ElectrochemistryElectrochemistry
Unit I
ElectrochemistryElectrochemistry
� Electrochemistry
� deals with interconversion between chemical and
electrical energy
ElectrochemistryElectrochemistry
� Electrochemistry
� deals with the inter conversion between chemical
and electrical energy
� involves redox reactions� involves redox reactions
ElectrochemistryElectrochemistry
� Electrochemistry
� deals with inter conversion between chemical and
electrical energy
� involves redox reactions� involves redox reactions
• electron transfer reactions
Redox reactions (quick review) Redox reactions (quick review)
� Oxidation
� Reduction
� Reducing agent
� Oxidizing agent
Redox reactions (quick review)Redox reactions (quick review)
� Oxidation
� loss of electrons
� Reduction
� Reducing agent
� Oxidizing agent
Redox reactions (quick review)Redox reactions (quick review)
� Oxidation
� loss of electrons
� Reduction
� gain of electrons
� Reducing agent
� Oxidizing agent
Redox reactions (quick review)Redox reactions (quick review)
� Oxidation
� loss of electrons
� Reduction
� gain of electrons
� Reducing agent
� donates the electrons and is oxidized
� Oxidizing agent
Redox reactions (quick review)Redox reactions (quick review)
� Oxidation
� loss of electrons
� Reduction
� gain of electrons
� Reducing agent
� donates the electrons and is oxidized
� Oxidizing agent
� accepts electrons and is reduced
Redox ReactionsRedox Reactions
� Direct redox reaction
� Oxidizing and reducing agents are mixed together
� Indirect redox reaction
� Oxidizing and reducing agents are separated but
connected electrically
• Example
– Zn and Cu2+ can be reacted indirectly
� Basis for electrochemistry– Electrochemical cell
Redox ReactionsRedox Reactions
� Direct redox reaction
� Oxidizing and reducing agents are mixed together
Zn rod
CuSO4(aq)
(Cu2+)
Zn rod
Zn rod
Deposit of Cu
metal forms
CuSO4(aq)
(Cu2+)
Electrochemical CellsElectrochemical Cells
e- e-
Cu cathodeZn anode
Zn2+Cu2+
Zn →→→→ Zn2+ + 2e- Cu2+ + 2e-→→→→ Cu
Salt
bridge
Electrochemical CellsElectrochemical Cells
� Voltaic Cell
� cell in which a spontaneous redox reaction generates
electricity
� chemical energy → electrical energy
Electrochemical CellsElectrochemical Cells
� Electrolytic Cell
� electrochemical cell in which an electric current
drives a non-spontaneous redox reaction
� electrical energy → chemical energy� electrical energy → chemical energy
Cell PotentialCell Potential
� Cell Potential (electromotive force), Ecell (V)
� electrical potential difference between the two
electrodes or half-cells
• Depends on specific half-reactions, concentrations, and • Depends on specific half-reactions, concentrations, and
temperature
• Under standard state conditions ([solutes] = 1 M, Psolutes =
1 atm), emf = standard cell potential, E°cell
• 1 V = 1 J/C
� driving force of the redox reaction
Cell PotentialCell Potential
Ecell = Ecathode - Eanode = Eredn - Eox
E°cell = E°cathode - E°anode = E°redn - E°ox
(Ecathode and Eanode are reduction potentials by definition.)
Cell PotentialCell Potential
� E°cell = E°cathode - E°anode = E°redn - E°ox
� Ecell can be measured
• Absolute Ecathode and Eanode values cannot be found.
� Reference electrode� Reference electrode
� has arbitrarily assigned E
� used to measure relative Ecathode and Eanode for half-cell reactions
� Standard hydrogen electrode (S.H.E.)
