UNIT-II

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ELECTROCHEMISTRY AND BATTERIES

Concept of electrochemistry: The branch of science which deals with the relationship between electrical

energy and chemical energy and their inter-conversion of one form to another is called electrochemistry.

Electrolysis: The changes in which electrical energy causes chemical reaction to occur.

Electro chemical cells: The changes in which electrical energy is produced as a result of chemical

change.

Conductors: The substances which allow the passage of electric current are called conductors.

Eg. Metals like Cu, Ag, Sn etc.

Insulators: The substances which do not allow the electric current to pass through them are called

non-conductors or insulators. Eg. Rubber, wood, wax, wool, glass etc.

Metallic conductors: These are the metallic substances which allow the electric current to pass through

them without undergoing any chemical change. The flow of electric current in metallic conductors is due

the flow of electrons in the metal atoms. Eg. Metals like Cu, Ag, Sn etc.

Electrolytes: These are the substances which allow the electric current to pass through them in their

molten states or in the form of their aqueous solutions and undergo chemical decomposition. The flow of

electric current through an electrolytic solution is called electrolytic conduction in which charge is carried

by ions. These substances do not conduct electricity in the solid state but conduct electricity in the molten

state or aqueous solutions due to movement of ions. Eg. Acids, bases and salts.

Non-electrolytes: The substances which do not conduct electricity either in their molten state or through

their aqueous solutions are called non-electrolytes. Eg. Sugar, glucose, urea, ethyl alcohol etc.

Strong electrolytes: The electrolytes which are almost completely dissociated into ions in are called

strong electrolytes. Eg: NaCl, KCl, HCl, NaOH, NH4NO3

Weak electrolytes: The electrolytes which do not ionize completely in solution are called weak

electrolytes. Equilibrium is established between unionized electrolyte and the ions formed in the

solution.

Types of Electrodes

Electrode potential (or) Single Electrode potential (E): It is the measure of tendency of a metallic

electrode to lose or gain electrons, when it is in contact with a solution of its own salt.

Standard Electrode potential (E°): It is the measure of tendency of a metallic electrode to lose or gain

electrons, when it is in contact with a solution of its own salt of one molar concentration at 25°C.

Types of Electrodes: An electrochemical cell consists of two electrodes, positive and negative. Each

electrode constitutes a half cell or a single electrode.

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CONSTRUCTION AND FUNCTIONING OF SOME STANDARD ELECTRODES:

The single electrode potential is conveniently measured by combining the half cell with a standard

electrode and measuring the total EMF of the cell.

Ecell = Eright - Eleft

Ecell = E right - Eo (if standard electrode is anode)

Ecell = Eo - Eleft (if standard electrode is cathode)

Where Eo is the standard electrode potential.

Eg. Standard hydrogen electrode, Calomel electrode, Quinhydrone electrode

1. Standard Hydrogen Electrode (S.H.E)

Standard Hydrogen Electrode (S.H.E) is a primary reference electrode.

Construction: S.H.E consists of a platinum electrode coated with Pt-black immersed in 1 Molar

solution of H+ ions maintained at 25oC. Hydrogen gas at 1atm pressure is constantly bubbled into the

glass hood over the Pt-electrode. Hydrogen gets oxidized at the Pt-electrode and passes into the

solution forming H+ ions and electrons. The standard electrode potential of Hydrogen electrode at

these conditions is considered as zero.

Depending upon the half cell to which it is combined, hydrogen electrode can act as either anode

(or) cathode.

When S.H.E acts as anode:

H2(g) 2H+ + 2e-

When S.H.E acts as cathode:

2H+ + 2e- H2(g)

Therefore S.H.E is regarded as reversible electrode.

H2(g) 2H+ + 2e-

Applications of S.H.E:

Determination of pH for the unknown solution:

For a Hydrogen electrode at 298K, the Nernst equation is given as follows:

2H+ (aq) + 2e- H2(g)

H+ (aq) + e- H2(g)

1/2o 2

Cell +

1/2 o

Cell 2+

+

Cell

Cell

[H ]2.303RTE = E - log

nF [H ]

0.0592 1E = - log ( [H ] = 1 and E = 0)

nF [H ]

E = - 0.0592 (-log[H ])

E = - 0.0592 pH

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In order to determine the pH of unknown solution, a Pt-electrode is immersed into the solution and

connected to S.H.E.

Cell representation: Pt, H2(g)│ H+ (aq) (?) ║ H+(aq) (1M)│ H2(g) (1atm), Pt

2. Calomel Electrode

Calomel Electrode is a secondary reference electrode.

Construction: It is a mercury-mercurous chloride (Hg-Hg2Cl2) electrode. The calomel electrode is

used as only reducing electrode i.e. as cathode only.

