Galvanic cell, Nernst equation, Calomel electrode, Quinhydrone electrode, Glass electrode, Dry cell, Lead Acid battery, Li ion cell, Methanol oxygen fuel cell

1. MATRUSRI ENGINEERING COLLEGE
DEPARTMENT OF SCIENCES AND HUMANITIES
SUBJECT NAME: CHEMISTRY
FACULTY NAME: VISHNU THUMMA
MATRUSRI
ENGINEERING COLLEGE
TOPIC: ELECTROCHEMISTRY AND BATTERIES

2. CHEMISTRY
COURSE OBJECTIVES:
➢Correlate the properties of materials with their internal structure and use
the for Engineering applications
➢Apply the principles of electrochemistry in storage of electrical energy in
batteries.
➢Gains knowledge in causes of corrosion and its prevention.
➢Attains knowledge about the disadvantages of hard water for domestic
and industrial purposes.
➢Also learns the techniques of softening of hard water and treatment of
water for drinking purpose.
➢Exposed to qualitative and quantitative parameters of chemical fuels.
➢Aware eco-friendly materials and processes.
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ENGINEERING COLLEGE
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3. CHEMISTRY
COURSE OUTCOMES: After completion of course students will be able to
➢Analyze and apply knowledge of electrodics in calculation of cell
potentials of batteries.
➢Identify the different types of hardness and alkalinities in water and
make use of softening methods, analyze and apply the knowledge of
corrosion for its prevention.
➢Discuss different types of polymers based on their end on use and the
need to replace the conventional polymers with polymers of engineering
applications.
➢Identify and analyze different types of chemical fuels for domestic and
automobile applications.
➢Outline the principles of green chemistry for sustainable environment
and preparation of biodiesel from renewable sources.
MATRUSRI
ENGINEERING COLLEGE
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4. 4
Course Title Chemistry
Course Code BS105CH
Programme Common for all branches
Semester I (CSE, EEE) II (ECE, MECH, IT)
Course Type Basic Sciences
Regulation AICTE Model curriculum
Course Structure
Theory Practical
Lectures Tutorials Credits Laboratory Credits
3 1 4 3 1.5
Marks Distribution
CIE SEE - CIE SEE
30 70 - 25 50
COURSE DESCRIPTION
MATRUSRI
ENGINEERING COLLEGE

5. UNIT-I ELECTROCHEMISTRY AND BATTERIES
MATRUSRI
ENGINEERING COLLEGE
INTRODUCTION:
•The electrochemistry is the subject which deals with the application of electricity with
chemical species.
•It is the branch of physical chemistry which deals with conversion of electrical energy
into chemical energy or vice versa.
•In electrolytic cells the conversion of electrical energy to chemical energy takes place;
whereas the chemical energy is converted to electrical energy in a galvanic cell.
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OUTCOMES: After completion of course students will be able to analyze and apply knowledge
of electrodics in calculation of cell potentials of batteries.

6. MATRUSRI
ENGINEERING COLLEGE
ELECTRIC CONDUCTOR:
Any substance which allows the passage of electrical current through it is called an
electric conductor.
Metallic conductors
• Involve the flow of current with
which does not undergo any
chemical change.
• The flow of current in the form of
flow of electrons, hence they are
also called as electronic
conductors.
• Passage of electricity does not
cause any change except a small
rise in temperature.
• They obey Ohm’s law but not
Faraday’s law.
Eg: Cu, Al etc.
Electrolytic conductors
• Involve the flow of current followed
by a chemical change.
• In solutions the current flows in
the form of movement of ions.
• Passage of electricity causes the
transfer of matter, and rise in
temperature increases the rate of
dissociation of electrolyte, thus,
conductance increases.
• Ex: solutions of acids, bases and
salts.
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7. CONTENTS: ELECTROLYTE – ELECTROLYSIS - ELECTROLYTIC CELL
MODULE-1: ELECTROLYTIC CELL
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ENGINEERING COLLEGE
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Electrolysis: The process of chemical
decomposition of an electrolyte by the passage of
electricity through its molten or dissolved state.
Electrolyte: A substance that produces an electrically conducting solution
when dissolved in a polar solvent, such as water.
It splits up into charged particles called ions.
The positively charged ions are called cations while the negatively charged
ions are called anions.
Electrolytic cell: The device in which the
process of electrolysis is carried out.

