Download Redox Reactions and Electrochemistry

Redox Reactions and
Electrochemistry
1
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Electrochemical processes are oxidation-reduction reactions
in which:
• 
the energy released by a spontaneous reaction is
converted to electricity or
• 
electrical energy is used to cause a nonspontaneous
reaction to occur
0
0
2Mg (s) + O2 (g)
2Mg
O2 + 4e-
2+ 2-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2-
Reduction half-reaction (gain e-)
2
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1.  Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2.  In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3.  The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
3
4.  The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
5.  Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the molecule
or ion.
HCO3−
Identify the oxidation numbers of
all the atoms in HCO3− ?
O = −2
H = +1
3x(−2) + 1 + ? = −1
C = +4
4
Balancing Redox Equations
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
1.  Write the unbalanced equation for the reaction ion ionic form.
Fe2+ + Cr2O72-
Fe3+ + Cr3+
2.  Separate the equation into two half-reactions.
+2
Oxidation:
Reduction:
+3
Fe2+
Fe3+
+6
Cr2O72-
+3
Cr3+
3.  Balance the atoms other than O and H in each half-reaction.
Cr2O72-
2Cr3+
5
Balancing Redox Equations
4.  Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
Fe2+
Fe3+ + 1e6e- + Cr2O722Cr3+
5.  For reactions in acid, add H+ to balance electronic charge and
H2O to balance O atoms and H atoms
6e- +14H+ + Cr2O722Cr3+
6e- + 14H+ + Cr2O72-
2Cr3+ + 7H2O
6.  If necessary, equalize the number of electrons in the two halfreactions by multiplying the half-reactions by appropriate
coefficients.
6Fe2+
6Fe3+ + 6e6
6e- + 14H+ + Cr2O722Cr3+ + 7H2O
Balancing Redox Equations
7.  Add the two half-reactions together and balance the final
equation by inspection. The number of electrons on both
sides must cancel.
Oxidation:
6Fe2+
Reduction: 6e- + 14H+ + Cr2O72-
6Fe3+ + 6e2Cr3+ + 7H2O
14H+ + Cr2O72- + 6Fe2+
6Fe3+ + 2Cr3+ + 7H2O
8.  Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6 x 2 = 24 = 6 x 3 + 2 x 3
9.  For reactions in basic solutions, add OH- to instead of H+ to
balance electronic charges.
10. Balance the reaction in the molecular form.
7
Galvanic Cells
anode
oxidation
cathode
reduction
spontaneous
redox reaction
8
Galvanic Cells
The difference in electrical
potential between the anode
and cathode is called:
•  cell voltage
•  electromotive force (emf)
•  cell potential
Zn (s) + Cu2+ (aq)
Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M and [Zn2+] = 1 M
Cell Diagram
phase boundary
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
salt bridge
cathode
9
Standard Reduction Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s)
Cathode (reduction): 2e- + 2H+ (1 M)
Zn (s) + 2H+ (1 M)
Zn2+ (1 M) + 2eH2 (1 atm)
Zn2+ + H2 (1 atm)
10
Standard Reduction Potentials
Standard reduction potential (E°) is the voltage associated
with a reduction reaction at an electrode when all solutes
are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M)
H2 (1 atm)
E° = 0 V
Standard hydrogen electrode (SHE)
11
Standard Reduction Potentials
0 = 0.76 V
Ecell
° )
Standard emf (Ecell
°
° = E°
Ecell
cathode - Eanode
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
° = E ° + - E ° 2+
Ecell
H /H 2
Zn /Zn
° 2+
0.76 V = 0 - EZn
/Zn
° 2+
EZn
/Zn = -0.76 V
Zn2+ (1 M) + 2e-
Zn
E° = -0.76 V
12
Standard Reduction Potentials
° = 0.34 V
Ecell
°
° = E°
Ecell
E
cathode
anode
° = E ° 2+
°
Ecell