� conventional reference electrode
Standard Hydrogen ElectrodeStandard Hydrogen Electrode
� E° = 0 V (by
definition; arbitrarily
selected)
2H+ + 2e- → H� 2H+ + 2e- → H2
Standard Electrode PotentialsStandard Electrode Potentials
� Standard Reduction Potentials, E°
� E°cell measured relative to S.H.E. (0 V)
• electrode of interest = cathode
� If E° < 0 V:� If E° < 0 V:
• Oxidizing agent is harder to reduce than H+
� If E° > 0 V:
• Oxidizing agent is easier to reduce than H+
Standard Reduction PotentialsStandard Reduction PotentialsReduction Half-Reaction E°°°°(V)
F2(g) + 2e- →→→→ 2F-(aq) 2.87
Au3+(aq) + 3e- →→→→Au(s) 1.50
Cl2(g) + 2 e- →→→→ 2Cl-(aq) 1.36
Cr2O72-(aq) + 14H+(aq) + 6e- →→→→ 2Cr3+(aq) + 7H2O 1.33
O2(g) + 4H+ + 4e- →→→→ 2H2O(l) 1.23
Ag+(aq) + e- →→→→Ag(s) 0.80
Fe3+(aq) + e- →→→→ Fe2+(aq) 0.77
Cu2+(aq) + 2e- →→→→ Cu(s) 0.34
Ox. ag
ent
stre
ngth
incr
ease
sR
ed. ag
ent stren
gth
increases
Cu2+(aq) + 2e- →→→→ Cu(s) 0.34
Sn4+(aq) + 2e- →→→→ Sn2+(aq) 0.15
2H+(aq) + 2e- →→→→ H2(g) 0.00
Sn2+(aq) + 2e- →→→→ Sn(s) -0.14
,i2+(aq) + 2e- →→→→ ,i(s) -0.23
Fe2+(aq) + 2e- →→→→ Fe(s) -0.44
Zn2+(aq) + 2e- →→→→ Zn(s) -0.76
Al3+(aq) + 3e- →→→→Al(s) -1.66
Mg2+(aq) + 2e- →→→→ Mg(s) -2.37
Li+(aq) + e- →→→→ Li(s) -3.04
Ox. ag
ent
stre
ngth
incr
ease
sR
ed. ag
ent stren
gth
increases
Uses of Standard Reduction Uses of Standard Reduction
PotentialsPotentials
� Compare strengths of reducing/oxidizing agents.
� the more - E°, stronger the red. agent
� the more + E°, stronger the ox. agent
Uses of Standard Reduction Uses of Standard Reduction
PotentialsPotentials
� Determine if oxidizing and reducing agent react
spontaneously
� diagonal rule
ox. agent
red. agent
Uses of Standard Reduction Uses of Standard Reduction
PotentialsPotentials
� Determine if oxidizing and reducing agent react
spontaneously
more +
Cathode
(reduction)
Anode
(oxidation)
more -
Uses of Standard Reduction Uses of Standard Reduction
PotentialsPotentials
� Calculate E°cell
� E°cell = E°cathode - E°anode
• Greater E°cell, greater the driving force
� E° > 0 : spontaneous redox reactions� E°cell > 0 : spontaneous redox reactions
� E°cell < 0 : nonspontaeous redox reactions
Cell PotentialCell Potential
� Is there a relationship between Ecell and ∆G for a
redox reaction?
Cell PotentialCell Potential
� Relationship between Ecell and ∆G:
� ∆G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn.
Cell PotentialCell Potential
� Relationship between Ecell and ∆G:
� ∆G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn.
• 1 J = CV• 1 J = CV
• ∆G < 0, Ecell > 0 = spontaneous
Equilibrium Constants from EEquilibrium Constants from Ecellcell
� Relationship between Ecell and ∆G:
� ∆G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn
• 1 J = CV• 1 J = CV
• ∆G < 0, Ecell > 0 = spontaneous
� Under standard state conditions:
� ∆G° = -nFE°cell
Equilibrium Constants from EEquilibrium Constants from Ecellcell
� Relationship between Ecell and ∆G:
� ∆G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn
• 1 J = CV• 1 J = CV
• ∆G < 0, Ecell > 0 = spontaneous
� Under standard state conditions:
� ∆G° = -nFE°cell
Equilibrium Constants from EEquilibrium Constants from Ecellcell
� Relationship between Ecell and ∆G:
� ∆G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
• 1 J = CV
• ∆G < 0, Ecell > 0 = spontaneouscell
� Under standard state conditions:
� ∆G° = -nFE°cell
and
� ∆G° = -RTlnK
so
� -nFE°cell = -RTlnK
∆∆∆∆H° ∆∆∆∆S°
Calorimetric Data
∆∆∆∆G°Electrochemical
DataComposition
Data
E°cell
Equilibrium
constants
K