8. MATRUSRI
ENGINEERING COLLEGE
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Electrolytic cell
The cell contains aqueous solution of an electrolyte, in which two metallic rods
(electrodes) are dipped which are connected to a battery.
ANODE (+): Positive electrode.
Current enters the cell.
CATHODE (-): Negative electrode
Current leaves the cell.
When electric current is passed through the solution, the
ions respond to the applied potential difference and their
movement is directed towards oppositely charged
electrodes.
The cations move towards the negatively charged
electrode while anions move towards the positively
charged electrode.
Products are formed at the respective electrodes takes place
due to oxidation at the anode and reduction at the cathode. `

9. ELECTRICAL ENERGY IS CONVERTED INTO CHEMICAL ENERGY
A CHEMICAL CHANGE IS BROUGHT UP BY PASSING ELECTRICITY.
In a Electrolytic cell:
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ENGINEERING COLLEGE
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OR
HCl(l) → H+ + Cl-
Anode: Cl- → ½ Cl2 + e- (Oxidation)
Cathode: ½ H+ + e- → H2 (Reduction)

10. 1. Electrolytic cell converts _________ energy into _______ energy.
Quiz MATRUSRI
ENGINEERING COLLEGE
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a. chemical, electrical
b. electrical, chemical
c. electrical, mechanical
d. mechanical, electrical
2. Electrolytic conduction involves
a. Chemical change
b. Migration of ions
c. Mobility of ions
d. All the above
3. At anode ________ reaction occurs.
a. reduction
b. oxidation
c. redox
d. All the above
4. Reduction reaction occurs at __________
a. anode
b. cathode
c. Both a&b
d. None

11. A redox reaction is utilized to get electrical energy.
CONTENTS: GALVANIC CELL – CONSTRUCTION - CELL REACTION – SALT BRIDGE – EMF OF CELL
MODULE-2: GALVANIC CELL
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ENGINEERING COLLEGE
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Chemical energy is converted into electrical energy.
It consists of two half cells connected through
an external wire and are separated with a salt
bridge.
Each half cell consists of an electrode which
is dipped in a suitable electrolytic solution.

12. MATRUSRI
ENGINEERING COLLEGE
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DANIEL CELL
It consists of two electrodes.
Zn rod dipped in Zinc Sulphate solution. (Anode)
Cu rod dipped in Copper Sulphate solution.
(Cathode)
Both the solutions are separated with a semi
permeable membrane or salt bridge.
Salt bridge prevents the diffusion of the two liquids
but allows the passage of ions through it, when the
flow of electric current takes place.
When the circuit is complete the flow of electric
current takes place.
At anode (-): Zn → Zn2+ + 2e- (Oxidation) (Metal dissolved)
At cathode (+): Cu2+ + 2e- → Cu (Reduction) (Metal deposited)

13. MATRUSRI
ENGINEERING COLLEGE
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The movement of electrons from zinc to copper produces a current in the circuit.
The net chemical change is described as the cell reaction.
At Anode(-) : Zn → Zn2+ + 2e- (Oxidation half cell reaction)
At Cathode(+): Cu2+ + 2e- → Cu (Reduction half cell reaction)
Cell reactions: Zn + Cu2+ → Zn2+ + Cu (Redox reaction)
Salt Bridge: It is an inverted ‘U’ shape tube open at both the ends, it
contains inert electrolytes such as KCl, KNO3, NH4NO3 etc., mixed
with agar-agar gel to make it as semi solid paste.
Role of salt bridge:
•It connects the two solutions and prevents their intermixing.
•It prevents the accumulation of charges around the electrode.
•It allows the movement of anions from cathodic solution to anodic
solution.
•It maintains electrical neutrality of solution and this completes the
circuit.

14. The negative electrode is written on the extreme left and the positive electrode is
on the extreme right.
Cell Notation (Representation):
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ENGINEERING COLLEGE
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The single vertical line indicates the electrode-electrolyte interface.
The double vertical lines between the two liquids signifies the salt bridge or semi
permeable membrane separating the two half cells.
The emf of the cell is written on extreme right.
Zn ZnSO4(1.0 M) CuSO4(1.0 M) Cu 1.1V
An inert electrode is indicated in ( ).
(Pt) Q,QH2 /H+ (Pt) H2 /H+ (Pt) Hg, Hg2Cl2(s) /KCl

