Electrochemistry and Corrosion

1. ENGINEERING CHEMISTRY
Dr. S. Suresh
Assistant Professor of Chemistry
Block-I, Room No. 204
Email: dr.s.suresh@cmrcet.ac.in
Contact: 9959499849
Dr. Suresh Siliveri

2. Dr. Suresh Siliveri
Engineering Chemistry
I. Electrochemistry, Batteries and Corrosion
II. Materials Chemistry-Polymers
III. Energy Sources
IV. Water Technology
V. Engineering Materials.

3. Dr. Suresh Siliveri
Course Outcomes
After completion of the course students will be able to
1. Apply the concept of electrochemistry and corrosion science in various practical
applications.
2. Predict the different engineering applications by preparing various polymers.
3. Summarize the manufacturing process of various fuels and their applications in
daily life.
4. Understand the benefits of treated water as source in steam generation in industrial
application.
5. Illustrate the importance and applications of various advanced engineering
materials.

4. Dr. Suresh Siliveri

5. Learning objectives
In this topic students are going to learn
1. Introduction about electrochemistry
2. Understanding the important Terms and its definitions
3. Galvanic cell.
4. Nernest Equation.
5. Electrodes.
6. Batteries.
7. Fuel cells.
8. Solar Cells and its applications.
Dr. Suresh Siliveri

6. What is electrochemistry about:
 It is a branch of physical chemistry which deals with the study of conversion of chemical energy into
electrical energy and vice versa.
 Electrochemistry deals with the study of electrical properties of solutions of electrolytes and with
the interrelation of chemical phenomenon and electrical energies.
Dr. Suresh Siliveri
Electrolytic Cell:
This device can convert electrical energy in to chemical
energy.
Electrochemical Cell
This device can convert chemical energy in to
electrical energy.

7. Substances around us can be divided into two classes based on their ability of conduct electricity:
Conductors and Non-Conductors
Non-Conductors: Those substances which do not allow electric current to pass through them are
called non-conductors or insulators.
Example: - wood, plastic glass, rubber etc.
Conductors: Those substances which allow electric current to flow through them are called
conductors.
Examples: Copper, Iron, Gold, Silver, Graphite, salt solution etc.
Dr. Suresh Siliveri

8. Metallic Conductors
Conductors
Electrolytes
Metallic Conductors: These conductors
conduct electricity or electric current by
movement of electrons without
undergoing any chemical change during the
process. These conduct electricity in both
solid as well as molten state.
Example: All the metals and Graphite
Electrolytes: Those substances which
conduct electricity only when they are
present in aqueous solution and not in
solid form are called electrolytes.
These conduct electricity by
movement of ions in solutions.
Dr. Suresh Siliveri

9. Conductivity
Electrolytic Conductivity
 Electric current flows by
movement of ions.
 Ions are oxidized or reduced at
the electrodes.
 It involves the transfer of matter
in the form of ions.
 Ohm’s law is followed.
 Resistance decreases with
increase of temperature.
Metallic Conductivity
 Electric current flows by
movement of electrons.
 No chemical change occurs.
 It does not involve the transfer
of any matter.
 Ohm’s law is followed.
 Resistance increases with
increase of temperature.
Dr. Suresh Siliveri

10. Examples of electrolytes are: NaCl, KCl, Na2SO4 etc.
Non-ionic compound or covalent compounds do not conduct electricity in aqueous solution and hence
they are called non-electrolytes. Examples of non- electrolytes are: Urea, Glucose, Sugar etc.
For a substance to conduct electricity; it must either have free electrons or ions which
carry electricity with them.
Electrolytes neither have free electrons nor free ion in solid state although they are ionic
compound. This is because the oppositely charged ions are held together by strong
electrostatic attraction and are not free to move.
But when they are dissolved in water, the two ions split up and become free to move in
solution and now they are free to conduct electricity.
Why do electrolytes not conduct electricity in solid form?
Dr. Suresh Siliveri

11. Strong Electrolytes
Electrolytes
Weak Electrolytes
 Strong Electrolytes are those
electrolytes which dissociate or
ionizes completely in aqueous
solution to give constituent ions.
 For example:
Inorganic salts like NaCl, KCl,
Strong Acid like HCl, H2SO4,
Strong bases like NaOH, KOH etc.
 Weak Electrolytes are those
electrolytes which partially
dissociate or ionizes in aqueous
solution to give constituent ions.
 For example:
weak acid like CH3COOH
weak bases like NH4OH.
Dr. Suresh Siliveri

12. Specific Conductance
The conductance of all the ions
present in 1 cm3 of electrolyte
solution
Equivalent conductance
The conductance of all the ions produced
by one gram equivalent of an electrolyte
in a given volume of solution.
Molar conductance
The conductance of all the ions produced
by ionization of 1 g mole of an electrolyte
when present in V mL of solution.
Conductance
Property of the conductor which facilitates
the flow of electricity
Units of conductance :
ohm-1 or mho or Ω-1
𝐶𝑜𝑛𝑑𝑢𝑐𝑡𝑎𝑛𝑐𝑒(𝐶) =
1
𝑅𝑒𝑠𝑖𝑠𝑡𝑎𝑛𝑐𝑒
=
1
𝑅
Units of Specific conductance : ohm-1 cm-1 or Ω-1 cm-1
Units of Equivalent conductance : ohm-1 cm-1 * cm3 𝑔𝑟. eq−1
ohm-1 cm2 𝑔𝑟. eq−1
Siemens metre-squared per mole
(Sm2mol−1Sm2mol−1)
Dr. Suresh Siliveri

13. Galvanic cell or Voltaic cell
Dr. Suresh Siliveri

14. REPRESENTATION OF A GALVANIC CELL
The following conventions are used in representing an electrochemical cell:
/ 𝑹𝒆𝒑𝒓𝒆𝒔𝒆𝒏𝒕𝒔 𝒑𝒉𝒂𝒔𝒆 𝒃𝒐𝒖𝒏𝒅𝒂𝒓𝒚 // 𝑹𝒆𝒑𝒓𝒆𝒔𝒆𝒏𝒕𝒔 𝒔𝒂𝒍𝒕𝒃𝒓𝒊𝒅𝒈𝒆
1. A galvanic cell is represented by writing the anode (where oxidation occurs) on the left hand side and cathode
(where reduction occurs) on the right hand side.
Anode // Cathode
2. The anode of the cell is represented by writing metal first and then the electrolyte (or the cation of the
electrolyte) with concentrations.
Zn(s)/Zn2+(aq)
(C1M)
3. The cathode is represented by writing the electrolyte first with concentrations and thenmetal.
Cu2+(aq)/Cu(s)
(C2M)
4. The two half cells are separated by a salt bridge, which is indicated by two vertical lines.
Zn(s)/Zn2+(aq)//Cu2+(aq)/Cu(s)
(C1M) (C2M)
5.The value of emf of a cell is written on the right of the cell diagram.
Zn(s)/Zn2+(aq)//Cu2+(aq)/Cu(s) Ecell=1.1V
(C1M) (C2M)
Dr. Suresh Siliveri