It consists of a glass tube having a side tube on each side. One side tube acts as salt bridge and other

is used to fill KCl solution. The high purity mercury is placed at the tip of this tube and connected to

the circuit by the means of Pt-wire, sealed in a glass tube. The surface of Hg is covered with a paste of

mercurous chloride (calomel) and Hg in KCl solution. The electrolyte is the solution of KCl. The

electrode is connected with the help of side tube on the left through salt bridge with the other electrode

whose potential has to be determined. The potential of the calomel electrode depends upon the

concentration of KCl.

Representation of calomel electrode:

Hg, Hg2Cl2(s) │ KCl (saturated solution)

Concentration of KCl Electrode potential (V)

0.1N 0.3335

1.0N 0.2810

Saturated 0.2422

+2

Cell cathode anode

Cell S.H.E H /H

Cell

Cell

Cell

E = E - E

E = E - E

E = 0- (-0.0592pH)

E = 0.0592pH

EpH =

0.0592

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Determination of pH of unknown solution by Saturated Calomel electrode:

Cell representation: Pt, H2 (1atm) │ H+ (?) ║ KCl (saturated) │ Hg2Cl2, Hg

+2Cell Calomel H /H

Cell

Cell

Cell

E = E - E

E = 0.2422 - (-0.0592pH)

E = 0.2422 + 0.0592pH

E 0.2422pH =

0.0592

3. Ion Selective Electrode- Glass electrode

They possess an ability to respond to only certain specific ions thereby developing a potential with

respect to that species only in a given mixture and ignore all the other ions totally. The potential

developed by ion selective electrode depends on the concentration of specific ion of interest.

Eg: Glass electrode specific for H+ ions and pressed pallet of Ag2S + AgCl specific for Cl- ions.

Glass electrode:

Construction: When two solutions of different pH values are separated by a thin glass membrane, a

potential difference develops between the two surfaces of the membrane. This potential difference

developed is proportional to the difference in pH value. The glass membrane functions as ion-

exchange resin.An equilibrium is established between Na+ ions of the glass and H+ ions in the solution.

The potential of the glass electrode is given by:

EG = EoG + 0.0592 pH

Where pH range of the test solution is 1-10.

The glass electrode consists of a thin walled glass tube containing AgCl

coated Ag-electrode or simply Pt-electrode in 0.1M HCl.

Cell representation of glass electrode is given by:

Ag │ AgCl(s) H+ (0.1M) │ Glass or Pt, HCl (0.1M)│ Glass

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Glass electrode & determination of pH by glass electrode:

HCl in the bulb furnishes constant H+ ion

concentration. Thus it is Ag-AgCl electrode

reversible with respect to Cl- ions. It is used as

internal reference electrode for determining the pH of

the solution especially the colored solution

containing oxidizing or reducing agents.

Cell representation: Pt, HCl (0.1M)│ Glass ║KCl (sat) Hg2Cl2(s) │Hg

o o

Cell Calomel Glass

o

Cell Glass

o

Cell GlassCell

E = E - E

E = 0.2422 - [E + 0.0592pH]

0.2422 - E - EE =

0.0592

Since the resistance is very high, a special electron tube voltmeter is used to measure the EMF of the

above cell.

Advantages:

1. It is simple and can be easily used.

2. The results are accurate.

3. Equilibrium is rapidly reached.

Limitations:

1. The range of solution pH is 1 to 10.

2. The resistance of the membrane is very high.

Nernst Equation

Standard electrode potentials are measured at standard states i.e at 1M concentration of

electrolyte, temperature at 298 K and 1atm pressure. However the electrode potentials depend upon

concentration of electrolyte solutions and temperature. Nernst equation gives the relationship between

electrode potentials and concentration of electrolytic solutions.

For the general reduction reaction occurring at an electrode,

Mn+(aq) + ne- M(s)

(s)on+ n+

M /M M /M n+

(aq)

[M ]2.303RTE = E - log

nF [M ]

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Where E° = Standard EMF of the cell for 1 M solution at 298 K;

R = Gas constant;

T = Kelvin temperature;

n = number of electrons involved in the cell reaction;

F = Faraday of electricity;

E = electrode potential of the metal;

[M] = activity of metal in the metal phase and is taken as unity;

[Mn+] = activity of metal ions in the solution is taken equal to their molarities;

But at STP conditions, R = 8.314JK-1mol-1; T = 298K; F = 96500 coulomb charge. Then the Nernst

equation becomes

o n+n+n+

M /M M /M (s)n+

(aq)

on+ n+

M /M M /M n+

(aq)

o n+n+ n+

M /M M /M

2.303 x 8.314 x 298 1E = E - log considering [M ] = 1

n x 96500 [M ]

0.0592 1 E = E - log

n [M ]

0.0592 E = E + log[M ]

n

Nernst equation enables us to calculate the following:

1. Half cell potential or single electrode potential.

2. Cell potential or EMF of the cell.

3. Equilibrium constant for the cell reaction.

Electrochemical series and its applications

Definition: When the various electrodes (metals) are arranged in the order of their increasing values of

standard reduction potential on the hydrogen scale, then the arrangement is called electrochemical

series.