15. EMF OF THE CELL:
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ENGINEERING COLLEGE
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It is equal to the sum of oxidation potential and reduction potential.
Ecell = EOx + ERed
= E0
Zn/Zn
2+
(SOP) + E0
Cu/Cu
2+
(SRP)
= +0.76 + 0.34 = +1.10 V
When Standard reduction potentials (SRP) are taken into account
Ecell = Ecathode - Eanode
When emf of cell is positive then the cell reaction is feasible
Since, ΔG = - nFE

16. MATRUSRI
ENGINEERING COLLEGE
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1. Which of the following is not related to galvanic cell
Quiz
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ENGINEERING COLLEGE
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a. Chemical energy is converted to electrical energy
b. It requires source of electrical energy
c. Flow electrons takes place from anode to cathode
d. As salt bridge is required
2. The purpose of salt bridge in galvanic cell is
a. Prevents accumulation of charge.
b. Provide a path for mobility of electrons
c. Prevents electrical neutrality.
d. All the above
3. Standard reduction potential of copper is
a. -0.76 V b. +0.76 V c. +0.34 V d. -0.34V
4. For spontaneous cell reaction emf of cell should be
a. positive b. negative c. 0.0V d. none

17. CONTENTS: NERNST EQUATION – DERIVATION – APPLICATIONS –NUMERICAL PROBLEMS
MODULE-3: NERNST EQUATION
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ENGINEERING COLLEGE
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Let us consider a galvanic cell whose cell reaction is written as:
aA + bB → cC + dD
Nernst equation establishes the relationship between concentrations or activities
of electrolyte solutions with cell emf.
The equilibrium constant for above equation:
= Activity coefficient = Q

18. MATRUSRI
ENGINEERING COLLEGE
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Nernst equation can be derived from Vant Hoff’s isotherm:
From the thermodynamics:
By substituting ∆G and ∆G0 values in equation (1), we get
By dividing the equation (2) with -nF on both sides
ΔG = ∆G0 + RTlnQ ---------- (1)
∆G = - nFE and ∆G0 = -nFE0
-nFE = -nFE0 + RTlnQ ---------- (2)
E = E0 -
RT
nF
lnQ
E = E0 -
2.303RT
nF
log
[C]c[D]d
[A]a[B] b
E = E0 -
2.303RT
nF
log
[Products]
[Reactants]

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ENGINEERING COLLEGE
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R = Gas constant = 8.314 J K-1mol-1
T = Temperature in Kelvins (at 250C = 298 K)
F = Faraday = 96500 C
n = number of electrons (Faradays)
At 250C, by substituting the R, T & F values,
Nernst equation can be simplified as:
The standard cell emf is equal to the cell emf when the activities of both
reactants and products are equal to unity.

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For Redox reaction: Ex: Zn + Cu2+ → Zn2+ + Cu
Since, [Zn] and [Cu] are unity.
For a reduction reaction: Mn+ → M + ne-
At 250C,
Since, [M] = 1

21. MATRUSRI
ENGINEERING COLLEGE
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Applications of Nernst Equation:
It is used to:
Study the effect of electrolyte concentration on electrode potential.
To calculate single electrode potential of a half cell.
Used for calculation of cell potential under non standard conditions.
pH of a solution can be calculated. (Ecell = E°cell – 0.0592 pH)
Helpful to determine the unknown concentration of one of the ionic species
of cell if E0 cell and concentration of other species is known.
Used for finding the valence of number of e- involved in a reaction.

22. MATRUSRI
ENGINEERING COLLEGE
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Q: Calculate the emf of a cell in which iron is in contact with FeSO4 (0.1M) electrolyte and
Cu which is placed in CuSO4 (0.01M) solution. The SRPs of Fe and Cu are -0.44V and
+0.34V respectively.
Numerical Problems
Cell Notation: Fe FeSO4 CuSO4 Cu
Cell reactions: At anode: Fe → Fe+2 + 2e- (Oxidation)
At cathode: Cu2+ + 2e- → Cu ( Reduction)
Cell reaction: Fe + Cu2+ → Fe2+ + Cu
E0
cell = E0
cathode - E0
anode
= 0.34 – (-0.44) V = 0.78v
= 0.78 V
At 250C,

23. Questions
MATRUSRI
ENGINEERING COLLEGE
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1. Derive Nernst equation? Write its applications?
2. Calculate the single electrode potential of Cu dipped in 0.01M CuSO4
Solution, the given E0
Cu
2+
/Cu = 0.34 V.
3. Write the cell reaction and calculate the emf of the following cell at 250C:
Zn(s) l Zn2+(0.001M) ll Ag+(0.0001M) l Ag(s)
Given E0
Zn
2+
/Zn= -0.76 V and E0
Ag
+
/Ag = 0.80 V.