Cu /Cu – EH +/H2
° 2+
0.34 = ECu
/Cu - 0
° 2+
ECu
/Cu = 0.34 V
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
Anode (oxidation):
H2 (1 atm)
Cathode (reduction): 2e- + Cu2+ (1 M)
H2 (1 atm) + Cu2+ (1 M)
2H+ (1 M) + 2eCu (s)
Cu (s) + 2H+ (1 M)
13
• 
E° is for the reaction as
written
• 
The more positive E° the
greater the tendency for the
substance to be reduced
• 
The half-cell reactions are
reversible
• 
The sign of E° changes
when the reaction is
reversed
• 
Changing the stoichiometric
coefficients of a half-cell
reaction does not change
the value of E°
14
What is the standard emf of an electrochemical cell made
of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr
electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e-
Cd (s) E° = -0.40 V
Cr3+ (aq) + 3e-
Cr (s)
Anode (oxidation):
Cd is the stronger oxidizer
E° = -0.74 V
Cr3+ (1 M) + 3e- x 2
Cr (s)
Cathode (reduction): 2e- + Cd2+ (1 M)
2Cr (s) + 3Cd2+ (1 M)
Cd will oxidize Cr
Cd (s)
x3
3Cd (s) + 2Cr3+ (1 M)
°
° = E°
Ecell
cathode - Eanode
° = -0.40 – (-0.74)
Ecell
° = 0.34 V
Ecell
15
The electrochemical cell
The electrochemical cell
Redox reactions can be used to generate electric current
Electrode processes
Electrode processes
The metallic electrode is dipped into a solution containing a salt of the metal.
Some atoms of the metal can leave the electrode and form the cation in
solution, leaving electrons in the metal. This form a double layer of opposite
charges to the electrode surface. The electrochemical potential of the metal and
its ion should be the same at the equilibrium.
M(s)  Mn+(aq) + neThere is the formation of an electric potential proportional at the ion
concentration in solution
+
+
+
+
+ + - +
+
- +
- +
- +
- +
- +
- +
+
Electrode potential
Nernst law
M+(aq) + e-  M(s)
GM = G°M + RTlnaM
GM+ = G°M+ + RTlnaM+
ΔG = G°M + RTlnaM - G°M+ + RTlnaM+
ΔG = ΔG° + RTln(aM/aM+)
ΔG = -nFE
-nFE = ΔG° + RTln(aM/aM+)
E = -ΔG°/nF + RT/nF ln(aM+/aM)
E = E° + RT/nF ln aM+
aM = 1
The Nernst equation
The potential of an electrode is expressed by the Nernst law:
2.3RT
[Ox]
E=E +
log
nF
[Re d]
Where Ox and Red are oxidized and reduced forms of
Red-Ox couple in equilibrium:
0
!
Oxn+ + ne-  Red0
R is the universal gas constant, T is the absolute temperature in Kelvins, n is a number of electrons
transferred in reaction, F is Faraday constant (~ 96500 C)
The electromotive force
•  It is useful to separate the overall redox reaction in two
separated processes: the oxidation and reduction semireaction.
•  In the electrochemical cell we have two electrodes and
we indicate as Cathode the electrode where the
reductions occur and Anode the electrode of oxidation
processes.
•  The electromotive force (EMF) of the cell is the electric
potential difference among the cathode and anode.
The Standard Hydrogen Electrode
We cannot know the absolute potential of a single electrode (it is not possible to
measure half reaction), so the E° = 0 V was assigned to the semi-reaction
2 H3O+ + 2e-  H2
Reference electrodes: Ag/AgCl and SCE
Metallic electrodes
3 main groups:
• 
First kind - wire of active metal
immersed in solution, contained the ions
of this metal (Cu, Zn, Co, Fe, etc)
Al, Cu, Sn, inox, brass and Fe electrodes
•  Second kind – wire of metal covered by precipitate of hardly soluble salt
or oxide: M/Mn+ (Ag/AgCl for instance)
•  Third kind - inert metallic electrodes (Pt, Au, etc)
Pt electrode
Electrochemical cell
Minimum 2 electrodes are
required for electrochemical
measurements. Dipped in
electrolyte solution these
electrodes constitute an
electrochemical cell.