(or)

When the various electrodes (metals) are arranged in the order of their decreasing values of standard

oxidation potential on the hydrogen scale, then the arrangement is called electrochemical series.

Applications of electrochemical series:

1. Calculation of standard EMF of the cell.

2. Relative ease of oxidation (or) Reduction.

3. Displacement of one element by the other.

4. Hydrogen displacement behavior.

5. Corrosion tendency of the elements.

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Electro chemical cell

An electrochemical cell is a device which converts chemical energy into electrical energy. The redox

reaction is utilized for generation of electrical energy.

The electrochemical cells are commonly referred as Voltaic or Galvanic cells. The electrochemical cell

is divided into two half cells. The half cell electrode where oxidation (loss of electrons) occurs is called

anode (negative electrode). The half cell electrode where reduction (gain of electrons) occurs is called

cathode (positive electrode). An electrochemical cell is the coupling of these two half cells.

Construction of Daniel cell / voltaic cell

The Daniel cell is a typical example of Galvanic cell. Daniel cell consists of a beaker containing copper

rod dipped in CuSO4 solution which is connected to another beaker containing zinc rod dipped in

ZnSO4 solution by a salt bridge. Salt bridge is an inverted U-tube containing saturated solution of some

electrolyte such as KCl, KNO3, and NH4NO3 which does not undergo chemical change during the

process. The saturated solution is generally taken in agar-agar jelly or gelatin.

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The salt bridge allows the flow of ions to pass through it when the flow of electric current takes place.

It completes the electrical circuit and maintains the electrical neutrality of two half cell solutions. When

the circuit is completed, electric current flows through the external circuit as indicated by an ammeter.

The following observations are made:

1. Zinc rod gradually loses its weight.

2. The concentration of Zn+2(aq) in the ZnSO4 (aq) solution increases.

3. Copper gets deposited on the electrode.

4. The concentration of Cu+2(aq) in CuSO4 (aq) solution decreases.

5. The flow of electrons is from Zn-electrode (anode) to Cu-electrode (cathode).

6. The flow of electric current is from Cu-electrode (cathode) to Zn-electrode (anode).

The above observations can be explained as follows:

Zn is oxidized to Zn+2 ions which go in the solution and therefore Zn rod gradually loses its weight

(oxidation half cell).

Zn Zn+2 + 2e- (At anode)

The electrons released at the Zn electrode move towards Cu electrode through external circuit. These

electrons are accepted by the Cu+2 ions in the solution and get reduced to copper which gets deposited

on the Cu-electrode (reduction half cell).

Cu+2 + 2e- Cu (At cathode)

The oxidation of Zn occurs at anode (negative terminal) and reduction of Cu+2 occurs at cathode

(positive terminal). The flow of electrons is from negative terminal (anode) to positive terminal

(cathode).

Flow of electrons

(-) Anode Zn │ ZnSO4 ║ CuSO4 │ Cu Cathode (+)

Flow of electric current

Cell reactions: At Anode: Zn Zn+2 + 2e-

At cathode: Cu+2 + 2e- Cu

Overall Reaction: Zn + Cu+2 Zn+2 + Cu

Cell representation: Zn(s) │ ZnSO4 (aq) (1M) ║ CuSO4 (aq) (1M)│ Cu(s)

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Electro Motive Force

The flow of electric current in Galvanic cell is due to potential difference between two electrodes which

is known as electrode potential. Electrode potential is a measure of tendency of an electrode to gain or

lose electrons. Each electrode has a different tendency to lose or gain electrons.

Reduction potential (conventionally used) is the measure of tendency of an electrode to lose electrons

and oxidation potential is the measure of tendency of an electrode to lose electrons. Electrodes with

high reduction potential readily gains electrons and gets reduced. Electrodes with low reduction

potential (oxidation potential) readily loose electrons and gets oxidized. The difference between the

electrode potentials of the two electrodes constituting a galvanic cell is known as EMF or cell

potential (expressed in volts).

EMF = Ecell = Eocathode – Eo

anode = Eoright - E

oleft

The reduction potential of an electrode measured at 25oC, 1atm and 1M concentration of electrolyte

is called standard reduction potential. The positive value of Ecell indicates that cell reaction is

feasible.