24. MODULE-4: CALOMEL ELECTRODE
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ENGINEERING COLLEGE
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It is a secondary reference electrode consisting of mercury-mercurous chloride (Hg-Hg2Cl2)
paste in contact with KCl solution.
It consists of a plug at the bottom of which a
small amount of Hg is placed.
Hg is covered with a paste of solid mercury-
mercurous chloride (calomel paste).
The solution of KCl is placed over the paste.
A platinum wire is employed for electrical contact.
The potential of calomel electrode depends
on KCl solution.
It is represented as: Pt | Hg, Hg2Cl2 | KCl

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Electrode reaction: Hg2Cl2 + 2e- 2Hg + 2Cl-
Electrode potential:
The calomel electrode is reversible to Cl- ion concentration.
The standard reduction potential of calomel electrode at 250C is
0.1M KCl l Hg2Cl2(s) l Hg, Pt 0.3338V Decinormal Calomel Electrode(DNCE)
1.0M KCl l Hg2Cl2(s) l Hg, Pt 0.2800V Normal Calomel Electrode (NCE)
Sat. KCl l Hg2Cl2(s) l Hg, Pt 0.2415V Saturated Calomel Electrode (SCE)

26. MATRUSRI
ENGINEERING COLLEGE
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QUIZ
1. Calomel electrode is reversible to _____ ions
a) K+ b) Cl- c) Hg d) Hg2Cl2
2. Calomel is a paste of
a) Hg2Cl2 b) Hg c) KCl d) all
3. The reduction reaction of calomel electrode is
a) Hg2Cl2 + 2e- 2Hg + 2Cl-
b) 2Hg + 2Cl- Hg2Cl2 + 2e-
c) Pt/Hg,Hg2Cl2/KCl
d) none

27. MATRUSRI
ENGINEERING COLLEGE
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MODULE-5: QUINHYDRONE ELECTRODE
Quinhydrone is a 1:1 equimolar mixture of quinone and hydroquinone which
exist in equilibrium in presence of H+ ions.
O
O
2H 2e
OH
OH
Quinone (Q) Hydroquinone (QH2 )
+
+
+ -
Quinhydrone electrode is setup by adding a pinch of Quinhydrone to the solution whose
pH is to be measured. A Pt electrode is placed in it to acquire potential.
Electrode representation: Pt |Q, QH2 | H+
(unknown)

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The potential developed on a Pt electrode immersed in this system is given by Nernst
equation as:
EQHE = E°QHE _ 2.303RT
2F
log
[QH2]
[Q] [H ]
+ 2
at 250C the equation may written as
EQHE = E°QHE + 0.0592
2
log
[QH2]
[Q] [H ]
+ 2
Since, Quinhydrone is sparingly soluble salt, concentration terms [Q] and [QH2] are
termed as unity.
Hence, EQHE = E0
QHE + 0.0592 log [H+]
Since, pH = - log [H+]
EQHE = E0
QHE – 0.0592pH OR EQHE = 0.6996 – 0.0592pH
Thus, the potential of Quinhydrone electrode depends on the pH of the solution with
which it is in contact. Therefore it can be used for the measurement of pH.

29. MATRUSRI
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Determination of pH using Quinhydrone electrode:
The potential of Quinhydrone electrode(QHE) is determined by connecting it with a
saturated calomel electrode(SCE)
Pt | Hg, Hg2Cl2(s) |KCl(saturated) || H+
(unknown) | Q, QH2 | Pt
The emf of the cell Ecell = EQHE - ESCE
Ecell = (0.6996 – 0.0592pH) – 0.242 V
pH =
0.4576 - Ecell
0.0592

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ENGINEERING COLLEGE
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Advantages:
•The Quinhydrone electrode is simple to set up and needs no removal of air.
•The reversibility equilibrium is achieved faster than hydrogen gas electrode there by
allowing a quick measurement.
•pH values of solutions containing reducible substances like Cu2+, Cd2+, unsaturated
acids, NO3
- etc., and catalytic poisons can be measured using Quinhydrone electrode,
where hydrogen electrode cannot be used.
Limitations:
The Quinhydrone electrode cannot be used when pH of the solution is more than 8.
Because, at this pH hydroquinone dissociates and undergoes oxidation by the
atmospheric oxygen. This alters the equilibrium between quinone and hydroquinone.