INDICATOR (or WORKING) electrode is an electrode responding to a
target analyte
REFERENCE electrode has a stable well defined potential value,
independent on analyzed solution composition
Reference electrodes: Ag/AgCl and SCE
Spontaneity of Redox Reactions
ΔG = -nFEcell
ΔG°
=
°
-nFEcell
n = number of moles of electrons in reaction
J
F = 96,500
= 96,500 C/mol
V • mol
°
ΔG° = -RT ln K = -nFEcell
°
Ecell
(8.314 J/K•mol)(298 K)
RT
ln K =
ln K
=
nF
n (96,500 J/V•mol)
°
Ecell
=
°
Ecell
0.0257 V
ln K
n
0.0592 V
log K
=
n
29
Spontaneity of Redox Reactions
°
ΔG° = -RT ln K = -nFEcell
30
What is the equilibrium constant for the following reaction
at 25°C?
Fe2+ (aq) + 2Ag (s)
Fe (s) + 2Ag+ (aq)
°
Ecell
0.0257 V
ln K
=
n
Oxidation:
2Ag
2Ag+ + 2e-
Reduction: 2e- + Fe2+
Fe
n=2
° 2+
°
E° = EFe
–
E
/Fe
Ag +/Ag
E° = -0.44 – (0.80)
E° = -1.24 V
K=e
°
Ecell
xn
0.0257 V
=e
-1.24 V x 2
0.0257 V
K = 1.23 x 10-42
31
The Effect of Concentration on Cell Emf
ΔG = ΔG° + RT ln Q
ΔG = -nFE
ΔG° = -nFE °
-nFE = -nFE° + RT ln Q
Nernst equation
E = E° -
RT
ln Q
nF
At 298 K
E =E° -
0.0257 V
ln Q
n
E =E° -
0.0592 V
log Q
n
32
Will the following reaction occur spontaneously at 250C if
[Fe2+] = 0.60 M and [Cd2+] = 0.010 M?
Fe2+ (aq) + Cd (s)
Fe (s) + Cd2+ (aq)
Oxidation:
Cd
Cd2+ + 2e-
Reduction: 2e- + Fe2+
2Fe
n=2
° 2+
°
E° = EFe
/Fe – ECd 2+/Cd
E° = -0.44 – (-0.40)
E° = -0.04 V
0.0257 V
ln Q
n
0.010
0.0257 V
ln
E = -0.04 V 2
0.60
E = 0.013
E =E° -
E>0
Spontaneous
33
Concentration Cells
Galvanic cell from two half-cells composed of the same
material but differing in ion concentrations.
34
Electrolysis
•  Electrolysis is the process in which
electrical energy is used to cause a
nonspontaneous chemical reaction
to occur.
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35
Electrolysis
•  Previously our lectures on electrochemistry were
involved with voltaic cells i.e. cells with Ecell > 0
and ΔG < 0 that were spontaneous reactions.
•  Today we discuss electrochemical cells where
Ecell < 0 and ΔG > 0 that are non-spontaneous
reactions and require electricity for the reactions
to take place. We can take a voltaic cell and
reverse the electrodes to make an electrochemical
cell.
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36
Voltaic
Electrolytic
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Electrolytic conductors
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42
Fig. 21.18: Car battery, both voltaic and electrochemical
cell.
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43
Increase
oxidizing
power
Increase
reducing
power
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A standard electrolytic cell. A power source forces the
opposite reaction
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Electrolysis
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46
(a) A silver-plated teapot.
(b) Schematic of the electroplating of a spoon.
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47
Schematic of the
electroplating of a spoon.
AgNO3(aq)
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48
The electrolysis of water produces hydrogen gas at the
cathode (on the right) and oxygen gas at the anode
(on the left).
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Electrolysis of water
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Electrolysis of water
•  At the anode (oxidation):
•  2H2O(l) + 2e-  H2(g) + 2OH-(aq)
• 
• 
• 
• 
• 
E= -0.42V
At the cathode (reduction):
O2(g) + 4H+(aq) + 4e-  2H2O(l)
E= 0.82V
Overall reaction after multiplying anode reaction by 2,
2H2O(l)  2H2(g) + O2(g)
Eocell = -0.42 -0.82 = -1.24 V
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51
Electrolysis: Consider the electrolysis of a solution
that is 1.00 M in each of CuSO4(aq) and NaCl(aq)
• 
Oxidation possibilities follow:
•  Cl2(g) + 2e–  2Cl–(aq)
E° = +1.358 V
•  S2O82–(aq) + 2e–  2SO42–(aq)
E° = +2.010 V
•  O2(g) + 4H+(aq) + 4e–  2H2O E° = +1.229 V
• 
Reduction possibilities follow:
•  Na+(aq) + e–  Na(s)
•  Cu2+(aq) + 2e–  Cu(s)
•  2H2O + 2e–  H2(g) + 2OH–(aq)
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E° = –2.713 V
E° = +0.337 V
E° = -0.428 V
52
Electrolysis
•  We would choose the production of O2(g) and Cu(s).