Concentration cells

Definition: The electrical energy in a concentration cell is due to 'transfer of substance' from the

solution of high concentration to the solution of low concentration.

The concentration cell is made up of two half cells having identical electrodes and identical

electrolytes. However the concentrations of the reactive ions at the two electrodes are different.

Cell representation: M │ Mn+(aq) (C1) ║ Mn+

(aq) (C2)│ M

Where C1 and C2 are the concentrations of the active metal ions in contact with the two electrodes

respectively.

Eg. Silver concentration cell

When a metal electrode is dipped in solution containing Mn+ ions, then a potential E is developed at the

electrode. The value of E varies with the concentration (C) of the ions in accordance with Nernst

equation.

o

Cell

2.303RTE = E + logC

nF

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Electrolytic concentration cell

These cells consist of two identical electrodes immersed in two solutions of the same electrolyte at

different concentrations. The electrodes of such cells are reversible to one of the ions of the electrolyte.

Eg. Zinc concentration cell: Zn │ Zn+2(aq) (C1) ║ Zn+2

(aq) (C2)│ Zn

where C1 and C2 are the concentration of Zn+2 at anode and cathode respectively (C2 ≥ C1).

Thus EMF so developed is due to the transference of metal ions from the solution of higher

concentration (C2) to the solution of lower concentration (C1)

n+ n+

Cell Cathode Anode

o o

Cell 2 1M /M M /M

Cell 2 1

Cell 2 1

2Cell

1

Cell

E = E - E

2.303RT 2.303RTE = E + logC - [E + logC ]

nF nF

2.303RT 2.303RTE = logC - logC

nF nF

2.303RTE = [logC - logC ]

nF

C2.303RTE = log

nF C

0.0E =

2

1

C592 log (At 298K)

n C

If C1 = C2 , then E cell = 0.

Applications of concentration cells:

1. To determine the solubility of a sparingly soluble salts.

2. To calculate the valency of cations.

3. To calculate the extent of corrosion in metals.

4. To determine the transition point.

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Batteries

The term battery is generally used for two or more galvanic cells connected in series. Thus a battery is a

series of portable electrochemical cells which are capable of generating electrical energy.

Batteries are of three types: (1) Primary cell (2) Secondary cell (3) Fuel cell.

Applications of batteries:

The portability of electronic equipments in the form of handsets has been made possible by

batteries. A variety of electronic gadgets with more reliability and service have been made more useful

with the introduction of rechargeable storage batteries.

Comparison between primary, secondary and fuel cells

Sl.

No.

Primary cells Secondary cells Fuel cells

1. Cell reaction is irreversible

Cell reaction is reversible.

Energy can be withdrawn

indefinitely as long as

outside supply of fuel is

maintained.

2. Must be discarded after use. May be recharged. Cannot be reused.

3. Have relatively short shelf

life.

Have long shelf life. Have life as long as the fuel

is supplied.

4. Function only as galvanic

cells.

Functions both galvanic Cell

& as electrolytic cell.

Functions as only galvanic

cell.

5. They cannot be used as

storage devices

They can be used as energy

storage devices (e.g. solar/

thermal energy converted to

electrical energy)

Do not store energy.

6. They cannot be recharged They can be recharged. They cannot be recharged.

7. Eg. Dry cell, Alkaline cell

and Li- battery.

Eg. Lead acid storage cell,

Ni-Cd battery.

Eg: H2-O2 fuel cell and

CH3OH-O2 fuel cell

Primary Batteries

In primary cells the cell reaction is not reversible. No electricity is produced after complete conversion

of the reactants to products and the cell becomes dead. These batteries are used as source of DC power.

Eg: 1. Dry cell (Leclanche cell),

2. Alkaline cell

3. Lithium cell

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1. Dry Cell (Leclanchi Cell) - Primary Battery

Anode: Zn-container

Cathode: Graphite rod

Electrolyte: Mixture of MnO2(s), NH4Cl (aq) and ZnCl2(s) to which starch is added to make a paste.

Voltage: 1.5 V

Construction:

The anode of the cell is a Zn-container containing

an electrolytic mixture of NH4Cl, ZnCl2 and

MnO2 to which starch is added to make a thick

paste. The graphite rod which is immersed in the

centre of the cell acts as cathode. This is a

primary cell without fluid component. The anode

of the cell is a Zn-container containing an

electrolytic mixture of NH4Cl, ZnCl2 and MnO2 to

which starch is added to make a thick paste. The

graphite rod which is immersed in the centre of

the cell acts as cathode.