31. MATRUSRI
ENGINEERING COLLEGE
QUIZ
1. What is quinhydrone?
a) 1:1 mixture of quinone and hydroquinone
b) 1:2 mixture of quinone and hydroquinone
c) 2:1 mixture of quinone and hydroquinone
d) none
2. Quinhydrone electrode is sensitive to
a) pH b) quinone
c) hydroquinone d) none
3. The standard potential of quinhydrone electrode is
a) 0.6996 V b) 0.2419 V
c) -0.6996 V d) -0.2419 V

32. MATRUSRI
ENGINEERING COLLEGE
MODULE-6: GLASS ELECTRODE
Glass electrode consists of thin walled glass bulb
containing Pt wire of AgCl coated, dipped in 0.1N HCl
solution.
The glass membrane acts as an ion exchange resin which
is highly pH sensitive.
Glass bulb, when in contact with acid test solution furnishes
a constant H+ ion concentration on both sides which is
responsible for potential difference.
This potential difference is directly proportional to the pH
difference between solutions on both sides of glass.
It is represented as:
Pt | Ag, AgCl | 0.1N HCl | Glass membrane

33. MATRUSRI
ENGINEERING COLLEGE
MEASUREMENT OF pH USING GLASS ELECTRODE
The glass electrode is coupled with a saturated calomel electrode so as to construct a
galvanic cell.
Pt |Ag, AgCl(s) | 0.1NHCl | Glass membrane | H+ (unknown) | Sat.KCl | Hg2Cl2,Hg Pt
The emf of the cell Ecell = ESCE – EGlass
Ecell = 0.242 - (E0
glass - 0.0592 pH)
Electrode Reaction: AgCl + e- AgCl
Eglass = E0
glass - 0.0592 pH

34. QUIZ
34
MATRUSRI
ENGINEERING COLLEGE
1. Glass electrode is a type of
a) inert electrode b) ion selective electrode
c) gas electrode d) metal insoluble electrode
2. The bulb in glass electrode is filled with
a) 0.1N HCl b) 1.0N HCl
c) 0.1 AgCl d) 1.0N AgCl
3. In glass electrode the semi permeable membrane is sensitive to
a) pH b) Ag ions
c) Chloride ions d) solvent
4. The standard potential of glass electrode depends on
a) nature of glass membrane
b) nature of electrolyte
c) nature of solvent
d) all the above

35. MODULE-8: THERMODYNAMICS OF EMF OF CELL
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Enthalpy and EMF:
From thermodynamics ∆G = ∆H - T∆S
According to Gibbs-Helmholtz equation:
∆𝐆 = ∆𝐇 + 𝐓
𝛛 ∆𝐆
𝛛𝐓 𝐏
Since, ∆G = -nFE
−𝐧𝐅𝐄 = ∆𝐇 + 𝐓
𝛛 −𝐧𝐅𝐄
𝛛𝐓 𝐏
−𝐧𝐅𝐄 = ∆𝐇 − 𝐧𝐅𝐓
𝛛𝐄
𝛛𝐓 𝐏
∆𝐇 = −𝐧𝐅𝐄 + 𝐧𝐅𝐓
𝛛𝐄
𝛛𝐓 𝐩
Here, the
𝜕E
𝜕T P
is called
Temperature Coefficient of the
cell.
If
𝜕E
𝜕T P
= 0 , then Electrical
energy is equal to Enthalpy of
the cell reaction. EMF of the cell
increases with temperature.

36. MODULE-8: THERMODYNAMICS OF EMF OF CELL
MATRUSRI
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Entropy and EMF:
From thermodynamics ∆G = ∆H - T∆S --1
Since,
∆𝐇 = −𝐧𝐅𝐄 + 𝐧𝐅𝐓
𝛛𝐄
𝛛𝐓 𝐩
and
∆G = -nFE
Then equation (1) can be written as
−𝐧𝐅𝐄 = −𝐧𝐅𝐄 + 𝐧𝐅𝐓
𝛛𝐄
𝛛𝐓 𝐩
− 𝐓∆𝐒
Then,
𝐓∆𝐒 = 𝐧𝐅𝐓
𝛛𝐄
𝛛𝐓 𝐩
And
∆𝐒 = 𝐧𝐅
𝛛𝐄
𝛛𝐓 𝐩