•  But the voltage for producing O2(g) from solution is considerably
higher than the standard potential, because of the high activation
energy needed to form O2(g).
•  The voltage for this half cell seems to be closer to –1.5 V in reality.
•  The result then is the production of Cl2(g) and Cu(s).
anode, oxidation: Cl2(g) + 2e–  2Cl–(aq) E° = +1.358 V
•  cathode, reduction: Cu2+(aq) + 2e–  Cu(s) E° = +0.337 V
•  overall: CuCl2(aq)  Cu(s) + Cl2(g)
E = –1.021 V
•  We must apply a voltage of more than +1.021 V to cause this reaction
to occur.
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Stoichiometry of electrolysis: Relation between
amounts of charge and product
•  Faraday s law of electrolysis relates to the amount
of substance produced at each electrode is directly
proportional to the quantity of charge flowing
through the cell (half reaction).
•  Each balanced half-cell shows the relationship
between moles of electrons and the product.
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Application of Faraday s law
•  1. First balance the half-reactions to find number
of moles of electrons needed per mole of product.
•  2. Use Faraday constant (F = 9.65E4 C/mol e-) to
find corresponding charge.
•  3. Use the molar mass of substance to find the
charge needed for a given mass of product.
–  1 ampere = 1 coulomb/second or 1 A = 1 C/s
–  A x s = C
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Stoichiometry of Electrolysis
  How
much chemical change occurs with the
flow of a given current for a specified time?
•  current and time → quantity of charge →
•  moles of electrons → moles of analyte →
•  grams of analyte
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Fig. 21.20
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58
Doing work with
electricity.
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59
Electrolysis and Mass Changes
charge (C) = current (A) x time (s)
1 mol e- = 96,500 C
60
How much Ca will be produced in an electrolytic cell of
molten CaCl2 if a current of 0.452 A is passed through the
cell for 1.5 hours?
Anode:
Cathode:
2Cl- (l)
Ca2+ (l) + 2eCa2+ (l) + 2Cl- (l)
Cl2 (g) + 2eCa (s)
Ca (s) + Cl2 (g)
2 mole e- = 1 mole Ca
C
s 1 mol e- 1 mol Ca
mol Ca = 0.452
x 1.5 hr x 3600 x
x
s
hr 96,500 C 2 mol e= 0.0126 mol Ca
= 0.50 g Ca
61
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62
Batteries
Dry cell
Leclanché cell
Anode:
Cathode:
Zn (s)
2NH4+ (aq) + 2MnO2 (s) + 2e-
Zn (s) + 2NH4 (aq) + 2MnO2 (s)
Zn2+ (aq) + 2eMn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
63
Alkaline ba*ery •  Electrolyte is a concentrated solu5on of KOH •  The anode is inside the ba*ery as a powder paste •  MnO2 is pasted with graphite around the Zn anode and contacted with the external steel electrode Alkaline ba*ery •  Lower polariza5on •  Higher dura5on •  Lower self discharge Rechargeable Alkaline Ba*ery •  Rechargeable Alkaline Manganese (RAM) cell •  The interest is due to the higher A: Alkaline manganese 2 – 3 Ah Nickel / cadmium 0.5 – 1.0 Ah Nickel / metal hydride 1 – 1.5 Ah The number of charge-­‐discharge is lower than usual Ni/Cd or Ni/MH cells Button Batteries
High energy and stable discharge, ideal for long time operation