Cell reactions: Anode: Zn(s) Zn2+

(aq)+ 2e-

Cathode: MnO2(s) + NH4+

(aq) + 2e- MnO(OH)-(s) + NH3

Net Reaction: Zn(s) + NH4+

(aq) + MnO2(s) Zn2+(aq) + MnO(OH)-

(s) + NH3(g) + Elect.Energy

Reaction between NH4+ (from NH4Cl) and OH- formed at cathode evolves NH3 gas which disrupts the

current flow. This is prevented by the reaction of NH3 (g) with ZnCl2 to form complex

[Zn(NH3)2Cl2].

ZnCl2+ 2NH3 [Zn(NH3)2] Cl2

The reactions involved in the dry cell cannot be reversed by passing electricity back through the cell.

Hence dry cell is a primary cell.

Advantages:

Low price; gives voltage of about 1.5 V; normally works without leaking (leak proof cells); possess

high energy density; non- toxic; contains no liquid electrolytes.

Applications:

Dry cell is used in consumer electronic devices like calculators, transistor radios, flash lights, quartz

wall clocks, walkman etc. and in small portable appliances where small amount of current is needed.

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2. Alkaline Cell

Anode: Zn-container, Cathode: Graphite rod

Electrolyte: Mixture of KOH (aq), Zn-powder and MnO2(s) to which starch is added to make a gel.

Voltage: 1.5 V

Construction: It is an improved form of the dry

cell, in which the electrolyte NH4Cl is replaced by

KOH. In alkaline battery, the powdered zinc is

mixed with KOH and MnO2 to get a gel. A carbon

rod (graphite), acts as cathode, is immersed in the

electrolyte in the centre of the cell. The outside

cylindrical zinc body is made of zinc, act as anode.

The EMF of the cell is 1.5 V.

Cell reactions: Anode: Zn(s) + 2OH-(aq) Zn2+

(aq) + 2e-

Cathode: 2MnO2(s) + H2O (l) + 2e- Mn2O3(s) + 2OH-(aq)

Net Reaction: Zn(s) + 2MnO2(s) + H2O (l) Zn2+ (aq) + Mn2O3(s) + 2OH-

(aq) + Elect.Energy

Advantages: 1. Zinc does not dissolve readily in a basic medium.

2. The life of alkaline battery is longer than the dry cell, because there is no corrosion on zinc.

3. Alkaline battery maintains its voltage, as the current is drawn from it.

Applications: Alkaline cell is used in consumer electronic devices like calculators, transistor radios,

flash lights, quartz wall clocks, walkman etc. and in small portable appliances where small amount

of current is needed.

3. lithium Cell - Primary Battery

Anode: Lithium metal, Cathode: Metal oxide or TiS2, Electrolyte: Solid polymer and Voltage: 3.0 V

Construction: Lithium cell is a primary cell, Lithium metal acts as anode. Metal oxide or TiS2 acts as

cathode. Solid polymer will act as an electrolyte; this polymer is packed in between the electrodes. The

electrolyte permits the passage of ions but not that of electrons. The EMF of the cell is 3.0V

Cell reactions: Anode: Li(s) Li+ + e-

Cathode: TiS2(s) + e- TiS2-

Net Reaction: Li(s) + TiS2(s) Li+ + TiS2- + Elect.Energy

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Advantages:

1. In this cell voltage is high, 3.0V.

2. Lithium is a light weight metal (7gms).

3. Lithium has the most negative E° value.

3. In Lithium cell all constituents are in solid form so

that there is no leakage.

4. The battery can be made in a variety of sizes and

shapes.

Applications: 1. Button sized Lithium batteries are used in calculators, watches, cameras, mobile phones, laptop

computers, etc.,

Secondary Batteries

The cells in which the cell reaction is reversed by passing direct current in opposite direction i.e. it

can operate both as a voltaic cell and as an electrolytic cell. The secondary batteries can be used

through a large number of cycles of discharging and charging. They are used as a source of DC power.

Eg. Lead –acid storage cell, Ni-Cd battery and Lithium ion batteries.

1. Lead acid storage cell - Secondary Battery

Anode: Pb(s), Cathode: PbO2(s), Electrolyte: 20% of H2SO4 (aq) with density of 1.20g/cm3

Voltage: 6.0 V for 3 pairs, 12.0 V for 6 pairs

Construction: It consists of lead –antimony alloy coated with lead dioxide (PbO2) as cathode and

spongy lead as anode. The electrolyte is a 20% solution of H2SO4. The storage cell can operate both

as voltaic cell and electrolytic cell. It acts as voltaic cell when supplying energy and as a result

eventually becomes rundown. The cell operates as electrolytic cell when being recharged. The cell

consists of a series of Pb-plates (negative plates) and PbO2 plates (positive plates) connected in parallel.

The plates are separated from adjacent one by insulators like wood, rubber or glass fiber.