The Nernst EquationThe Nernst Equation
� ∆G depends on concentrations
� ∆G = ∆G° + RTlnQ
and
� ∆G = -nFEcell and ∆G° = -nFE°cell� ∆G = -nFEcell and ∆G° = -nFE°cell
thus
� -nFEcell = -nFE°cell + RTlnQ
or
� Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
The Nernst EquationThe Nernst Equation
� Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
� At 298 K (25°C), RT/F = 0.0257 V
so
� Ecell = E°cell - (0.0257/n)lnQ
or
� Ecell = E°cell - (0.0592/n)logQ
Commercial Voltaic CellsCommercial Voltaic Cells
� Battery
� commercial voltaic cell used as portable source of
electrical energy
� types
� primary cell� primary cell
• Nonrechargeable
• Example: Alkaline battery
� secondary cell
• Rechargeable
• Example: Lead storage battery
How Does a Battery WorkHow Does a Battery Work
Seal/cap
Assume a generalized battery
cathode (+)
anode (-)
Electrolyte
Paste
BatteryBattery
Electrolyte paste:
ion migration occurs
Placing the battery into a flashlight,
etc., and turning the power on
completes the circuit and allows
electron flow to occur
cathode (+):
Reduction occurs
here
anode (-):
oxidation
occurs here
e- flowion migration occurs
here
How Does a Battery WorkHow Does a Battery Work
� Battery reaction when producing electricity (spontaneous):
Cathode: O1 + e- → R1
Anode: R2 → O2 + e-
Overall: O1 + R2 → R1 + O2Overall: O1 + R2 → R1 + O2
� Recharging a secondary cell
� Redox reaction must be reversed, i.e., current is reversed (nonspontaneous)
Recharge: O2 + R1 → R2 + O1
� Performed using electrical energy from an external power source
BatteriesBatteries
� Read the textbook to fill in the details on
specific batteries.
� Alkaline battery
� Lead storage battery� Lead storage battery
� Nicad battery
� Fuel cell
Alkaline Dry CellAlkaline Dry Cell
Alkaline Dry CellAlkaline Dry Cell
Plated steel (+)
Cathode:
Mixture of
MnO2 and C
Brass rod
Plated steel (-)
Anode:
Mixture of Zn
and KOH(aq)
MnO2 and C
(graphite)
Paper or fabric
Separator
Insulators
Alkaline Dry CellAlkaline Dry Cell
Half-reactions
Alkaline Dry CellAlkaline Dry Cell
Half-reactionsanode: Zn(s) + 2OH-(aq) --> ZnO(s) + H2O(l) + 2e-
Alkaline Dry CellAlkaline Dry Cell
Half-reactionsanode: Zn(s) + 2OH-(aq) --> ZnO(s) + H2O(l) + 2e-
cathode: 2MnO2(s) + H2O(l) + 2e- -->
Mn2O3(s) + 2OH-(aq)Mn2O3(s) + 2OH (aq)
Alkaline Dry CellAlkaline Dry Cell
Half-reactionsanode: Zn(s) + 2OH-(aq) --> ZnO(s) + H2O(l) + 2e-
cathode: 2MnO2(s) + H2O(l) + 2e- -->
Mn2O3(s) + 2OH-(aq)Mn2O3(s) + 2OH (aq)
overall: Zn(s) + 2MnO2(s) --> Mn2O3(s) + ZnO(s)
Ecell = 1.54 V
Lead Storage BatteryLead Storage Battery
(anode)
(cathode)
(anode)
6 x 2V = 12 V
Lead Storage BatteryLead Storage BatteryHalf-reactions
Lead Storage BatteryLead Storage BatteryHalf-reactionsanode: Pb(s) + SO4
2-(aq) --> PbSO4(s) + 2e-
Lead Storage BatteryLead Storage BatteryHalf-reactionsanode: Pb(s) + SO4
2-(aq) --> PbSO4(s) + 2e-
cathode: PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- -->
PbSO4(s) + 2H2O(l)
Lead Storage BatteryLead Storage BatteryHalf-reactionsanode: Pb(s) + SO4
2-(aq) --> PbSO4(s) + 2e-
cathode: PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- -->
PbSO4(s) + 2H2O(l)
overall: Pb(s) + PbO2(s) + 2H2SO4(aq) -->
2PbSO (s) + 2H O(l) 2PbSO4(s) + 2H2O(l)
Lead Storage BatteryLead Storage BatteryHalf-reactionsanode: Pb(s) + SO4
2-(aq) --> PbSO4(s) + 2e-
cathode: PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- -->
PbSO4(s) + 2H2O(l)
overall: Pb(s) + PbO2(s) + 2H2SO4(aq) -->
2PbSO (s) + 2H O(l) 2PbSO4(s) + 2H2O(l)
Cell reaction reversed during recharging.
