Electrochemical energy systemsBatteries and Fuel cells

 Unit V: Electrochemical energy systems: Basic concepts of electrochemistry and electrochemical energy systems. Conventional primary batteries: Dry cell. Advanced primary batteries: Lithium and alkaline primary batteries. Conventional secondary batteries: Lead-acid, nickel-cadmium secondary batteries. Advanced secondary batteries: Nickel-Metal hydride and lithium-ion secondary batteries. Fuel cells: Key issues – Hydrogen-oxygen fuel cells – new generation fuel cells – electric vehicle application – solid oxide fuel cells.

Electrochemistry• Electrochemistry - study of the relationships which exist

between the flow of electrons and chemical reactions• Types of electrochemical systems

▫ electrolytic - chemical reaction which occurs when electrical current is passed through solution

▫ voltaic/galvanic - spontaneous reactions able to generate a supply of electricity (e.g., batteries)

• Spontaneous redox reactions are coupled in such a way (i.e., an electrochemical cell) as to allow electrons to flow through an external circuit

• The electrochemical cell design: half-cells (2); salt bridge; potentiometer; electrodes; electrolyte solutions; conducting wire

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Electrochemistry• A complete redox reaction takes place in a galvanic cell• Overall reaction separated into half-reactions which take

place at the anode and cathode• Given the following reaction:

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)the two half-reactions are:oxidation half-reaction: Zn(s) Zn2+(aq) + 2e- (anode)reduction half-reaction: Cu2+(aq) + 2e- Cu(s) (cathode)

• Electrode reactions:▫ Anode: site of oxidation; electrons originate there; neg.

pole of cell (anions migrate toward)▫ Cathode: site of reduction; electrons consumed there;

pos. pole of cell (cations migrate toward)

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Two electrodes are connected by an external circuit 4

Daniel Cell Design

Cell Notation• Cell notation is used to describe structure of galvanic cell• For the Zn/Cu cell, the galvanic cell notation is:

Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)

= phase boundary = salt bridgeanode reaction: to the left of the salt bridgecathode reaction: to the right of the salt bridgeboth half-cell reactions in order of spontaneous reactionZinc solid reacts to form zinc(II) ion at the anodeCopper(II) ion reacts to form copper metal at the cathode

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Electromotive Force• Electrons are driven (“pushed”) through conducting

wire in the direction of anode cathode by cell force• Origin of cell force is maximum electric potential

difference between electrodes or electromotive force (Ecell) or cell potential

• Potential difference - difference in electrical potential (electrical pressure) between two electrodes; standard unit of cell potential difference is the Volt

• Electrical workelectrical work = charge moved X potential

differenceJ = C X V

wmax = G = -nFEcell

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Standard Cell Potentials• Standard potential of galvanic cell = sum of standard

half-cell potentials of ox at anode and red at cathodeEo

cell = Eoox(anode) + Eo

red(cathode)

Since -Eored(anode) = Eo

ox(anode)Eo

cell = Eored(cathode) - Eo

red(anode)

• For a spontaneous cell reaction, Eocell is positive (since

Gibbs free energy change must be < 0)• Method developed to estimate standard cell potentials

under standard conditions (1 atm, 1 M, 25 oC)• Standard cell potentials termed standard reduction

potentials according to above formula• SHE used to determine standard reduction potentials

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Electrode•Electrically conductive•Can be inert electrode or reactive•Charge transfer takes place at the interface• Ion intercalation electrode used in secondary Li-

ion batteries•Every electrode has its electrode potential,

depends on the material and the electrolyte which is contact

Electrode•Anode (positive terminal), Cathode (negative

terminal) in electrolytic cell•Anode ( negative ), Cathode ( positive ) in

Galvanic cell•Three electrode system for electrolytic cell

Electrolyte• Ionically conductive medium•Electronically insulating•Can solid, semisolid,(gel), liquid •Should have high ionic conductivities•The solvent should have high potential window,

specially organic electrolytes have high potential window ( in Li-ion batteries)

• In this potential window electrolytes are inert under the electrochemical conditions

Slide 12 of

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Ecell, ΔG, and Keq

•Cells do electrical work.▫Moving electric charge.

•Faraday constant, F = 96,485 C mol-1

elec = -nFE

ΔG = -nFE

ΔG° = -nFE°

Slide 13 of

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Spontaneous Change•ΔG < 0 for spontaneous change.•Therefore E°cell > 0 because ΔGcell = -

nFE°cell •E°cell > 0

▫Reaction proceeds spontaneously as written.

•E°cell = 0▫Reaction is at equilibrium.

•E°cell < 0▫Reaction proceeds in the reverse direction

spontaneously.