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Cell reactions

Discharging (voltaic cell): During the discharging the battery, H2SO4 is consumed the density of H2SO4

falls to 1.20g/cm3, then battery needs recharging. In discharging, the cell acts as a voltaic cell where

oxidation of lead occurs. In charging, the cell is operated like an electrolyte cell and electric energy is

supplied to it.

At Anodic (-): Pb(S) → Pb2+ + 2e-

Pb2+ + SO42-

(aq) → PbSO4(S)

At Cathodic (+): PbO2(S) + 4H++ 2e- → Pb2+ +2H2O

Pb2+ + SO42-

(aq) → PbSO4(S)

Net reaction: Pb(S) + PbO2(S) + 4H++ 2SO42-

(aq) → 2PbSO4(S) +2H2O + Elect.Energy

Charging (electrolytic cell): During this process, lead is deposited at the cathode, PbO2, is formed at

the anode and H2SO4 is regenerated in the cell.

At Anodic (+): PbSO4(S) + 2H2O → PbO2(S) + 4H+ + SO42-

(aq)+ 2e-

At Cathodic (-): PbSO4(S) + 2e- → Pb(S) + SO4

2- (aq)

Net reaction: Elect.Energy + 2PbSO4(S) + 2H2O → Pb(S) + PbO2(S) + 4H+ +2SO42

(aq)

Applications: The lead storage cells are used to supply current for electrical vehicles, gas engine

ignition, telephone exchanges, electric trains, mines, laboratories, hospitals, broadcasting stations,

automobiles and power station.

2. Nickel-Cadmium cell - Secondary Battery

Anode: Cd (s), Cathode: NiO2 (s), Electrolyte: 20-28% aqueous KOH jelled with a jelling agent

Voltage: 1.4 V

Construction: This battery consists of Cd anode and cathode is composed of a paste of nickel dioxide

[NiO2]. These cells are rechargeable and give a voltage of 1.4V each. They can be connected in series

to give Ni-Cd battery. Electrolyte is 20-28% aqueous KOH jelled with a jelling agent.

Discharging (voltaic cell):

Anode(-): Cd(s) + 2OH-(aq) Cd(OH)2(s) + 2e-

Cathode(+): NiO2(s) + 2H2O + 2e- Ni(OH)2(s) + 2OH-(aq)

Net Reaction: Cd(s) + NiO2(s) + 2H2O Cd(OH)2(s)+ Ni(OH)2(s)+ Elect.Energy

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Charging (electrolytic cell):

Cathode(-): Cd(OH)2(s) + 2e- Cd(s) + 2OH-(aq)

Anode(+): 2Ni(OH)2(s) + 2OH-(aq) NiO2(s) + 2H2O +2e-

Net reaction: Elect.Energy + Cd(OH)2(s) + 2Ni(OH)2(s) Cd(s) + NiO2(s) + 2H2O

Applications:

1. It is used - in flash lights, photoflash units and portable electronic equipments.

2. In emergency lighting systems, alarm systems.

3. In aircrafts and space satellite power systems.

4. For starting large diesel engines and gas turbines etc.

3. Lithium ion cell - Secondary Battery

Anode: Layers of porous carbon, Cathode: layers of lithium metal oxide, Electrolyte: Polymer gel

Voltage: 3.8 V – 4.2V

Construction: It is rechargeable secondary cell. In lithium battery consists of a layers of porous carbon

acts as anode and layers of lithium metal oxide acts as cathode. Electrolytes, generally a polymer gel,

act as separator between the electrodes. The separator permits the passage of ions but not that of

electrons.

Cell reactions during discharging (voltaic cell):

During discharging, the Li+ ions flow back through

the electrolyte form negative electrode to the

positive electrode. Electrons flow from the negative

electrode to the positive electrode. The Li+ ions and

electrons combine at the positive electrode and

deposit there as Li.

Li1-xCoO2 + CLix + E.E LiCoO2 + C + Elect.Energy

Cell reactions during charging (electrolytic cell):

During charging, the Li+ ions flow from the

positive electrode to negative electrode through

electrolyte.

LiCoO2 + C + Elect.Energy Li1-xCoO2 +

CLix

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Advantages:

1. Lithium –ion batteries are high voltage and light weight batteries.

2. It is smaller in size.

3. It produces three times the voltage of Ni-Cd batteries.

Applications:

1. It is used in cell phone, note PC, portable LCD TV, semiconductor driven audios.

Fuel cells

A fuel cell is a galvanic cell in which chemical energy of a fuel – oxidant system is converted directly

into electrical energy in a continuous electrochemical process. The chemical energy provided by the

fuel cell is stored outside the cell. The fuel and oxidants are continuously and separately supplied to the

electrodes of the cell where they undergo reactions. Fuel cells are capable of supplying current as long

as reactants are supplied.