2PbSO4(s) + 2H2O(l) --> Pb(s) + PbO2(s) + 2H2SO4(aq)
Lead Storage BatteryLead Storage BatteryHalf-reactionsanode: Pb(s) + HSO4
2-(aq) --> PbSO4(s) + H+ + 2e-
cathode: PbO2(s) + 3H+(aq) + HSO42-(aq) + 2e- -->
PbSO4(s) + 2H2O(l)
overall: Pb(s) + PbO2(s) + 2H+ + 2HSO4-(aq) -->
2PbSO (s) + 2H O(l) 2PbSO4(s) + 2H2O(l)
Cell reaction reversed during recharging.
Lead Storage BatteryLead Storage Battery
Half-reactions during recharging (nonspontaneous)
cathode: PbSO4(s) + H+ + 2e- --> Pb(s) + HSO42-(aq)
anode: PbSO4(s) + 2H2O(l) -->
PbO2(s) + 3H+(aq) + HSO42-(aq) + 2e-PbO2(s) + 3H (aq) + HSO4 (aq) + 2e
overall: 2PbSO4(s) + 2H2O(l) -->
PbO2(s) + Pb(s) + 2H+ + 2HSO4-(aq)
Cell converted into electrolytic cell via application of
external electrical energy.
Fuel CellsFuel Cells
� Voltaic-like cell that operates with continuous
supply of energetic reactants (fuel) to the
electrodes
� utilize combustion reactions� utilize combustion reactions
� do not store chemical energy
• Not self-contained since reactants must be supplied to the
electrodes
� Example: Hydrogen-Oxygen fuel cell
HydrogenHydrogen--Oxygen Fuel CellOxygen Fuel Cell
HydrogenHydrogen--Oxygen Fuel CellOxygen Fuel Cell
Half-reactions
HydrogenHydrogen--Oxygen Fuel CellOxygen Fuel Cell
Half-reactionsanode: 2H2(g) + 4OH-(aq) --> 4H2O(l) + 4e-
HydrogenHydrogen--Oxygen Fuel CellOxygen Fuel Cell
Half-reactionsanode: 2H2(g) + 4OH-(aq) --> 4H2O(l) + 4e-
cathode: O2(g) + 2H2O(l) + 4e- --> 4OH-(aq)
HydrogenHydrogen--Oxygen Fuel CellOxygen Fuel Cell
Half-reactionsanode: 2H2(g) + 4OH-(aq) --> 4H2O(l) + 4e-
cathode: O2(g) + 2H2O(l) + 4e- --> 4OH-(aq)
overall: 2H2(g) + O2(g) --> 2H2O(l)overall: 2H2(g) + O2(g) --> 2H2O(l)
What is a Fuel Cell, and how does it What is a Fuel Cell, and how does it
work?work?
H2
H2
O2
O2
H2OH2O
e-e- e-e-
•A fuel cell is an electrochemical device
that combines hydrogen fuel and oxygen
from air to produce electricity and water.
•In a Polymer Electrolyte Fuel Cell, Hydrogen
ions form at the anode, and
diffuse through the electrolyte and react
with oxygen at the cathode.