15. What is Electrode Potential?
• When a metal is placed in a solution of its own ions, the metal acquires either a positive or negative
charge with respect to the solution. On account of this, a definite potential difference is developed
between the metal and the solution. This potential difference is called electrode potential. (V)
Reduction: Mn+ + ne– → M
• If the metal undergoes oxidation, then the positive metal ions may pass into the solution
• If the metal undergoes reduction, then the positive metal ions from the solutions may get deposited
over the metal.
Oxidation: M → Mn+ + ne–
Dr. Suresh Siliveri

16. When a plate of zinc is placed in a solution having Zn2+ ions, it becomes negatively
charged with respect to solution and thus a potential difference is set up between zinc
plate and the solution.This potential difference is termed the electrode potential of zinc.
Zn metal dipped in a ZnSO4 solution
First, a plate of zinc is placed in a ZnSO4 solution, Zn goes
into the solution as Zn2+ ions.
Zn2+
ZnSO4
solution
Now, the Zn electrode attains the negative charge, due to the
accumulation of valence electrons on the metal.
The negative charges developed on the electrode attract the
positive ions from solution.
Due to this attraction the positive ions remain close to the
metal.
Dr. Suresh Siliveri

17. Cu metal dipped in a CuSO4 solution
Cu2+ solution
Cu2+
when copper is placed in a solution having Cu2+ ions, it becomes
positively charged with respect to solution. A potential difference is
set up between the copper plate and the solution.The potential
difference thus developed is termed as electrode potential of copper.
First, a plate of Cu is placed in a CuSO4 solution, Cu2+ ions in the
solution deposit over the metal.
Now, the Cu electrode attains the positive charge, due to the
accumulation of Cu2+ ions on the metal.
The positive charges developed on the electrode attract the
negative ions from solution.
Due to this attraction, the negative ions remain close to
the metal.
Dr. Suresh Siliveri

18. Depending on the nature of the metal electrode to lose or gain electrons, the electrode
potential may be of two types:
Oxidation potential: When electrode is negatively charged with respect to solution,
i.e., it acts as anode. Oxidation occurs.
M → Mn+ + ne-
Reduction potential: When electrode is positively charged with respect to solution,
i.e., it acts as cathode. Reduction occurs.
Mn+ + ne- → M
The EMF of the cell is equal to the sum of potentials on the two electrodes.
Emf of the cell = EAnode + ECathode
= Oxidation potential of anode + Reduction potential of cathode
Dr. Suresh Siliveri

19. In order to compare the electrode potentials of various electrodes, it is
necessary to specify the concentration of the ions present in solution in
which the electrode is dipped and the temperature of the half-cell.
 The potential difference developed between metal electrode and the solution of
its ions of unit molarity (1M) at 25°C (298 K) is called standard electrode potential.
 Standard emf of a cell is represented by the symbol Eo.
Standard electrode potential.
Single Electrode Potential.
Note: According to IUPAC convention the reduction potential alone is taken as the electrode potential
Dr. Suresh Siliveri

20.  The mere production of electrons is not enough to get the electric current.
emf = Oxidation potential of anode + Reduction potential of cathode
emf = Reduction potential of cathode -Reduction potential of anode
 𝒄𝒆𝒍𝒍= 𝒄𝒂𝒕𝒉𝒐𝒅𝒆− 𝒂𝒏𝒐𝒅𝒆
 𝒄𝒆𝒍𝒍= 𝑹𝒊𝒈𝒉𝒕− 𝑳𝒆𝒇𝒕
emf = Oxidation potential of anode -Oxidation potential of cathode
The emf of cell potential is measured in units of volts (V) and is also referred to as cell voltage.
“ The potential difference between the electrodes which is driving force for the flow of
electrons is known as electromotive force (emf)”.
The electrons liberated at anode have to flow towards cathode through external circuit. It becomes
possible only when there exists a potential difference between the electrodes.
Electromotive force (emf)
 Reduction potential of an electrode is
equal in magnitude but opposite in sign
to its oxidation potential
Dr. Suresh Siliveri

21. Dr. Suresh Siliveri
Nernst equation
 The potential of an electrode depends on concentration and temperature.Therefore a quantitative equation relating the electrode Potential
with these parameters can be tailored.
 Walter Herman Nernst has deduced such an equation for the electrode potential.
 As the reaction proceeds, there is a moment of charges. Hence some amount of electrical work is done. This value becomes maximum at
equilibrium. It is represented as Wmax.
Wmax depends on 1. the no. of coulombs of charge flowing across the interphase and 2. the energy available per coulombs of charge
Wmax = No. of coulombs of charge flowing across the interphase x Energy available per coulombs of charge
Therefore, Wmax = (nF)(E)= nFEcell
As per Faraday’s second law, 1 mole of electrons= 1 Faraday of charge so n mol = nF charge.
As per definition, energy available per coulomb is called potential, measured in volts.
As the reaction is spontaneous, there is decrease in the free energy, the maximum work done by the galvanic cell is equal to decrease in its Gibb’s
energy
-∆G= Wmax=nFEcell
∆G = -nFEEcell
Similarly, ∆G0 = -nFE0
Ecell
G0 and E0 are change in the free energy and electrode potential under standard conditions respectively.

22. Dr. Suresh Siliveri
From thermodynamic equation ∆G=∆G0 + RT ln K
-nFE cell = -nFE0
cell+RT ln K
E cell= E0
cell – ln K K= Equilibrium Constant=
[ ]
[ ]
E cell= E0
cell – ln
[ ]
[ ]
E cell= E0
cell –
.
l𝑜𝑔
[ ]
[ ]
R= 8.314 Jk-1mol-1
T= 298 K
F= 96500 C
E cell= E0
cell –
.
l𝑜𝑔
[ ]
[ ]
Applications of Nernst Equation:
The potential of an electrode and EMF of a cell can be calculated at any temperature and
concentration.
Knowing potential of an electrode, the concentration of the reactant can be calculated.
The concentration of the solution in a galvanic cell can be determined.
The pH of a solution can be calculated by measuring the EMF.