37. MODULE-8: BATTERY CHEMISTRY
Battery is an electrochemical cell or often several electrochemical
cells connected in series that can be used as a source of direct
electric current at constant voltages.
Battery is a commercial electrochemical cell.
Classification:
Primary Cells. Eg.: Dry cell
Secondary Cells. Eg.: Lead acid cell.
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38. Primary Cells
Produce electricity from chemicals that are sealed into it.
Electrical energy can be obtained at the expense of chemical energy
only as long as the active materials are still present.
Cannot be recharged as the cell reaction cannot be reversed
efficiently by recharging.
The cell must be discarded after discharging.
Ex: Zinc – Carbon battery (Dry cell)
Zinc – Mercuric oxide cell.
Zinc – Silveroxide cell
38
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39. Secondary Cells
Generation of electric energy, that can be restored to its original
charged condition after its discharge by passing current flowing in
the opposite direction.
These cells have a large number of cycles of discharging and charging.
They are known as rechargeable cells, storage cells, or accumulators.
Ex: Lead storage cell.
Nickel- cadmium cell.
Lithium- ion batteries.
39
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40. Differences
Primary Batteries
➢ Cell reaction is irreversible.
➢ Must be discarded after use.
➢ Have relatively short shelf life.
➢ Function only as galvanic cell.
➢ They cannot be used as storage devices.
➢ They cannot be recharged.
Ex: Dry cell.
Li-MnO2battery.
Secondary Batteries
➢ Cell reaction is reversible.
➢ May be recharged
➢ Have long shelf life.
➢ Functions both as galvanic & electrolytic
cell.
➢ They can be used as energy storage
devices. Ex: solar/ thermal energy
converted to electrical energy
➢ They can be recharged.
Ex: Lead acid,
Ni-Cd battery.
40
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41. Anode: Zinc metal container.
Cathode: MnO2 + Carbon (powdered graphite)
Electrolyte: Aqueous paste of NH4Cl and ZnCl2
Cell Notation:
Zn(s) | ZnCl2(aq),NH4Cl(aq) |MnO2(s) |Mn2O3(s) |C
Output voltage = 1.5 V
41
DRY CELL(LECLANCHE CELL)
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42. DRY CELL(LECLANCHE CELL)
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ENGINEERING COLLEGE

43. Working:
Primary Electrode Reactions:
Anode: Zn(s)→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)
43
Secondary Reactions:
2NH4
+
(aq)+2OH-
(aq) → 2NH3(g)+2H2O(l)
Zn2+
(aq)+2NH3(s)+2Cl- → [Zn(NH3)2 Cl2]
Over all reaction:
Zn + 2MnO2 + 2NH4Cl → [Zn(NH3)2Cl2] + H2O + Mn2O3
MATRUSRI
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44. Applications:
❖ In small portable appliances where small amount of
current is needed.
❖ In consumer electronic devices – flash lights, clocks,
remote controllers, walkman etc.
44
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45. Advantages:
❖ Dry cell is cheap.
❖ Normally works without leaking (leak proof cells).
❖ Has a high energy density.
❖ It is not toxic
❖ It contains no liquid electrolytes.
45
MATRUSRI
ENGINEERING COLLEGE

46. Disadvantages:
❖ Voltage drops due to build up of reaction products around the
electrodes when current is drawn rapidly from it .
❖ It has limited shelf life because the zinc is corroded by the
faintly acid, ammonium chloride.
❖ The shelf life of dry cell is 6-8 months.
❖ They cannot be used once they get discharged.
❖ Its emf decreases during use as the material is consumed.
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MATRUSRI
ENGINEERING COLLEGE