with low A
Mercury Battery
(Silver Oxide)
Anode:
Cathode:
Zn(Hg) + 2OH- (aq)
HgO (s) + H2O (l) + 2eZn(Hg) + HgO (s)
ZnO (s) + H2O (l) + 2eHg (l) + 2OH- (aq)
ZnO (s) + Hg (l)
67
Batteries
Lead storage
battery
Anode:
Cathode:
Pb (s) + SO2-4 (aq)
PbSO4 (s) + 2e-
PbO2 (s) + 4H+ (aq) + SO24 (aq) + 2e
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq)
PbSO4 (s) + 2H2O (l)
2PbSO4 (s) + 2H2O (l)
68
Acidic ba*eries •  Pb ba*eries have been first reported in 1859 •  Electrode reac5ons: Pb + H2SO4 PbSO4 + 2H+ + 2e-­‐ ( -­‐0.356 V) PbO2 + H2SO4 + 2H+ + 2e-­‐ PbSO4 + 2H2O (1.685 V) Total reac5on: Pb + PbO2 + H2SO4 2PbSO4 + 2H2O (2.041 V) Pb ba*eries •  Advantages: Low cost Well known Technology Good Ah •  Disadvantages: Low energy density Deposi5on of low soluble PbSO4 Ni-­‐Cd Cd + 2NiOOH + 4H2O  Cd(OH)2 + 2Ni(OH)2.H2O e.f.m. = 1.20 V High number of cycles, reliable, low maintenance Energy density not high Cd is toxic and costly Ni-­‐MH •  Alterna5ve to Ni-­‐Cd cells •  Developed ader Ni-­‐H2 cells, for military applica5ons •  Electrode reac5ons: H2 + 2OH-­‐ 2H2O + 2e 2NiOOH + 2H2O + 2e 2Ni(OH)2 + 2OH-­‐ e.f.m. = 1.2 – 1.3 V Metallic hydride is used as hydrogen source Li Ba*ery •  Lightest metal; •  High nega5ve standard poten5al but: •  Easy to oxidize •  Unstable and not compa5ble with water Li Ba*ery •  Need non aqueous solvents •  Advantages: •  High voltage ( >4V) •  Uniform T discharge •  Long shelf-­‐life •  Loss of capacity < 10% •  Wide range of working T •  Cathode MnO2 Li Ba*ery •  Long self discharge (up to 10 years) •  Working T around -­‐40 °C and 60 °C Batteries
Solid State Lithium Battery
76
Li-­‐ion Ba*eries Batteries
A fuel cell is an
electrochemical cell
that requires a
continuous supply of
reactants to keep
functioning
Anode:
Cathode:
2H2 (g) + 4OH- (aq)
O2 (g) + 2H2O (l) + 4e2H2 (g) + O2 (g)
4H2O (l) + 4e4OH- (aq)
2H2O (l)
78
History
• 
The fuel cells had been conceived in 1839
by the British scientist Mr. William Grove.
• 
Developed practical applications during
years 60 and 70, for NASA.
• 
The American astronauts consumed the
water produced for the electric generators
of its ships.
• 
These generators had constituted the first
operational use of fuel cells.
What they are…
•  Electrochemical cell that converts chemical energy into electric
energy;
•  It can have taxes of conversion in the order of 90%;
•  Cathode + anode + electrolyte + catalyst;
•  Ex.: Combustible H2 and oxidant O2
Anode – H2(g) → 2H+(aq) + 2eCathode – 1/2O2(g) + 2H+(aq) + 2e- → H2O(g)
•  It is important the selection of the electrolyte, and the dimensions
of this and the electrodes.
and its operating…
Types of Fuel Cells
• 
• 
• 
• 
• 
• 
Polymer Electrolyte Fuel Cell (PEMFC)
Alkaline Fuel Cell (AFC)
Phosphoric Acid Fuel Cell (PAFC)
Molten Carbonate Fuel Cell (MCFC)
Intermediate Temperature Solid Oxide Fuel Cell (ITSOFC)
Solid Oxide Fuel Cell (SOFC)
PEMFC
• 
Operating Temperature: 50-100ºC
• 
Appropriate for electric vehicles (Automobile
Industry)
• 
Anode – Platinum (0.4mg/Pt cm2)
H2(g) → 2H+ + 2e-
• 
Cathode – Platinum (0.4 mg/Pt cm2)
1/2O2(g) + 2H+ + 2e- → H2O(aq)
• 
Common electrolyte:
- Solid organic polymer poly- perfluorosulfonic
acid;
- Membrane of Nafion.