Relationship Between E°cell and Keq

ΔG° = -RT ln Keq = -nFE°cell

E°cell = nF

RTln Keq

Slide 15 of

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Ecell as a Function of ConcentrationΔG = ΔG° -RT ln Q

-nFEcell = -nFEcell° -RT ln Q

Ecell = Ecell° - ln QnF

RT

Convert to log10 and calculate constants

Ecell = Ecell° - log Qn

0.0592 VThe Nernst Equation:

Galvanic cellElectrolytic cell

HISTORY OF BATTERIES

1800 Voltaic pile: silver zinc1836 Daniell cell: copper zinc1859 rechargeable lead-acid cell1868 Leclanché: carbon zinc wet cell1888 Gassner: carbon zinc dry cell1899 Junger: nickel cadmium cell

1946 Neumann: sealed NiCd1960s Alkaline, rechargeable NiCd1970s Lithium, sealed lead acid1990 Nickel metal hydride (NiMH)1991 Lithium ion1992 Rechargeable alkaline1999 Lithium ion polymer

HISTORY OF BATTERIES

Battery :- A battery is a storage device used for the storage of chemical

energy and for the transformation of chemical energy into electrical energy

Battery consists of group of two or more electric cells connected together electrically in series.

Battery acts as a portable source of electrical energy.

Energy produced by an electrochemical cell is not suitable for commercial purposes since they use salt bridge which produce internal resistance which results in drop in the voltage. The drop in voltage is negligible only for a small interval of time during which it is being used. Batteries are of 3 types. Namely

• Primary Batteries (or) Primary Cells• Secondary Batteries (or) Secondary Cells• Reserve Batteries• Fuel Cells (or) Flow Batteries

Primary (Disposable) Batteries

Leclanché Cells (zinc carbon or dry cell) Alkaline Cells

Mercury Oxide Cells Zinc/MnO2 Cells Aluminum / Air Cells

Lithium Cells Liquid cathode lithium cells Solid cathode lithium cells Solid electrolyte lithium cells

Secondary (Rechargeable) Batteries

Lead–acid Cells Nickel/Cadmium Cells Nickel/Metal Hydride (NiMH) Cells Lithium Ion Cells

Daniel cell – primary batteries

AnodeZn Zn2+

CathodeCu2+ + 2e- Cu

Cell reactionZn + cu2+ + SO4- ZnSO4 + Cu

EMF – 1.1 V

The Leclanché (Dry) Cell

Slide 24 of

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(Zn/MnO2 + C) Leclanché cell

Cathode2MnO2 + 2H2O + 2e- 2MnO(OH) + 2OH-

AnodeZn Zn2+ + 2e-

Secondary reactionNH4Cl + 2OH- NH3 + 2Cl- + 2H2OZn2+ + NH3 + 2Cl- [Zn(NH3)2] Cl2

MnO2 + C cathode was dipped into 20% NH4Cl In dry cell - Electrolyte in form of paste is used

Advantages: 1) These cells have voltage ranging from 1.25v to 1.50v. 2) Primary cells are used in the torches, radios, transistors,

hearing aids, pacemakers, watches etc.3) Price is low.

Disadvantages: These cells does not have a long life, because the acidic

NH4Cl corrodes the container even when the cell is not in use.

Alkaline Dry Cell• Zn/MnO2 – with KOH as electrolyte • Anode : Zn + 2OH- ZnO + H2O +

2e-

• Cathode 2MnO2 + 2H2O + 2e- 2MnO(OH) +2OH-

• Alkaline cells have high output capacity and current carrying ability• Less variation in output capacity• Zn electrodes must be very pure to avoid hydrogen evolution reaction

The Silver-Zinc Cell: A Button Battery

Zn(s),ZnO(s)|KOH(sat’d)|Ag2O(s),Ag(s)

Zn(s) + Ag2O(s) → ZnO(s) + 2 Ag(s) Ecell = 1.8 V

Lithium primary batteries• Li/ Ethylene carbonate + Propylene carbonate+

Li+/MnO2 .

• Li/ Ethylene carbonate + Propylene carbonate+ Li+/organic sulphides

• Li/ SOCl2 , Li/ electrolyte/SO2 , Li/ Electrolyte/I2+ polyvinyl pyridine

AnodeLi Li+ + e-

CathodeMnO2 + xLi+ + xe- LixMnO2

Give high energy densityLi metal have high capacityPassive layer formation on anode

Lead-acid batteryElectrolyte – 20 % H2SO4

H2SO4 Concentration decreases with discharging and regained on charging

This can tested by specific gravity measurement of H2SO4

Cell voltage 1.88 – 2.15 V

PbO2 + Pb + H2SO4 2PbSO4

+ 2H2O

dischargingcharging

Basics-Cell Chemistry• At the positive plate: PbO2 + 4H+ + SO4

2- + 2e- PbSO4 + 2H2O

• At the negative plate: Pb + SO42- PbSO4 + 2e-

• Total Cell Reaction: PbO2 + Pb +2H2SO4 2PbSO4 +2H2O

Note: Active materials include lead dioxide, lead and sulfuric acid.