Eg. Hydrogen - oxygen fuel cell and Methanol - oxygen fuel cell

1. Hydrogen – Oxygen Fuel cell

Anode & Cathode: Porous graphite electrodes impregnated with finely divided Pt/Pd

Electrolyte: 25% KOH held in asbestos matrix

Voltage: 0.8 V – 1.0V

Construction: Both anode and cathode consists porous

graphite electrodes impregnated with finely divided Pt/Pd.

The electrolyte is 25% KOH held in asbestos matrix. The

reactants (fuel + oxidant) are constantly supplied from

outside and the products are removed at the same rate as they

are formed. The product discharged is water and the standard

EMF of the cell is 0.8 V – 1.0V.

Cell reactions: Anode: 2H2(g) + 40H- (aq) 4H2O(l) + 4e-

Cathode: O2(g) + 2H2O(l) + 4e- 4OH-(aq)

Net Reaction: 2H2(g) + O2(g) 2H2O(l) + Elect.Energy

Applications:

1. It is used as energy source in space shuttles e.g. Apollo spacecraft.

2. Used in small- scale applications in submarines and other military vehicles.

3. Suitable in places where environmental pollution and noise are objectionable.

4. Drinking water source in space crafts, submarines and other military applications.

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2. Methyl Alcohol – Oxygen Fuel cell

Anode & Cathode: Nickel sheet covered by Pt-Pd catalyst

Electrolyte: 25 % of KOH solution

Voltage: 1.23V

Construction: The cell consists of methanol and used as the

fuel and oxygen (or) air as the oxidant. Anode: Nickel sheet

covered by Pt-Pd catalyst, Cathode: Silver on nickel sheet

Electrolyte: KOH solution. In this cell methanol and oxygen

are fed to the electrolyte. CO2 produced is absorbed by the

electrolyte KOH and converted in carbonates. Formation of

carbonates decreases the cell efficiency because of the

increasing, concentration and polarization at the electrode

surface and decreases the conductivity of the electrolyte.

Cell reactions:

At Anode: CH3OH(l) + 6OH− → CO2(g) + 5H2O (l) + 6e−

At Cathode: 3/2O2(g) +3H2O (l)+ 6e−→ 6OH−(aq)

Net Reaction: CH3OH(l) +3/2O2→ CO2(g) +2H2O (l) + Elect.Energy

Advantages of methyl alcohol-oxygen fuel cell:

1. Easy to transport.

2. Do not require complex steam reforming operations.

3. These fuel cells are targeted to portable applications

4. It is excellent fuel due to presence of high concentration of hydrogen in methanol.

5. There is zero emission by the cells hence the fuel cells are eco friendly.

Applications:

1. Because of light weight these fuel cells are preferred for space crafts and product H2O is a valuable

fresh water source for astronauts.

2. The major application of methyl alcohol-oxygen fuel cells is a fuel for fuel cell motor vehicles like

NECAR-5 in Japan, USA.

Advantages of Fuel cells over other cells

1. High efficiency in energy conversion.

2. The product H2O is a drinking water source for astronauts.

3. Noise and thermal pollution are low.

4. Low maintenance cost.

5. It is an energy storage system for space applications.

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Important questions

1.Define a reference electrode and explain Construction and functioning of standard hydrogen

electrode.

2. Explain Construction and functioning of Calomel electrode.

3. Explain with a neat diagram the function of Reference and ion Selective electrodes?

4. What is electrochemical cells and explain construction and functioning of Galvanic cell.

5. Explain Nernst equation and write its applications.

6. What is electrochemical series and write its applications.

7. Define concentration cells? Explain Electrolyte concentration cell in detail.

8. A zinc road is placed in 0.1M-ZnSO4 solutions at 298 K. write the electrode reaction and calculate

the potential of the electrode. Zn = - 0.7 V?

9. What are primary batteries and explain construction and functioning of Dry/Leclanchi battery.

10. Explain Construction and functioning of alkaline cell.

11. Why primary lithium battery best with other primary batteries and explain construction and

functioning of lithium battery.

12. What are secondary batteries and explain construction and functioning of Lead acid-storage

battery.

13. Explain Construction and functioning of Ni-Cd cell.

14. Why secondary lithium ion batteries are best with other secondary batteries and explain

construction and functioning of lithium ion battery.

15. What is a fuel cell and explain Hydrogen – Oxygen Fuel cell.

16. What are the advantages of fuel cells over primary and secondary batteries and Explain

Methanol – Oxygen Fuel cell.