• 
System Output: < 1kW - 250kW
• 
Efficiency Electrical:
- 53-58% (transportation)
- 25-35% (stationary)
[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm
Polymer Electrolyte Fuel Cell (PEMFC)
Applications :
• 
Backup power
• 
Portable power
• 
Small distributed generation
• 
Transportation
Advantages :
• 
Solid electrolyte reduces corrosion & electrolyte management problems
• 
Low temperature
• 
Quick start-up
Disadvantages :
• 
Requires expensive catalysts
• 
High sensitivity to fuel impurities
• 
Low temperature waste heat
• 
Waste heat temperature not suitable for combined heat and power (CHP)
AFC
• 
Operating Temperature: 90-100ºC
• 
Anode – Zn
H2(g) + 2OH-(aq) → 2H2O + 2e-
• 
Cathode – MnO2
1/2O2(g) + H2O + 2e- → 2OH-(aq)
• 
Common electrolyte:
- Aqueous solution of potassium
hydroxide soaked in a matrix
• 
System Output: 10kW - 100kW
• 
Efficiency Electrical: 60%
[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm
Alkaline Fuel Cell (AFC)
Applications :
• 
• 
Military
Space
Advantages :
•
Cathode reaction faster in alkaline electrolyte, higher performance.
Disadvantages :
• 
Expensive removal of CO2 from fuel and air streams required (CO2 degrades
the electrolyte).
PAFC
• 
Operating Temperature: 150-200ºC
• 
Anode – Platinum (0.1 mg/Pt cm2)
H2(g) → 2H+ + 2e-
• 
Cathode – Platinum (0.5 mg/Pt cm2)
1/2O2(g) + 2H+ + 2e- → H2O(aq)
• 
Common electrolyte:
- Liquid phosphoric acid soaked in a matrix
• 
System Output: 50kW – 1MW (250kW module
typical)
• 
Efficiency Electrical: 32-38%
[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm
Phosphoric Acid Fuel Cell (PAFC)
Applications :
• 
Distributed generation
Advantages :
• 
Higher overall efficiency with CHP
• 
Increased tolerance to impurities in hydrogen
Disadvantages :
• 
Requires expensive platinum catalysts
• 
Low current and power
• 
Large size/weight
MCFC
• 
Operating Temperature: 600-700ºC
• 
Anode: Nickel
H2(g) + CO32- → H2O(g) + CO2(g) + 2e-
• 
Cathode: Nickel
• 
1/2O2(g)+CO2(g)+2e-→ CO32-
• 
Common electrolyte:
- Carbonate salt
• 
System Output: < 1kW – 1MW (250kW
module typical)
• 
Efficiency Electrical: 45-47%
[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm
Molten Carbonate Fuel Cell (MCFC)
Applications :
• 
• 
Electric utility
Large distributed generation
Advantages :
• 
• 
• 
• 
High efficiency
Fuel flexibility
Can use a variety of catalysts
Suitable for CHP
Disadvantages :
• 
High temperature speeds corrosion and breakdown of cell components
• 
Complex electrolyte management
• 
Slow start-up
TSOFC
• 
Operating Temperature: 800 -1000ºC
• 
• 
Anode: Co-ZrO2 or Ni-ZrO2 cermet
H2(g) + O2- → H2O(l) + 2e-
• 
• 
Cathode: Sr-doped LaMnO3
1/2O2(g) + 2e- → O2-
• 
Common electrolyte:
- Solid zirconium oxide to which a small
amount of Yttria is added
• 
System Output: 5kW – 3MW
• 
Efficiency Electrical: 35-43%
[2] http://www.treehugger.com/files/2007/06/biogas-powered_fuel_system.php
Solid Oxide Fuel Cell (SOFC)
Applications :
• 
• 
• 
Auxiliary power
Electric utility
Large distributed generation
Advantages :
• 
High efficiency
• 
Fuel flexibility
• 
Can use a variety of catalysts
• 
Solid electrolyte
reduces electrolyte
• 
• 
Suitable for CHP
Hybrid/GT cycle
management problems
Disadvantages :
• 
High temperature enhances corrosion and breakdown of cell components
• 
Slow start-up
• 
Brittleness of ceramic electrolyte with thermal cycling
ITSOFC
• 
Operating Temperature: 600-800ºC
• 
• 
Anode: Co-ZrO2 or Ni-ZrO2 cermet
H2(g) + O2- → H2O(l) + 2e-
• 
• 
Cathode: Sr-doped LaMnO3
1/2O2(g) + 2e- → O2-
• 
Lower temperatures ⇒ increase the internal
resistance of the cell
• 
Common electrolyte:
- Solid zirconium oxide to which a small