23. Dr. Suresh Siliveri
The emf of the newly constructed cell, E, is determined with a voltmeter.
Measurement of single electrode potential (Or)
Determination of emf of a half-cell
 The emf of a cell that is made of two half-cells can be determined by connecting them to a
voltmeter. However, there is no way of measuring the emf of a single half-cell directly.
A convenient procedure to do so is to combine the given half-cell with another standard half-cell.
Reference electrode: “Reference electrode are the electrode with reference to those, the electrode
potential of any electrode can be measured” It can acts both as an anode or cathode depending upon the
nature of other electrode.
The reference electrodes can be classified in to two types
i) Primary reference electrodes Eg: Standard hydrogen electrode
ii) Secondary reference electrodes Eg: Calomel and Ag/AgCl electrodes

24. Standard Hydrogen Electrode
• The SHE consists of platinum wire in a inverted glass tube.
• Hydrogen gas is passed through the tube at 1atm
• A platinum foil is attached at the end of the wire
• The electrode is immersed in 1M H+ (HCl) solution at 25oC.
• The electrode potential of SHE is zero at all temperatures.
It is represented as
Pt,H2(1atm)/H+(1M)
Dr. Suresh Siliveri

25. Dr. Suresh Siliveri
It is not possible to directly measure the potential of isolated electrode.
But, the potential of a cell can be experimentally determined.
For measuring the single electrode potential, the experimental electrode is coupled with
the reference electrode (say, SHE) by means of a salt bridge and the cell is constructed.
The EMF of the cell constructed is measured and the potential of the electrode is
calculated on the basis of the cell representation.
Measurement of single electrode potential

26. Dr. Suresh Siliveri
Case 1: If the experimental electrode is an oxidation electrode, the cell construction shall be
Pt,H2(1atm)/H+(1M)||Mn+(x M)|M
M|Mn+(x M)||H+(1M)/H2(1atm), Pt
For the measurement of EMF, a sensitive device which draws smaller current for its working is preferred,
vacuum volt meter is employed in place of an ordinary voltmeter. Let the EMF recorded be E1. According to cell
representation.
E1 = ESHE – Eexpt
ESHE = 0, Eexpt= -E1
Case 2: If the experimental electrode is a reduction electrode, the cell construction shall be
Let the EMF recorded be E1. According to cell representation.
E1 = Eexpt – ESHE
ESHE = 0, Eexpt= E1
Note: The single electrode potential for real cells, operating at non-standard condition can be calculated using
Nernst equation.

27. Dr. Suresh Siliveri
Measurement of Zn electrode potential
Limitations
 Construction of SHE is difficult activity.
 It is difficult to maintain unit molar concentration of hydrogen throughout and pass hydrogen always at
exactly 1 atm pressure.
 Presence of arsenic compounds would easily got absorbed on platinum foil there poising the surface it
would be effect equilibrium of the reaction.
 In addition, SHE cannot be used in presence of strong oxidizing & reducing agents.

28. Dr. Suresh Siliveri
 Calomel electrode is the mercury-mercurous chloride electrode.
 It is a commonly used as secondary reference electrode.
 It consists of a thin layer of Pure mercury is placed at the bottom of the container and
it is covered with a paste of mercury- mercurous chloride (Hg+Hg2Cl2) i.e., calomel.
 The remaining portion of the cell is filled with a solution of normal known
concentration of KCl (1M or Saturated) and saturated with Hg2Cl2.
 A platinum wire sealed into a glass tube is dipped into mercury layer is used to
provide the external electrical contact.
 The electrode potential of saturated calomel electrode is 0.2412 V
 The electrode potential of 1M calomel electrode is 0.28 V
Calomel electrode
𝑯𝒈𝟐𝑪𝒍𝟐 𝑯𝒈𝟐
𝟐
+2 𝑪𝒍
−
(Ionisation)
𝑯𝒈𝟐
𝟐
+ 𝟐𝒆 𝟐𝑯𝒈 (Reduction)
𝑯𝒈𝟐𝑪𝒍𝟐 + 𝟐𝒆 𝟐𝑯𝒈 + 2 𝑪𝒍
−
Pt, Hg(l), Hg2Cl2(s)/KCl (xM) Saturated with Hg2Cl2
If it acts as cathode it involves reduction.
If it acts as anode it involves oxidation.
𝟐𝑯𝒈 𝑯𝒈𝟐
𝟐
+𝟐𝒆−
𝑯𝒈𝟐
𝟐
+2 𝑪𝒍
−
𝑯𝒈𝟐𝑪𝒍𝟐
𝟐𝑯𝒈 + 𝟐𝑪𝒍 𝑯𝒈𝟐𝑪𝒍𝟐 + 2 𝒆
−

29. E C.E= E0
C.E –
.
𝑙𝑜𝑔
[ ]
[ ]
E C.E= E0
C.E –
.
𝑙𝑜𝑔
[ ]
E C.E= E0
C.E –
.
𝑙𝑜𝑔 𝑐𝑙
− 2
E C.E= E0
C.E –
.
2 𝑙𝑜𝑔 10 𝑐𝑙
−
E C.E= E0
C.E –0.0591 x 𝑙𝑜𝑔 𝑐𝑙
− 2
The potential of calomel electrode is depends on the concentration of Cl- ions.
Advantages:
1. It is easy to construct and easy to carry.
2. It provides almost a constant potential value with varying temperature and finds application in
laboratories for measuring potential of electrodes.
3. It is used in corrosion studies.
Dr. Suresh Siliveri

30. Dr. Suresh Siliveri
Measurement of Zn electrode potential using SCE
V
V
1.0 V

31. Dr. Suresh Siliveri
The quinhydrone electrode is a type of redox electrode which can be used to measure the hydrogen ion
concentration (pH) of a solution in a chemical experiment. It provides an alternative to the commonly used glass
electrode in a pH-meter.
The quinhydrone electrode consists from a platinum dips into a solution saturated with quinhydrone.
Quinhydrone (HQ) is a slightly soluble compound formed by the combination of one mole of quinone (Q) and
one mole of hydroquinone (H2Q).
QUINHYDRONE ELETRODE
The electrode reaction is:
𝐸 = 0.699 𝑉
Q+2H++2e- QH2+ 0.699V
(Quinhydrone)
(Hydroquinone) (Quinone)

32. Dr. Suresh Siliveri
𝐸 = 𝐸 −
0.0591
n
log
𝑄𝐻2
𝑄 𝐻
+ 2
𝐸 = 𝐸 −
0.0591
2
log
1
𝐻
+ 2
Hydroquinone and quinone taken as equimolar, then [Q] = [QH2]
𝐸 = 𝐸 − 0.059 𝑝𝐻
𝐸 = 0.699 − 0.059 𝑝𝐻
This electrode can be employed for measuring pH of the solution as an indicator electrode and SCE as
the reference electrode
MEASUREMENT OF pH USING QUINHYDRONE ELETRODE
𝐸 = 𝐸 −
0.0591
n
log
𝑃
𝑅
𝐸 = 𝐸 −
0.0591
2
−2 log[𝐻
+
]
𝐸 = 𝐸 + 0.059 log[𝐻
+
]