47. QUIZ
1. Which of the following sentence is incorrect?
a) primary battery should be discarded after single use.
b) secondary battery should be discarded after single use.
c) battery is a portable source of electrical energy.
d) battery operate as galvanic cell during discharging.
2. Dry cell cannot be recharged because of
a) secondary reactions are not reversible.
b) primary reactions are not reversible.
c) battery dries after discharging.
d) battery leaks if connected to external source of emf.
3. The inert electrode in dry cell is
a) Zn b) Graphite c) MnO2 d) NH4Cl
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48. Anode: Spongy lead on lead grid.
Cathode: Porous PbO2.
Electrolyte: H2SO4(aq)( 20 %)
(density 1.21-1.30g/ml)
Cell Notation:
Pb/PbSO4;H2SO4(aq);PbSO4;PbO2/Pb
Voltage = 2V
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MODULE-9: Lead-Acid Battery
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49. 49
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50. Discharging (as voltaic cell):
Anode: Pb (s) → Pb2+
(aq) + 2e-
Pb2+
(aq) + SO4
2-
(aq) → PbSO4(s)
Pb(s)+ SO4
2-
(aq) → PbSO4(aq) + 2e-
Cathode: PbO2(s)+ 4H+
(aq)+2e- →Pb2+
(aq)+ 2H2O(l)
Pb2+
(aq)+SO4
2-
(aq)→PbSO4(s)
PbO2(s)+4H+
(aq)+SO4
2-
(aq)+2e- → PbSO4(s)+ 2H2O(l)
Overall:
Pb (s)+PbO2 (s)+4H+
(aq)+ 2SO4
2-
(aq) →2PbSO4(s) + 2H2O(l) + Energy
(2H2SO4) 50
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51. Charging (as electrolytic cell)
Cathode:
PbSO4(s)+2H2O(l)→PbO2(s)+ SO4
2-
(aq)+4H+
(aq) +2e-
Anode :
PbSO4(s) + 2e- → Pb(s)+ SO4
2-
(aq)
Net reaction:
2PbSO4 (s)+ 2H2O(aq) → Pb(s)+ PbO2(s) +2H2SO4
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52. A lead storage battery is highly efficient.
Voltage efficiency = average voltage during discharge
average voltage during charge
The voltage efficiency of the lead – acid cell is about 80 %.
A lead – acid battery provides a good service for several years. Its larger
versions can last 20 to 30 years, if carefully attended.
(i.e. longer design life)
It can be recharged. The number of recharges possible range from 300 to
1500, depending on the battery’s design and conditions.
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ADVANTAGES
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53. ✓ The battery’s own internal self – discharging is low.
✓ The length of time that is generally required for recharging process is
less.
✓ Low environmental impact of constituent materials is an added
advantage.
✓ It has sensitivity to rough handling and good safety characteristics.
✓ Ease of servicing as indicated by several local battery service points.
✓ It is a low- cost battery with facilities for manufacture throughout the
world using cheap materials.
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54. Limitations
Self discharge: They are subject to self discharge with H2 evolution at negative
plates and O2 evolution at positive plates.
Pb +H2SO4 → PbSO4 + H2
PbO2 + H2SO4 → PbSO4 +H2O +1/2 O2
SO4
2- +2 H+ (From dissociation of water) → H2SO4
H2O → H+ +OH-
Loss of Water: Due to evaporation, self discharge and electrolysis of water
while charging. Hence water content must be regularly checked and distilled
water must be added.
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55. Corrosion of Grid: Can occur due to overcharging when grid metal
gets exposed to the electrolyte which weakens the grid and
increases the internal resistance of the battery.
Effectiveness of battery is reduced at low temperature due to
increase in the viscosity of electrolyte.
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56. Applications:
• To start Automotive engines.
• Standby/Back-up/Emergency power for electrical
in hospitals, industries, households etc.
• Submarines.
• UPS (Uninterruptible Power Supplies).
• Lighting
• High current drain applications like trains.
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57. QUIZ
1. Which of the following is incorrect in lead acid battery
a) anode is spongy lead b) cathode is PbO2
c) H2SO4 is the electrolyte d) none
2. During discharging the change in oxidation number of Pb at cathode is
a) Pb(II) to Pb(IV) b) Pb(IV) to Pb(II)
c) Pb to Pb(II) d) none
3. During charging the lead acid battery acts as
a) voltaic cell b) galvanic cell
c) electrolytic cell d) all the above
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58. MODULE-10: Lithium Ion Battery
• It consists of a graphite anode and lithium oxide cathode.
• Cathode consists of a layered crystal into which the lithium is
intercalated.
• Experimental cells have also used lithiated metal oxide such as LiCoO2,
NiNi0.3Co0.7O2, LiNiO2, LiV2O5, LiV6O13, LiMn4O9, LiMn2O4, LiNiO0.2CoO2.
• Electrolytes are usually LiPF6 although this has a problem with aluminum
corrosion, and so alternatives are being sought one such is LiBF4.
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59. • Cell reactions in Lithium ion battery involve migration of lithium ions
between the positive and negative electrode.
• No chemical changes are observed at the two electrodes or in the
electrolyte.
• Lithium-Graphite intercalation can be charged and discharged
reversibly where lithium doping is charging reaction and un-doping is
discharging reaction.
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Lithium Ion Battery

