
pH
10
12
Walt Mon Feb 06 2006
Fe 3(OH) 4(5 ), Magnetite, Fe(OH)2 (ppd), FeOH , Fe(OH)2 , Fe(OH)3
3
3,
-
+
++
++++
Fe(OH)4, Fe(OH)2 , FeOH , Fe2 (OH)2
+++
+
, T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H O] = 1; Suppressed: FeO(c), Goethite, Fe(OH) (ppd), Fe(OH)
Fe
+
+++
Diagram Fe
Eh (volts)
1
14
35
,
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
36
0
2
Hematite
Fe
0
–.5
Magnetite
25°C
4
6
8
pH
10
12
Diagram Fe , T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H2 O] = 1; Suppressed: FeO(c)
–6
++
++
Eh (volts)
1
Fe
+++
.5
14
Walt Mon Feb 06 2006
37
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
38
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
39
0
2
Hematite
–.5
25°C
4
6
8
pH
10
-
–2
+
Fe
+++
.5
Fe
0
Siderite
Magnetite
12
Diagram Fe , T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H2 O] = 1, a [HCO 3] = 10 ; Suppressed: FeO(c), FeCH3 COO , Fe(CH3 COO)3 , Fe(CH COO)+, FeCH COO++
3
2
3
–6
++
++
Eh (volts)
1
14
Walt Mon Feb 06 2006
40
++
0
2
–.5
25°C
4
6
8
pH
10
--
Hematite
0
Pyrite
Troilite
Magnetite
FeO(c)
12
–5
Diagram Fe , T = 25 °C , P = 1.013 bars, a [main] = 10 , a [H2 O] = 1, a [SO 4 ] = 10
–6
Fe
++
(speciates)
Eh (volts)
1
Fe
+++
.5
14
Walt Fri May 05 2006
41
Evolution of Water Chemistry
42
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
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
44
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
45
Soil horizons
O horizon: surface layer predominately
organic matter
A horizon: highly weathered, high
organic matter, Fe/Al leached; high N
B horizon: accumulated clay, Fe/Al
hydroxides, humus (stable organic
matter; gaseous diffusion and aqueous
transport between B and C
C horizon: altered parent material,
solute and gases exchange with
saturated zone; periodically saturated
when water table high
Saturated zone
Zone of Leaching
Zone of
Accumulation
Partly decomposed
and unaltered
bedrock
46
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)
47
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
48
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
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 H2PO41 SO42H2O
Micronutrients: B, Cu, Mn,
Mo, Zn, Cl• Na+ only major ion not
involved in biological
activity
50