33. Dr. Suresh Siliveri
 Quinhydrone electrode can very easily be set up by adding a pinch of quinhydrone powder to the experimental solution with
stirring, until the solution is saturated and a slight excess of it remains undissolved. Then, indicator electrode (Pt) is inserted in
it.
 For determining pH value, this half-cell is combined with any other reference electrode, usually saturated calomel electrode and
the e.m.f., of the cell is determined potentiometrically.
The complete cell may be represented as:
Pt,Hg (l) | Hg2Cl2(s) | KCI(satd.) || H+
(unknown) | Q | H2Q |Pt
Calomel electrode Quinhydrone electrode
MEASUREMENT OF pH USING QUINHYDRONE ELETRODE
Calculations:
𝑪𝒆𝒍𝒍 Q 𝑺C𝑬
𝑪𝒆𝒍𝒍 (0.6996 − 0.0591 𝑝𝐻) 0.2415
𝑝𝐻 =
0.6996 − 0.2415 − 𝑬𝑪𝒆𝒍𝒍
.
𝒑𝑯 =
0.4581 − 𝑬𝑪𝒆𝒍𝒍
𝟎.𝟎𝟓𝟗𝟏

34. Dr. Suresh Siliveri
(1) Quinhydrone electrode is easily set up by simply immersing a platinum strip in the test
solution.
(2) The pH values are very accurate even in the presence of oxidizing ions which interfere with
the working of a hydrogen electrode.
(3) It does not give satisfactory results for solutions whose pH is more than 8.5 due to the
ionization or oxidation of hydroquinone.
Merits and demerits of Quinhydrone electrode

35. Dr. Suresh Siliveri
Glass electrode
It is an example for Ion selective electrode.
Ion selective electrode:
Ion-selective electrode possesses the ability to respond only to certain specific ions, thereby
developing a potential with respect to that specific ions in a solution and ignoring the other ions
totally. In other words, the potential developed by an ion–selective electrode depends only on the
concentration of ions of interest.
 Glass electrode is made up of a special type of glass having low melting point and high
electrical conductance.
 The bottom of the glass electrode is blown in the form a bulb, which is filled with a 0.1 N
HCl solution to provide constant hydrogen ion concentration.
 A silver wire coated with AgCl is inserted into the bulb for making electrical contact.

36. Dr. Suresh Siliveri
The representation of glass electrode is, Ag/AgCl/0.1N HCl/Glass membrane.
Glass membranes
• Almost all commercial pH-sensitive glasses in glass membrane electrodes respond to single-
charged ions such as H+, Na+, and Ag+ ions. pH-electrode is most common in that class.
• Few chalcogenide glass membranes are sensitive to double-charged ions such as Pb2+, and
Cd2+ ions.
• Most commercial glass membranes are manufactured based on a tetrahedral network of
silicon dioxide (SiO2) by adding oxides of sodium, potassium, lithium, aluminum, boron,
or calcium.
• (Composition: 72% SiO2, 6% CaO, and 22% Na2O)
𝐸 = 𝐸 + 0.0591 𝑝𝐻
𝐸 = 𝐸 − 0.0591 log[𝐻
+
]

37. Dr. Suresh Siliveri
 For determining pH value, this half-cell is combined with any other reference electrode, usually saturated
calomel electrode and the e.m.f., of the cell is determined potentiometrically.
The complete cell may be represented as:
Ag, AgCl | HCl | Glass membrane | H+
(unknown) || KCI(satd.) |Hg2Cl2(s), Hg(l), Pt,
Calomel electrode
Glass electrode
MEASUREMENT OF pH USING GLASS ELETRODE
Calculations:
𝑪𝒆𝒍𝒍 Cathode anode
𝑪𝒆𝒍𝒍 SCE g
𝑪𝒆𝒍𝒍 SCE −(𝑬𝒈
𝒐
+ 𝟎. 𝟎𝟓𝟗𝟏 𝒑𝑯)
𝑪𝒆𝒍𝒍 = SCE −𝑬𝒈
𝒐 − 𝟎. 𝟎𝟓𝟗𝟏 𝒑𝑯)
𝒑𝑯 = SCE 𝑬𝒈
𝒐− 𝑬𝑪𝒆𝒍𝒍
𝟎.𝟎𝟓𝟗𝟏

38. Dr. Suresh Siliveri
Advantages
• Easy to operate.
• Equilibrium attains very rapidly.
• Results are accurate.
Limitations:
• Glass membrane is fragile.
• Fluoride ions in the sample may attack the glass surface and alter the composition of the
membrane.

39. Dr. Suresh Siliveri
Batteries
 Primary
 Secondary
 Fuel cells

40. Dr. Suresh Siliveri
Battery:
Battery is a device consisting of one or more electrochemical cells connected parallelly or in
series that converts stored chemical energy into electrical energy.
What are primary and secondary batteries?
Primary batteries Secondary batteries (Rechargeable batteries)
 Batteries which are not rechargeable
after their use are called primary
batteries.
 The chemical reactions that take place
are not reversible.
Eg: Lithium Cells Battery
 The batteries which can be recharged
are called Secondary batteries.
 The chemical reaction that take place
are reversible.
Eg: Lead-Acid battery, Lithium ion battery.

41. Dr. Suresh Siliveri
Lithium Cell Battery
 Anode is composed of lithium
 Cathode is composed of heat treated MnO2
 Electrolyte contains a mixture of LiCl, LiBr, LiAlO4, and LiClO4 dissolved in organic
solvents like propylene carbonate and 2-dimethoxyethane.
 The Representation
At anode: Li Li++ e-
At Cathode: MnO2+ Li++ e- LiMn(III)O2
Net Reaction: Li + MnO2 LiMn(III)O2
The battery offers EMF of 3.0 V.
These batteries have following characteristics:
1. Light weight and compact.
2. Low maintenance and high energy density.
These batteries found to be in memory backups, automatic cameras, and calculators.