amount of Yttria is added
• 
System Output: 5kW – 3MW
• 
Efficiency Electrical: 35-43%
[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm
Applications
Corrosion Corrosion is a spontaneous and irreversible electrochemical process, which results in the degrada5on of a metallic material, upon interac5on with the environment. The corrosion could occur in the presence or in the absence of water: The first one is called wet corrosion, the second dry corrosion As for all the chemical processes, the corrosion depends on both thermodynamic (spontaneous or not process) and kine5c (rate of the process) factors The interac5on with the environment could lead to: 1. The corrosion of the metal (ac5ve condi5on): the process is both thermodynamic and kine5c favored. ΔE > 0 2. The forma5on of a protec5ve film (passive condi5on): the process is favored by thermodynamic but kine5cally inhibited 3. No modifica5on of the metal: : the process is not thermodynamic favored. ΔE < 0 The corrosion is an electrochemical process, where a cathode and an anode are formed The metal is oxidized in the anodic region and leaves the electrons that migrate to the cathodic region, the corrosive region, where molecular oxygen is reduced Anodic process: Me  Men+ + ne-­‐ Cathodic process: O2 + 2H2O + 4e-­‐  4OH-­‐ or O2 + 4H+ + 4e-­‐  2H2O The molecular oxygen is more concentrated at the surface than in the bulk of the droplet, leading to a concentra5on cell. The oxygen reduc5on produces the hydroxide ions that lead to the rust forma5on This effect produces the ring morphology for the metal corrosion The corrosion can be: Generalized: the anodic zone is big, while the cathodic zone is small Localized: is the reverse case of the generalized corrosion. It is the most dangerous Generalized corrosion This corrosion interests all the metallic surface and leads to a reduc5on of the metal thickness Uniform Not uniform Localized corrosion This corrosion interests only small parts of the metal surface and it is the most dangerous because it is impossible to evaluate the gravity of the corrosive a*ack from an external inspec5on. Temporal evolu5on of the corrosion Constant process: ex. Fe in HCl Self-­‐cataly5c process: the hydrolysis of iron in the presence of Cl-­‐ ion produces protons in the anodic zone that increases the corrosion rate Self-­‐inhibi5ng process: the forma5on of carbonate salts in the alkaline region can produce low soluble salts that par5ally protect the metal surface from the oxygen reduc5on Passiva5ng process: the forma5on in the anodic zone of a compact oxide film that protect the metal: ex. Al Galvanic corrosion The corrosion is produced by a junc5on of two metals having different E: the metal with lower E is oxidized. Lower the ra5o of the zone anode/cathode, higher and more penetra5ng the dissolu5on of the less noble metal. Example of corrosion: the brass The brass is an alloy of copper and zinc: the zinc is oxidized and copper forms the characteris5c colored powder Protec5on methods Cathodic protec5on: cathodic current or sacrificial anode Appica5on of films resistant to corrosion: metallic, non-­‐metallic, polymers Cathodic Protection of an Iron Storage Tank
107
Protec5on with metallic films Hot deposi5on: immersion or spray coa5ng Galvanic deposi5on: electrochemical deposi5on (problems: not homogeneous thickness) Chemical deposi5on: deposi5on of the film by redox reac5on Protec5on with non-­‐metallic films Conver5on layers: the film is formed in situ by forma5on of chemical bond with the metal surface Ex.: chromature Protec5on with organic layers Thick films: gums or polymers Thin films: paints