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There are four stages in the discharging−charging cycle:

•Fully Charged•Discharging•Fully Discharged•Charging

Positive plate covered with lead oxide(PbO2)

Negative plate covered with a sponge lead (Pb)

Electrolyte contains water (H2O) and a sulfuric acid (H2SO4)

FULLY CHARGED

Current flows in the cell from the negative to the positive plates.

Electrolyte separates into hydrogen (H2) and sulfate (SO4).

The free sulfate combines with the lead (both lead oxide and sponge lead)

and becomes lead sulfate (PbSO4). The free hydrogen and oxygen

combine to form more water, diluting the electrolyte.

DISCHARGING

Both plates are fully sulfated.

Electrolyte is diluted to mostly water.

DISCHARGED

Reverses the chemical reaction

that took place during discharging. Sulfate (SO4) leaves the positive

and negative plates and combineswith hydrogen (H2) to becomesulfuric acid (H2SO4).

Hydrogen bubbles form at thenegative plates; oxygen appears at the positive plates.

Free oxygen (O2) combines with lead (Pb) at the positive plate to become lead oxide (PbO2).

CHARGING

Ni-Cd batteryCd / CdO / KOH / NiO or Ni2O3 / Ni

Other systems with Fe/FeO,

These batteries comes in a discharged state

Cell voltage – 1.3 V

Nickel Cadmium (Ni-Cd)• 1.2V, 400 Cycles• Inexpensive − Simple charging• low energy density − Memory effect• high self discharge (20% month)• Toxic

Advantages and uses1. The Nickel-Cadmium cell has small size and high

rate charge/discharge capacity, which makes it very useful.

2. It has also very low internal resistance and wide temperature range (up to 70°C).

3. It produces a potential about 1.4 volt and has longer life than lead storage cell.

4. These cells are used in electronic calculators, electronic flash units, transistors etc.

5. Ni- Cd cells are widely used in medical instrumentation and in emergency lighting, toys etc.

Ni – Metal hydride batteries•Alloy electrodes - anodes•AB2 , A - Group IV metal (Ti), B –Group VIII (Ni) •AB5 , A - Lanthanum (La), B – Group VIII (Ni) •These alloys help in better M-hydrides formation•This have better

performance than Ni- Hydrogen batteries•Cell voltage = 1.35 V

Anode : MH+2OH- M+H2O + 2e-

Cathode : NiO(OH)+H2O+e- Ni(OH)2+OH-

Over all reaction : MH+ NiO(OH) M + Ni(OH)2

Chemistry:LaNi5, TiMn2, ZrMn2 (-), nickel hydroxide (+)Potassium hydroxide aqueous electrolyte Features:

• Higher energy density (40%) than NiCd• Nontoxic• Reduced life, discharge rate (0.2-0.5C)• More expensive (20%) than NiCd• reduced memory effect − Less-toxic

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D

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D

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Li-ion Battery

+-e-e-

e- e-

Li+ conducting electrolyte LiCoO2LixC6 Graphite

charge

dischargeLi+

Li+

• Electrode redox reactions on charge: Cathode oxidation : LiCoO2 Li1-xCoO2 + xLi+ + xe-

Anode reduction : xLi+ + xe- + C6 LiC6discharge is the opposite

Fuel cells

PEM – fuel cell

Parts of a Fuel Cell• Anode

▫ Negative post of the fuel cell. ▫ Conducts the electrons that are freed from the hydrogen

molecules so that they can be used in an external circuit. ▫ Etched channels disperse hydrogen gas over the surface of

catalyst.• Cathode

▫ Positive post of the fuel cell▫ Etched channels distribute oxygen to the surface of the

catalyst.▫ Conducts electrons back from the external circuit to the

catalyst▫ Recombine with the hydrogen ions and oxygen to form water.

• Electrolyte▫ Proton exchange membrane.▫ Specially treated material, only conducts positively charged

ions.▫ Membrane blocks electrons.

• Catalyst ▫ Special material that facilitates reaction of oxygen and

hydrogen▫ Usually platinum powder very thinly coated onto carbon paper

or cloth.▫ Rough & porous maximizes surface area exposed to hydrogen

or oxygen▫ The platinum-coated side of the catalyst faces the PEM.

