Non-Conventional Energy Sources and Applications-ppt

ElectrochemistryElectrochemistry

Unit I

� Electrochemistry

� deals with interconversion between chemical and

electrical energy

� deals with the inter conversion between chemical

and electrical energy

� involves redox reactions� involves redox reactions

� 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

� Oxidation

� Reduction

� gain of electrons

� Reducing agent

� Oxidizing agent

� Oxidation

� Reduction

� Reducing agent

� donates the electrons and is oxidized

� Oxidizing agent

� Oxidation

� Reduction

� 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

Deposit of Cu

metal forms

CuSO4(aq)

(Cu2+)

Electrochemical CellsElectrochemical Cells

Cu cathodeZn anode

Zn2+Cu2+

Zn →→→→ Zn2+ + 2e- Cu2+ + 2e-→→→→ Cu

bridge

� Voltaic Cell

� cell in which a spontaneous redox reaction generates

electricity

� chemical energy → electrical energy

� 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

Ecell = Ecathode - Eanode = Eredn - Eox

E°cell = E°cathode - E°anode = E°redn - E°ox

(Ecathode and Eanode are reduction potentials by definition.)

� 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

ed. ag

ent stren

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

ed. ag

ent stren

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

� Determine if oxidizing and reducing agent react

spontaneously

� diagonal rule

ox. agent

red. agent

� Determine if oxidizing and reducing agent react

spontaneously

more +

Cathode

(reduction)

(oxidation)

more -

� 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

� Is there a relationship between Ecell and ∆G for a

redox reaction?

� Relationship between Ecell and ∆G:

� ∆G = -nFEcell

• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s

transferred redox rxn.

� ∆G = -nFEcell

transferred redox rxn.

• 1 J = CV• 1 J = CV

• ∆G < 0, Ecell > 0 = spontaneous

Equilibrium Constants from EEquilibrium Constants from Ecellcell

� ∆G = -nFEcell

transferred redox rxn

• 1 J = CV• 1 J = CV

� Under standard state conditions:

� ∆G° = -nFE°cell

� ∆G = -nFEcell

transferred redox rxn

• 1 J = CV• 1 J = CV

� ∆G = -nFEcell

• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn

• 1 J = CV

• ∆G < 0, Ecell > 0 = spontaneouscell

� ∆G° = -RTlnK

� -nFE°cell = -RTlnK

∆∆∆∆H° ∆∆∆∆S°

Calorimetric Data

∆∆∆∆G°Electrochemical

DataComposition

E°cell

Equilibrium

constants

The Nernst EquationThe Nernst Equation

� ∆G depends on concentrations

� ∆G = ∆G° + RTlnQ

� ∆G = -nFEcell and ∆G° = -nFE°cell� ∆G = -nFEcell and ∆G° = -nFE°cell

� -nFEcell = -nFE°cell + RTlnQ

� 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

� Ecell = E°cell - (0.0257/n)lnQ

� 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

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

anode (-):

oxidation

occurs here

e- flowion migration occurs

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

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

Half-reactions

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)

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-

cathode: PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- -->

PbSO4(s) + 2H2O(l)

2-(aq) --> PbSO4(s) + 2e-

PbSO4(s) + 2H2O(l)

overall: Pb(s) + PbO2(s) + 2H2SO4(aq) -->

2PbSO (s) + 2H O(l) 2PbSO4(s) + 2H2O(l)

2-(aq) --> PbSO4(s) + 2e-

PbSO4(s) + 2H2O(l)

overall: Pb(s) + PbO2(s) + 2H2SO4(aq) -->

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) -->

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

Half-reactions

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?

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.

H2OH2O

AnodeAnodeElectrolyteElectrolyte

CathodeCathode

•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

� 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-

6x9=54 (cells) 100-300 (modules)

Installation Installation exampleexample

Owner can sell

excess power to

power utility.

Roof top of residence ( Grid connected )

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