42. Dr. Suresh Siliveri
 One electrode is Pb, the other electrode is made of lead oxide (PbO2).
 Number of Pb plates are connected in series (anode, -ve plates) and
number of PbO2 plates (cathode, +ve plates) also connected in parallel.
 The Pb plates fits in between the PbO2 plates.
 These plates are separated by insulators (wood, fiber, glass,
rubber).
 The entire combination is immersed in 20-21% dil H2SO4.
 The lead-acid battery was invented in 1859 by French physicist Gaston
Planté. It is the oldest rechargeable battery and It is a secondary battery.
LEAD-ACID BATTERY

43. Dr. Suresh Siliveri
𝑷𝒃 + 𝑺𝑶𝟒
𝟐
→ 𝑷𝒃𝑺𝑶𝟒 ↓ +𝟐𝒆
At Cathode: (Positive Terminal)
𝑷𝒃𝑶𝟐 + 𝑺𝑶𝟒
𝟐
+ 𝟒𝑯 + 𝟐𝒆 → 𝑷𝒃𝑺𝑶𝟒 ↓ +𝟐𝑯𝟐𝑶
𝑷𝒃 → 𝑷𝒃 𝟐 + 𝟐𝒆
𝑷𝒃 𝟐 + 𝑺𝑶𝟒
𝟐
→ 𝑷𝒃𝑺𝑶𝟒 ↓
𝑷𝒃𝑶𝟐 + 𝟒𝑯 + 𝟐𝒆 → 𝑷𝒃 𝟐 + 𝟐𝑯𝟐𝑶
𝑷𝒃 𝟐 + 𝑺𝑶𝟒
𝟐
→ 𝑷𝒃𝑺𝑶𝟒 ↓
At Anode: (Negative Terminal)
Net reaction:
𝑷𝒃 + 𝑷𝒃𝑶𝟐 + 𝟐𝑺𝑶𝟒
𝟐
+ 𝟒𝑯 → 𝟐𝑷𝒃𝑺𝑶𝟒 ↓ +𝟐𝑯𝟐𝑶 + 𝑬
Working
(During Discharging)
𝑷𝒃𝑺𝑶𝟒 → 𝑷𝒃 𝟐 + 𝑺𝑶𝟒
𝟐
𝑷𝒃 𝟐 + 𝟐𝒆 → 𝑷𝒃
𝑷𝒃𝑺𝑶𝟒 + 𝟐𝒆 → 𝑷𝒃 + 𝑺𝑶𝟒
𝟐
At Anode: (Positive Terminal)
(During Charging)
At Cathode: (Negative Terminal)
𝑷𝒃 𝟐 + 𝟐𝑯𝟐𝑶 → 𝑷𝒃𝑶𝟐 + 𝟒𝑯 + 𝟐𝒆
𝑷𝒃𝑺𝑶𝟒 → 𝑷𝒃 𝟐 + 𝑺𝑶𝟒
𝟐
𝑷𝒃𝑺𝑶𝟒 + 𝟐𝑯𝟐𝑶 → 𝑷𝒃𝑶𝟐 + 𝑺𝑶𝟒
𝟐
+ 𝟒𝑯 + 𝟐𝒆
During discharging, Lead-Acid battery acts as electrochemical cell
(Voltaic cell).
During charging, Lead-Acid battery acts as electrolytic cell.

44. Dr. Suresh Siliveri
 During discharging H2SO4 is consumed and its concentration decreases.
 During charging H2SO4 is regenerated and its original concentration is restored. The
variation of concentration of the acid and hence the extent of discharge or charge of the cell
can be easily monitored by changes in specific gravity of the acid.
 During charging, Specific gravity of H2SO4 increases whereas during discharge specific
gravity decreases.
 Voltage of each cell is 2V. In general, lead-Acid battery consists of such 6 cells which are
connected in
 series to get higher voltage(12V).
Applications:
Lead acid -batteries are used in telecommunication, power systems, radio, and television
systems, solar, UPS, electric vehicles, automobile, emergency lights.
Advantages: Low maintenance, Low Cost,
Disadvantages: Heavy in weight, Lead is not environmentally friendly.

45. Dr. Suresh Siliveri
Lithium – Ion Battery (LIB)
• Lithium ion battery is a secondary battery. Lithium ion batteries are rechargeable batteries.
The primary functional components of a Li-ion battery are,
Anode (-ve): Porous Graphite
Cathode (+ve): Lithium metal oxide (LiMO2)
Electrolyte: A non-aqueous medium used as electrolyte, usually
which is a mixture of organic carbonate (ethylene carbonate)
containing complex ([LiPF6]) of lithium ions.
Anode and cathode are separated by electrical insulating
separator (diaphragm) which permeable to lithium ions.
During discharging, Li ions travel from Anode to Cathode & transition metal get reduced from M+4, to M+3
At Cathode (+): 𝑳𝒊MO2 → 𝑳𝒊+ + 𝒆− +MO2
At Anode (-): C6 + 𝑳𝒊+ + 𝒆−→ 𝑳𝒊𝑪6
-------------------------------------------------------------------------
𝑵𝒆𝒕 𝑹𝒆𝒂𝒄𝒕𝒊𝒐𝒏 ∶ 𝑳𝒊MO2 + 𝑪6 → 𝑳𝒊𝑪6 + MO2
-------------------------------------------------------------------------
At Cathode (-): MO2+ 𝑳𝒊+ + 𝒆−→ 𝑳𝒊MO2
At Anode (+): 𝑳𝒊𝑪6 → 𝑪6 + 𝑳𝒊+ + 𝒆−
----------------------------------------------------------------------------
𝑵𝒆𝒕 𝑹𝒆𝒂𝒄𝒕𝒊𝒐𝒏 ∶ MO2 + 𝑳𝒊𝑪6 → 𝑳𝒊MO2 + 𝑪6
----------------------------------------------------------------------------
During Charging During Disharging

46. Dr. Suresh Siliveri
Voltage: 3.7 V
Depending on the material choices the voltage, capacity, life and safety of a Li-ion battery can
changes.
Advantages:
 Light weight, no self-discharge, compact
Disadvantages:
 Over charging may lead to overheating and battery may explode.
 Over charge and discharge makes this battery irreversible.
Applications:
 As mobile phone batteries, laptop batteries.
 As a battery in electric cars.

47. Dr. Suresh Siliveri
Fuel cells are electrochemical cells consisting of two electrodes and an electrolyte which convert
the chemical energy of chemical reaction between fuel and oxidant directly into electrical
energy.
Fuel Cells
Ordinary Combustion process of fuel is
Fuel Oxygen Combustion
Products
Heat
Fuel Oxygen Oxidation
Products
Electricity
The process of fuel cell
The cell consists of two electrodes made of porous graphite, Teflon, PVC.
They are placed in aqueous concentrated (25-35%) solution of NaOH or KOH.
Fuels: H2, CH3OH, CO2, CH4.
Oxidants: O2, H2O2, O3.
Electrode material: Platinum

48. HYDROGEN-OXYGEN FUEL CELL
❖ It consists essentially of an electrolyte solution
such as 25% KOH solution and two inert porous
electrodes H2 and O2 gases as bubbled through the
anode and cathode compartments respectively.
𝟐(𝒈)
−
𝒂𝒒 𝟐 (𝒍)
−
𝟐(𝒈) 𝟐 (𝒍)
− −
(𝒂𝒒)
𝟐 (𝒈) 𝟐(𝒈) 𝟐 (𝒍)
Cell reactions:
At Anode:
At Cathode:
Net reaction:
❖ The one cell produces an e.m.f. of about 1.23 V.
❖ Usually a large no. of these cells are stacked
together in series to make a battery called
“fuel cell battery”.
Dr. Suresh Siliveri