Objective questions

1. Which of the following is weak electrolyte [ ]

(a) HCl (b) NaOH (c) NH4OH (d) NaCl

2. Pure water does not conduct electricity because, it is [ ]

(a) Acidic (b) Basic (c) Almost unionized (d) Decomposed easily

3. The potential of standard hydrogen electrode is taken as [ ]

(a) 1 volt (b) 0 volt (c) 10 volts (d) None of these

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4. Which statement is correct about the Daniel cell [ ]

(a) The spontaneous cell reaction involves the reduction of Cu+2 to Cu.

(b) The spontaneous cell reaction involves the oxidation of Cu by Zn.

(c) The spontaneous cell reaction involves the reduction of Zn+2 to Zn.

(d) The Daniel cell is an example of an electrolytic cell.

5. A galvanic cell converts [ ]

(a) Electrical energy into chemical energy.

(b) Chemical energy into electrical energy.

(c) Electrical energy into heat energy.

(d) Chemical energy into heat energy

6. Electrochemical series the elements are arranged in the [ ]

(a) Decreasing order of standard reduction potentials.

(b) Increasing order of standard reduction potentials.

(c) Increasing order of equivalent weight.

(d) Increasing order of oxidation potentials.

7. The ion selective electrodes make use of [ ]

(a) Membrane

(b) Mercury

(c) KCl solution

(d) Platinum electrode

8. The electrode potential is the tendency of a metal [ ]

(a) To gain electrons.

(b) To loss electrons.

(c) Either to gain or loss electrons.

(d) None of the above

9. Calomel electrode is constructed using a solution of [ ]

(a) Saturated KCl

(b) Saturated CaCl2

(c) Saturated MgCl2

(d) Saturated NaCl

10. A storage cell is a device that can operate [ ]

(a) Both as voltaic cell and electrolytic cell.

(b) As voltaic cell

(c) As electrolytic cell.

(d) As concentration cell

11. Which one of the following is a primary cell? [ ]

(a) Ni-Cd cell (b) Lead-Acid cell (c) lithium ion cell (d) Alkaline cell

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12. The cathode of Ni-Cd battery is composed of [ ]

(a) Cd

(b) Ni

(c) Paste of NiO(OH)

(d) Paste of Cd (OH)2

13. A fuel cell converts [ ]

(a) Chemical energy of fuels directly to electrical energy.

(b) Chemical energy of fuels directly to heat.

(c) Chemical energy of fuels directly to pressure.

(d) Electrical energy of fuels directly to chemical energy.

14. What is the voltage produced by the H2 - O2 fuel cell. [ ]

(a) 1.0 V (b) 1.23 V (c) 2.0 V (d) 0.5 V

Fill in the blanks

1. The substance which allows electric current to pass through it is called____________________.

2. The substance which conducts electricity without decomposition is called____________________.

3. In an electro chemical cell electron flows from _______ to _______.

4. ________ are the cells which do not store the energy.

5. Electro chemical series is based on ______________.

6. A device which converts electrical energy to chemical energy is called__________________.

7. The substance which conducts electricity without decomposition is called____________________.

8. The standard electrode potential of saturated calomel electrode at 250c is________________.

9. When two identical electrodes are dipped in two electrolytic solutions of different concentrations,

the resulting obtained is called______________.

10. A cell whose reaction is not reversible is called _________.

11. In lithium primary cell, the anode is _______________.

12. In alkaline cell, electrolyte is ___________________________________.

13. Cathode in dry cell is _______.

14. The electrolyte in a Lead-acid battery is _____________.

15. In lead-acid storage battery during discharging operation the concentration of H2SO4 ___________.

16. In Ni-Cd cell the cathode is ________________.

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Two Marks and Three Marks Questions

1. Define single electrode potential mention the factors affecting it.

2. Define reference electrode give one example.

3. Write the construction of standard Hydrogen electrode

4. Write the construction of saturated calomel electrode

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5. Write the Nernst equation.

6. What is electro chemical series? Write its applications.

7. What is electro chemical cell give one example?

8. What is Concentration cell.

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9. What is EMF and calculate the potential of a zinc road is placed in 0.1M-ZnSO4 solutions at 298 K.

write the electrode reaction and calculate the potential of the electrode. Zn = - 0.7 V?

10. Write the chemical reactions involved in Methanol-Oxygen Fuel cell.

11. Write the differences between Primary and Secondary batteries

12. Write Discharge and charging reactions of lead-acid battery.

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13. Write Discharge and charging reactions of Ni-Cd battery.

Draw the neat diagrams of the following

1. Standard Hydrogen Electrode.

2. Calomel electrode:

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3. Glass electrode:

4. Daniel cell:

5. Electrolytic Concentration cell:

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6. Dry Cell / Leclanche Cell:

7. Alkaline Cell:

8. Lithium Battery:

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9. Lead-Acid Storage Battery:

10. H2-O2 fuel Cell:

11. CH3OH-O2 fuel Cell::