60. During discharging lithium ions are extracted from
anode by electrochemical oxidation and inserted
into the cathode by electrochemical reduction.
This cell produces an emf of 3.7V.
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Lithium Ion Battery
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61. During charging of lithium ion batteries, Lithium
ions are extracted by oxidation from LiCoO2
cathode and the extracted lithium ions are
doped by electrochemical reduction into carbon
anode to form Li-GIC.
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Lithium Ion Battery

62. Advantages
• The high energy density.
• Low self-discharge.
• They do not require and maintenance to ensure their
performance.
• The voltage produced by each lithium ion cell is about 3.7
volts.
• Variety of types of lithium ion cell available.
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63. Disadvantages
• They require protection from being over charged and discharged too
far.
• They need to have the current maintained within safe limits.
• Ageing.
• Transportation: Any lithium ion batteries carried separately must be
protected against short circuits by protective covers, etc.
• Cost: They are around 40% more costly to manufacture than Nickel
cadmium cells.
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64. Applications
• These are used in mobile phones and smart phones, laptops and
tablets, digital cameras, electronic cigarettes, hand held game
consoles and torches etc.
• Electric vehicles.
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65. QUIZ
1. Which of the following is correct sentence incase of lithium ion battery
a) lithium ions move between the electrodes during operation
d) Li ion battery consists of graphite anode and lithium oxide cathode
c) Electrolytes are usually LiPF6 or LiBF4 in Li ion battery.
d) all the above
2. During discharging in Li ion battery
a) Li ions are extracted from anode
b) Li ions are extracted from cathode
c) Li ions are inserted to anode
d) none
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66. MODULE-11: 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.
• Cell Schematic Representation:
Fuel/electrode/electrolyte/electrode/oxidant.
e.g. H2-O2; CH3OH-O2
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67. The reactants (i.e. fuel + oxidant) are constantly supplied from outside
and the products are removed at the same rate as they are formed.
Anode:
Fuel+ oxygen → Oxidation products+ ne-
Cathode:
Oxidant + ne- → Reduction products.
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68. Requirements Of Fuel Cell.
Electrodes: Must be stable, porous and good conductor.
Catalyst: Porous electrode must be impregnated with catalyst like Pt, Pd,
Ag or Ni, to enhance otherwise slow electrochemical reactions.
Temperature: Optimum.
Electrolyte: Fairly concentrated.
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69. CH3OH-O2 Fuel Cell
Both electrodes: Made of porous nickel plates impregnated with finely-
divided Platinum.
Fuel: Methyl alcohol.
Oxidant: Pure oxygen / air.
Electrolyte: Conc.Phosphoric acid/Aq.KOH
Operating Temperature: 150-200oC.
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70. CH3OH-O2 Fuel Cell
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71. The emf of the cell is 1.20 V at 25oC.
CH3OH is one of the most electro active organic fuels in the low
temperature range as
*It has a low carbon content
*It possesses a readily oxidizable OH group
*It is miscible in all proportions in aqueous electrolytes.
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72.  At anode:
CH3OH + 6OH- →CO2 + 5H2O + 6e-
• At cathode:
3/2 O2 +3H2O + 6e- →6OH-
Net Reaction:
CH3OH +3/2O2 →CO2 + 2H2O.
It is used in military applications as doesn't produce much noise &
gives high efficiency of energy conversion with a high long life
and in large scale power production. It has been used to power
television relay stations.
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73. Advantages Of Fuel Cells.
• High efficiency of the energy conversion process.
• Silent operation.
• No moving parts and so elimination of wear and tear.
• Absence of harmful waste products.
• No need of charging.
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74. Limitations Of Fuel Cells.
• Cost of power is high as a result of the cost of electrodes.
• Fuels in the form of gases and O2 need to be stored in tanks under
high pressure.
• Power output is moderate.
• They are sensitive to fuel contaminants such as CO,H2S, NH3 &
halides, depending on the type of fuel cell.
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75. QUIZ
1. In which of the following cell the reactants are not stored?
a) Lead-acid battery b) Ni-Cd battery
c) Methanol-oxygen fuel cell d) Dry cell
2. The products of the methanol-oxygen fuel cell are
a) CO2, CH3OH b) CO2, H2O
c) CH3OH, H2O d) H2O, KOH
3. Advantages of the methanol-oxygen fuel cell are
a) silent operation b) no harmful products
c) high efficiency d) all the above
4. Which of the following is not correct with respect to methanol-oxygen fuel cell
a) anode is fed with CH3OH b) cathode is fed with O2
c) KOH is the electrolyte d) electrodes are not filled with catalyst
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