49. CH3OH-O2 Fuel cell:
• Pure methanol is mixed with steam and fed directly to the anode and O2 is bubbled through the cathode.
• The electrolyte is a polymer and the charge carrier is the hydrogen ion (proton).
• The liquid CH3OH is oxidized in the presence of water at the anode generating CO2, hydrogen ions and the
electrons that travel through the external circuit as the electric output of the fuel cell.
• The hydrogen ions travel through the electrolyte and react with O2 from the air and the electrons from the
external circuit to form water at the anode completing the circuit.
• The electrode reactions are
At Anode: CH3 OH (𝒈) + 𝑯𝟐O(𝒍) → CO𝟐+ 𝟔𝑯 + 𝟔𝒆
At Cathode: 𝑂2(𝒈) + 𝟔𝑯 + 𝟔𝒆 → 3𝑯𝟐O(𝒍)
Net reaction: CH3OH +
𝟑
𝟐
02 → CO2+ 2H2O
Electrodes are made of C paper coated with a finely dispersed Pt catalyst. Electrons generated at the anode
travel through an external circuit providing direct current electric power and return to the cathode.
Dr. Suresh Siliveri

50. Applications:
In 2003 president Bush proposed the H2 fuel Initiative (HFI) which was later implemented in
2005. In 2009 president Obama proposed the development of fuel cell-H2 vehicles.
 Useful as power sources in remote locations, such as space craft, remote weather sections,
large parks, rural locations and in certain military applications.
 The energy conversion is very high (75-82%).
 The product H2O is a drinking water source for astronauts.
 They are used as auxiliary energy source in submarines.
 Noise and thermal pollutions are low.
 The maintenance cost is low for these fuels.
Dr. Suresh Siliveri

51. Solar Cells
 solar cell, also called photovoltaic cell, any device that directly converts the energy of light into electrical
energy through the photovoltaic effect.
 The photovoltaic effect can be defined as being the appearance of a potential difference (voltage) between
two layers of a semiconductor slice in which the conductivities are opposite, or between a semiconductor and
a metal, under the effect of a light stream.
Dr. Suresh Siliveri
photovoltaic effect

52. Advantages of Solar Cell
 No pollution associated with it.
 It lasts for a long time . No maintenance cost.
Dr. Suresh Siliveri
Disadvantages of Solar Cell
 It has high cost of installation.
 It has low efficiency.
 During cloudy day, the energy cannot be produced
and also at night we will not get solar energy.

53. 1. Solar Cell for Transportation: Solar energy is used in cars. This solar power is created by
photovoltaic cells. This electricity is transferred to the storage battery or powers the motor.
2. Solar Cells in Calculators: Solar-powered calculators use photovoltaic cells. These calculators
work with solar energy. The light from sun gives power for the operation of calculators. Solar
calculators work very well in outdoor light
3. Solar Cell Panels: On the rooftop, solar panels are kept. It is used as a solar heater which
heats the water. This water can be used for bathing. Also, another use it helps in generating
power. People can store this energy in the backup battery and can use during power cut
issues. Or people can store this energy and use it to generate electricity in their house and save
money by reducing the electricity bill
4. Solar Cell Advantages: Solar energy is a renewable form of energy. Saves money as it reduces
the electricity bill. Maintaining is simple and affordable so the maintenance cost is also low. It
is one of the best alternatives for non-renewable energy.
Dr. Suresh Siliveri

54. Dr. Suresh Siliveri
Corrosion and its control

55. Dr. Suresh Siliveri
CORROSION AND ITS CONTROL
 Corrosion is a process of formation of the unwanted compound of a pure metal by the chemical
reaction between metallic surface and its environment.
 It is an oxidation process. It causes loss of metal.
Introduction
Example:
Definition: Any process of destruction and consequent loss of a solid metallic material, through an
unwanted chemical and electrochemical attack by its environmental, stating at its surface is called
corrosion. or
Disintegration of a metal by its surrounding chemicals through a chemical or electrochemical reaction on
the surface of the metal is called corrosion.
1. Formation of rust (Fe3O4) on the surface of iron.
2. Formation of green film [CuCO3 + Cu(OH)2] on the surface of copper.

56. Dr. Suresh Siliveri
CONSEQUENCES (EFFECTS) OF CORROSION:
Consequences of corrosion cause a great loss of economy and life. The following harmful
effects are specific.
(a) Due to corrosion, properties of metals such as malleability, ductility and electrical
conductivity are lost.
(b) Due to corrosion efficiency of metal is reduced.
(c) Contamination of product will also take place if the corroded equipment is used.
(d) The replacement of corroded equipment is time consuming; maintenance cost
increases.

57. Dr. Suresh Siliveri
Based on the environment, corrosion is classified into
(i) Dry or Chemical Corrosion
(ii)Wet or Electrochemical Corrosion
CLASSIFICATION OR THEORIES OF CORROSION:
This type of corrosion is due to the direct chemical attack of metal surfaces by the atmospheric
gases such as oxygen, halogen, hydrogen sulphide, sulphur dioxide, nitrogen or anhydrous
inorganic liquid, etc.
DRY or CHEMICAL CORROSION:
Example:
1. Silver materials undergo chemical corrosion by atmospheric H2S gas.
2. Iron metal undergo chemical corrosion.

58. Dr. Suresh Siliveri
1.Corrosion by oxygen or oxidation corrosion
2.Corrosion by other gases
3.Liquid metal corrosion
TYPES OF DRY or CHEMICAL CORROSION:
3. The overall reaction is of oxide ion reacts with the metal ions to form metal oxide film.
2 𝑀 +
𝑛
2
𝑂 → 2 𝑀 + 𝑛𝑂
1. CORROSION BY OXYGEN or OXIDATION CORROSION
Metal + Oxygen → Metal Oxide (Corrosion Product)
Alkali metals (Li, Na, K etc.,) and alkaline earth metals (Mg, Ca, Sn, etc.,) are rapidly oxidized at low
temperature. At high temperature, almost all metals (except Ag, Au and Pt) are oxidized.
Mechanism:
1. Oxidation takes place at the surface of the metal forming metal ions
2𝑀 → 2𝑀 + 2𝑒 −
2. Oxygen is converted to oxide ion (O2-) due to the transfer of electrons from metal.
𝑛
2
𝑂 + 2𝑛𝑒− → 𝑛𝑂