Types of Fuel Cells

- Transportation applications

- Space application

- avoids the need of pure H2

- envisaged for stationary power plants

- high volumetric energy density

Anode Cathode

H2-O2 fuel cell(Alkaline Fuel Cell)

H2 O2

H+

Overall: H2 + ½ O2 H2O

½ O2 + 2H+ + 2e- H2OH2 2H+ + 2e-

Electrolyte

e-

H2-O2 fuel cellO2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)

2{H2(g) + 2 OH-(aq) → 2 H2O(l) + 2 e-}

2H2(g) + O2(g) → 2 H2O(l)

E°cell = E°O2/OH- - E°H2O/H2

= 0.401 V – (-0.828 V) = 1.229 V

Electrodes• Need of porous electrode substrate•Need of electro catalysts•Noble metal electro catalysts shows high

performance•Stability of electro catalysts under the

given experimental conditions•Good electrical conductivity of these

electrodes

Electrolyte in fuel cell•Alkaline electrolyte ( 85 % KOH)•Phosphoric acid electrolyte•Solid polymer electrolyte conducting

protons – proton exchange membranes ( Nafion)

•Recent studies on anion exchange membranes

Fuel cell electrode reactions

Direct methanol fuel cell

Solid oxide fuel cell

High temperature fuel cellOxide ion conductor as electrolyte

Solid oxide fuel cell• Electrolyte - fused mixture of

Yttrium dioxide + Zirconium dioxide• Cells operate 800 – 1000oC • Charge transfer by O2- ions• Electrode materials should have

good oxide ion diffusion coefficient• Cathode -LaMnO3 , Anode – Ni/ZrO2 • No need for noble metal catalysts,

no corrosion problem as in case of liquid electrolyte systems• H2 + CO (reformate gas), CH4 can be used at anode

Advantages/Disadvantages of Fuel Cells• Advantages

▫Water is the only discharge (pure H2)• Disadvantages

▫CO2 discharged with methanol reform▫Little more efficient than alternatives▫Technology currently expensive

Many design issues still in progress▫Hydrogen often created using “dirty” energy

(e.g., coal)▫Pure hydrogen is difficult to handle

Refilling stations, storage tanks, …

Other energy resources•Solar cells•Tidal power generation•Nuclear power•Geothermal energy

GENERAL APPLICATIONS•Emergency power - Lithium cells, water

activated batteries•Standby power - Lead acid•Medical implants , long life, low self

discharge, high reliability - Lithium primary, button and special cells

•Cordless equipment - NiCad, Lithium Ion•Hearing aids, watches, calculators, memory

back up, wireless peripherals: Button and coin cells, Zinc air, Silver oxide.

Distinction between Primary, Secondary & Fuel cellsPrimary Secondary Fuel cells

1) It only acts as galvanic or voltaic cell. i.e., produces electricity

1) It acts as galvanic or voltaic cell while discharging (produces electricity) and acts as electrolytic cell (consumes electricity)

1) It is a simple galvanic or voltaic cell. i.e., produces electricity

2) Cell reaction is not reversible.

2) Cell reaction is reversible.

2) Cell reaction is reversible.

3) Can’t be recharged. 3) Can be recharged 3) Energy can be withdrawn continuously

4) Can be used as long as the active materials are present

4) Can be used again and again by recharging.

4) Reactants should be replenished continuously. it does not store energy.

eg: Leclanche cell or Dry cell, Lithium cell.

eg: Lead storage battery, Ni-Cd battery, Lithium ion cell

eg: H2&O2 Fuel cellCH3OH &O2 Fuel cell

Uses: In Pace makers watches, Transistors, radios ect.

Uses: In electronic equipments, automobile equipments, digital cameras, laptops, flash light.

Uses: Great use in space vehicles due to its light weight (product of is source of fresh water for astronauts )

Reference :• Engineering Chemistry R.P.Mani,K.N.Mishra,B.RamaDevi,Cengage learning publications,New Delhi(2009).

• Engineering Chemistry by P.C.Jain & M.Jain, Dhanpatrai & Co., New Delhi (2005).

• Modern Aspects of Electrochemistry, J. O’M. Bockris and A. K. N. Reddy, Kluwer Academic, 2000.

• Electrochemistry, Prof. B. Viswanathan et al., S.Viswanathan Publishers, 2007

• Electrochemistry of Semiconductors, Adrian W. Bott, Current Separations 17 (1998) 87 – 91.

• Electrochemical capacitors, Brian E. Conway, http://electrochem.cwru.edu/ed/encycl