H2
H2
O2
O2
H2
H2
H2
H2
H2
H2
H2
H2
H2
H2
O2
O2
O2
O2
O2
O2
O2
O2
H+H+
H+H+
H2OH2O
H2OH2O
AnodeAnodeElectrolyteElectrolyte
CathodeCathode
O2
O2
•Anode: H2 → 2H + (aq) +2e-
•Cathode: ½ O2 + 2H + (aq) + 2e- → H2O (l)
Uses of Fuel CellsUses of Fuel Cells
� Transportation:� Phosphoric Acid Fuel Cell
• Anode: H2(g) → 2H + (aq) + 2e-
• Cathode: ½ O2 (g) + 2H+ (aq) + 2e- → H2O(l)
� Portable:� Portable:� Proton Exchange Membrane Fuel Cell
• Anode: H2(g) → 2H + (aq) + 2e-
• Cathode: Cathode: ½ O2 (g) + 2H+ (aq) + 2e- → H2O(l)
� Stationary:� Solid Oxide Fuel Cells
• Anode: H2(g) + O2→ H2O(g) + 2e-
• Cathode: ½ O2 (g) + 2e- → O2-
Fuel Cell usesFuel Cell uses
� Fuel Cell transit buses in Chicago
(Ballard Corp)� Anode: H2(g) → 2H + (aq) + 2e-
� Cathode: ½ O2 (g) + 2H+ (aq) + 2e- → H2O(l)
� Energy Research Corp.� Anode: H2(g) + 2CO3 → H2O(g) + CO2(g) + 2e-
� Cathode: ½ O2 (g) + CO2 + 2e- → 2CO3
Solar Energy Solar Energy –– A Bright Idea!A Bright Idea!
“I’d put my money on the sun and solar energy.
What a source of power! I hope we don’t have
to wait ‘til oil and coal run out before we
tackle that.”
- Thomas Edison- Thomas Edison
People have been harnessing
solar energy for a long time!
Solar collector for
heating water
A home in California in 1906
Sun Angles
Solar Electric Solar Electric
(Photovoltaic)(Photovoltaic)
� Photovoltaic (PV) systems convert light
energy directly into electricity.
� Commonly known as “solar cells.”
� The simplest systems power the small
calculators we use every day. More
complicated systems will provide a large
Solar Electric SystemsSolar Electric Systems
complicated systems will provide a large
portion of the electricity in the near
future.
� PV represents one of the most promising
means of maintaining our energy intensive
standard of living while not contributing to
global warming and pollution.
How Does it Work?How Does it Work?� Sunlight is composed of photons, or bundles of radiant
energy. When photons strike a PV cell, they may be
reflected or absorbed (transmitted through the cell).
Only the absorbed photons generate electricity. When
the photons are absorbed, the energy of the photons is
transferred to electrons in the atoms of the solar cell.transferred to electrons in the atoms of the solar cell.
How Does it Work?How Does it Work?• Solar cells are usually made of two thin pieces of silicon, the
substance that makes up sand and the second most common
substance on earth.
• One piece of silicon has a small amount of boron added to it,
which gives it a tendency to attract electrons. It is called the p-
layer because of its positive tendency.
• The other piece of silicon has a small amount of phosphorous • The other piece of silicon has a small amount of phosphorous
added to it, giving it an excess of free electrons. This is called the
n-layer because it has a tendency to give up negatively charged
electrons.
How Does it Work?How Does it Work?
Best Place For Solar Panels?Best Place For Solar Panels?
� South Facing roof, adequate
space
� No shading (time of year, � No shading (time of year,
future tree growth)
� Roof structure, condition
Large Scale PV Large Scale PV
Power PlantsPower Plants
Prescott AirportLocation: AZ
Operator: Arizona Public ServiceConfiguration: 1,450 kWp
SGS SolarLocation: AZ
Operator: Tucson Electric Power CoConfiguration: 3,200 kWp
Centralized WindCentralized Wind--Solar Hybrid Solar Hybrid
SystemSystem
� In hybrid energy
systems more than
a single source of
energy supplies the
electricity. electricity.
� Wind and Solar
compliment one
another
Various Various type of PV celltype of PV cell
� Hierarchy of PV
Array (10-50kW)
Volt Ampere Watt Size
Cell 0.5V 5-6A 2-3W about 10cm
Module 20-30V 5-6A 100-200W about 1m
Array 200-300V 50A-200A 10-50kW about 30m
Cell (2-3W)
Array (10-50kW)
Module, Panel (100-
200W)
6x9=54 (cells) 100-300 (modules)
Installation Installation exampleexample
Owner can sell
excess power to
power utility.
Roof top of residence ( Grid connected )
Most popular installation style in
Japan.
(Almost 85% PV in Japan )
� Roof top of school ,community-center building.(For education and emergency power)
84
Distant and independent power supply ( Off grid )
85
Relay station on top of
mountain
Advertising sign beside
highway
� Mountain lodge ( Off grid )
Inverter and
86
1.2kW system
Inverter and
controller