59. Dr. Suresh Siliveri
 When oxidation starts, a thin layer of oxide is formed on the metal surface and the nature of this film decides the
further reaction. If it is
(i) Stable oxide layer:
A Stable natured corrosion product do not posses any cracks and pores. Hence it stops the access of oxygen on underlying
metal. Example: The oxide films on Al, Sn, Pb, Cu, etc., are stable, tightly adhering and impervious in nature.
(ii) Unstable oxide layer:
This is formed on the surface of noble metals such as Ag, Au, Pt. As the metallic state is more stable than oxide, it decomposes
back into the metal and oxygen. Hence, oxidation corrosion is not possible with noble metals.
(iii) Volatile oxide layer:
The oxide layer film volatilizes as soon as it is formed. Hence, always a fresh metal surface is available for further attack. This
causes continuous corrosion. MoO3 is volatile in nature.
(iv) Porous layer:
The layer having pores or cracks. In such a case, the atmospheric oxygen have access to the underlying surface of metal,
through the pores or cracks of the layer, thereby the corrosion continues, till the entire metal is completely converted into its
oxide.
1. CORROSION BY OXYGEN or OXIDATION CORROSION

60. Dr. Suresh Siliveri
 Corrosion by other gases like SO2, CO2, H2S etc. The degree of attack depends on the formation of
protecting or non-protective films on the metal surface.
2. CORROSION BY OTHER GASES
B) If the film formation is non-protective or porous the surface of the whole is gradually destroyed.
E.g.: Dry Cl2 gas attacks on tin (Sn) forming volatile SnCl4 there by leaving fresh surface for further
corrosion.
Example:
A) If the film formation is protective or non-porous
(E.g.: AgCl film resulting from the attack of Cl2 on Ag).
The intensity if corrosion decreases because of the film formed will protect the metal from the further
corrosion.

61. Dr. Suresh Siliveri
The corrosion reaction involves either:
(i) Dissolution of a solid metal by a liquid metal or
(ii) Internal penetration of the liquid metal into the solid metal.
3) Liquid metal corrosion
 This is due to chemical action of flowing liquid metal at high temperatures on solid metal or
alloy. Such corrosion occur in devices used for nuclear power.
Pilling-Bedworth rule:
 If volume of metal oxide formed on the surface of a metal is more than or equal to the volume of metal
from which it is formed, the oxide layer will be protective or non porous.
 If volume of metal oxide formed on the surface of a metal is less than the volume of metal
from which it is formed, the oxide layer will be no-protective or porous.
 For example, specific volume ratio (volume of metal oxide /volume of metal) of W is 3.6, Cr = 2.0, Ni =
1.6. Hence, the rate of corrosion is very less in tungsten.

62. Dr. Suresh Siliveri
WET OR ELECTROCHEMICAL CORROSION
Wet or electro chemical corrosion is common type of corrosion which occurs under wet or
moist
conditions.
It is observed when
(i) a metal is in contact with a conducting liquid , and/or
(ii) dissimilar metals are dipped partially in aqueous corrosive environment.
The Wet or electrochemical corrosion involves:
1. Formation of anodic and cathodic areas
2. Presence of conducting medium
3. Corrosion of anodic parts only
4. Formation of Corrosion product closer to the cathodic area.

63. Dr. Suresh Siliveri
Electrochemical theory of corrosion:
According to this theory, when a metal comes in contact with a conducting liquid a galvanic cell is formed
within the metal. Some parts of the metal act as anode and rest act as cathode.
Atmospheric gases and humidity present in the corrosive environment act as an electrolyte.
Corrosion of anodic parts takes place due to oxidation at anode.
Electrochemical theory of corrosion is explained by taking rusting of iron as an example under
different environmental conditions.
The anodic reaction involves oxidation of Fe to Fe+2
The cathodic reaction depends on the availability of oxygen at different pH of the corrosive medium.
The cathodic reaction occurs in two ways
(i) Hydrogen evolution
(ii) Oxygen absorption
Fe Fe+2+ 2e-(oxidation)

64. Dr. Suresh Siliveri
All metals above hydrogen in the electrochemical series have a tendency to get dissolved
in acidic solution with simultaneous evolution of hydrogen.
It occurs in acidic environment. Consider the example of iron,
At anode: Fe → Fe2+ + 2e-
These electrons flow through the metal, from anode to cathode,
where H+ ions of acidic solution are eliminated as hydrogen gas.
At cathode: 2 H+ + 2 e- → H2↑
The overall reaction is: Fe + 2H+ → Fe2+ + H2
Hydrogen Evolution Type:

65. Dr. Suresh Siliveri
At Anode: Metal dissolves as ferrous ions with liberation of electrons.
Fe → Fe2+ + 2e-
At Cathode: The liberated electrons are intercepted by the dissolved oxygen.
1/2O2 +H2O+2e- → 2OH-
The Fe2+ ions and OH- ions diffuse and when they meet, ferrous hydroxide is precipitated.
Fe2+ + 2OH- → Fe(OH)2
(i) If enough oxygen is present, ferrous hydroxide is easily oxidized to ferric hydroxide.
4Fe(OH)2 +O2 +2H2O →4Fe(OH)3 (Yellow rust Fe2O3.H2O)
(ii) If the supply of oxygen is limited, the corrosion product may be even black anhydrous magnetite, Fe3O4.
Oxygen Absorption Type:
The surface of iron is usually coated with a thin film of iron oxide.
However, if this iron oxide film develops some cracks, anodic areas are
created on the surface; while the well metal parts acts as cathodes.
Rusting of iron in neutral aqueous solution of electrolytes (like NaCl solution) in the presence of atmospheric
oxygen is a common example of this type of corrosion.

66. Dr. Suresh Siliveri
Difference between dry and wet corrosion

67. Dr. Suresh Siliveri
Types of electrochemical corrosion- Galvanic Corrosion, Differential aeration corrosion,
Water line corrosion and pitting Corrosion.
Galvanic corrosion (Differential metal corrosion)
It is also called as bi-metallic corrosion. Galvanic corrosion occurs when two dissimilar metals (having
different potentials) are electrically connected and jointly exposed to aqueous corrosive environment. The
metal with the lower potential acts as anode and undergoes corrosion. The metal with higher potential acts
as cathode and is protected.
Example: Iron pipes coupled with copper couplings.
In the above example, iron (-0.44V) acts as anode and undergoes corrosion because its potential is less than
that of copper (0.34V).
If anodic region is small it leads to severe corrosion.
Galvanic corrosion is controlled by
(i) using a pure metal
(ii) using metal closer in galvanic series
(iii) Making sure that anodic metal should occupy large surface.

68. Dr. Suresh Siliveri
Differential aeration (oxygen corrosion cell) corrosion:
 It is a type of electrochemical corrosion that occurs when a metal in contact with conducting
medium (corrosive medium) is exposed to uneven supply of oxygen on its surface.
 As a result of uneven supply of oxygen, an oxygen concentration cell is formed within the
metal.
 The part of the metal which is exposed to less oxygen concentration acts as ‘anodic area’
and gets corroded; while the part of the metal which is exposed to relatively high oxygen
concentration acts as cathodic area of corrosion cell is protected.
 At anode (poorly oxygenated): M Mn++ne-
 At cathode (more oxygenated): 2H2O + O2 + 4e- 4OH-
 Examples of this type of corrosion include drop corrosion, crevice corrosion, waterline
corrosion, pitting corrosion.
 Also the areas below nuts, bolts, joints, areas covered by dust particles and welded areas
are suffered from this type of corrosion.

69. Dr. Suresh Siliveri
Waterline corrosion:
 Waterline corrosion a type of differential aeration corrosion.
 It occurs when a metal is partly submerged in water or a metallic tank is partially filled with water.
 The Part of metal below waterline is poorly oxygenated and acts as anodic area; while the part of the
metal above the waterline is more oxygenated and acts cathodic area.
 Corrosion occurs in the anodic area and simultaneously reduction of oxygen to OH- ions occurs at
cathodic area. The corrosion product is formed closer to cathodic area. This type of corrosion occurs in
water tanks, ocean liners, etc.,.

70. Dr. Suresh Siliveri
Pitting Corrosion:
 Pitting corrosion is a localized form of corrosion resulting in the formation of pin holes or pits.
 When a small cracking, breaking of protective film on the metal surface occurs, small anodic areas
and large cathodic areas are formed.
• The inner area of it is less aerated and the outer surface of the metal is well aerated.
 Therefore, the inner area of the pit acts as anode and under corrosion, while outer surface acts
cathode and protected.
 Pitting may also be due to the galvanic corrosion;
(i) The preferential attack of the environment on a metal of an alloy may lead to pitting. E.g.;
Alloy of cu and Sn
(ii) The scratches on cathodic coatings also favor pitting.

71. Dr. Suresh Siliveri
Factors affecting rate of Corrosion:
1. Nature of metal
2. Nature of the corrosion Product
3. Nature of corroding Environment.
1. Nature of metal
(i) Position of the metal in Galvanic Series: Galvanic series gives real and useful information
regarding the corrosion behavior of metals and alloys in a given environment.
When two metals are electrically combined and exposed to corrosive environment the metal
with lower potential acts as anode and undergoes corrosion where as the metal with higher
potential acts as cathode.
(ii) Purity of metal: In general higher the purity of the metal, lesser the rate of corrosion. Thus
zinc of 99.95 % purity undergoes corrosion at rate of about 5000 times more compared to
zinc of 99.999% purity.

72. Dr. Suresh Siliveri
(iii) Hydrogen Overvoltage: If the hydrogen over voltage of the metal is low, the rate of
corrosion will be high. (there is a difference in potential of the electrode in which the gas evolution of
practical value and with the theoretical value is called hydrogen over voltage)
(iv) Relative areas of anode and cathode: The rate of corrosion is more with the combination
of a large cathodic area and small anodic area.
Therefore , Rate of Corrosion
(v) Nature of surface oxide film
If the corrosion product formed is stable, non-porous and strongly supported layer it prevents
further corrosion.
If corrosion product formed is highly unstable the metals do not undergo corrosion.
If the corrosion product formed is porous or volatile metals undergo severe corrosion.
If the corrosion product is soluble in medium to which it is immersed, corrosion of the metal
occurs with a faster rate.

73. 2) Nature of the corrosion product
(a) Solubility of corrosion products:
In electrochemical corrosion, if the corrosion product is soluble in the corroding medium,
then corrosion proceeds at a faster rate. If the corrosion product is insoluble in the medium
or it interacts with the medium to form another insoluble product (e.g. PbSO4 formation in
case of Pb in H2SO4 medium), then the corrosion product functions as physical barrier,
there by decreases the rate of corrosion.
(b)Volatility of corrosion product :
If the corrosion product is volatile, it volatilizes as soon as it is formed, there by leaving
the underlying metal surface exposed for further attack. This causes rapid and continuous
corrosion, leading to excessive corrosion. Ex: molybdenum oxide (MoO3), the oxdation
product of "Mo" is volatile.
Dr. Suresh Siliveri

74. Dr. Suresh Siliveri
(3) Nature of the environment
(i) Effect of humidity:
Higher the humidity (moisture content) higher will be the rate of corrosion.
Moisture acts as solvent for O2, H2S, SO2 etc., to furnish the electrolyte essential for setting
corrosion cell.
Gases like H2S and SO2 increase the acidity of the medium by their dissolution in water.
Therefore, the rate of corrosion increases as humidity increases
(ii) Effect of temperature: The rate of corrosion increases with an increase in temperature.
(iii) Effect of pH: The rate of corrosion of metals is much faster in acidic pH than in
alkaline or neutral pH, but amphoteric metals like Al, Zn undergoes severe corrosion even
in basic medium.

75. Dr. Suresh Siliveri
Corrosion control methods:
Cathodic protection:
The principle involved in this method is to force the metal to be protected (Parent metal) to
behave like cathode. Therefore, corrosion of the parent metal is prevented. There are two types
of cathodic protections,
(a) Sacrificial anodic protection
(b) Impressed current cathodic protection.

76. Dr. Suresh Siliveri
a) Sacrificial anodic protection
 In this method, the metallic structure to be protected is electrically connected to a more active or anodic
metal than the metallic structure to be protected.
 The more active metal acts as anode and gets corroded slowly; while the parent structure (metallic
structure to be protected) is forced to act as cathode of galvanic cell, hence protected.
 As this more active metal is sacrificed its life in the process of saving metallic structure from corrosion, it
is known as sacrificial anode and, therefore, this method is called as sacrificial anodic protection.
 The metals which are commonly used as sacrificial anodes are Mg, Zn, Al and their alloys.
 This method is used to protect buried pipelines, underground cables, marine structures, ship hulls and
 metallic water tanks.

77. Dr. Suresh Siliveri
b) Impressed current cathodic protection
 In this method, an impressed D.C current is applied, through D.C current source by a
battery, between the metallic structure to be protected which is forced to act as cathode and
an insoluble electrode like platinum, graphite or nickel which is buried in conducting
medium adjacent to the metallic structure to be protected and acts as anode of an
electrolytic cell as shown in the figure.
 The impressed D.C current applied nullifies the corrosion current thus the metallic structure
is protected from corrosion.

78. Dr. Suresh Siliveri
 The metals which are commonly used as insoluble anodes are Graphite, Scrap iron, etc.
 This type of cathodic protection is applied to buried structures such as tanks and pipelines,
since, their operating and maintenance costs are less, and they are well suited for large
structures and long term operations.

79. Dr. Suresh Siliveri
SURFACE COATINGS:
METTALIC COATINGS AND METHODS OF APPLICATIONS
• Protecting the surface of an object by the application of coating is probably the oldest of the
common procedures for corrosion prevention. A coated -surface isolates the underlying metal
from the corroding environment.
• The only limitations of this method is the service behavior of the protective coatings.
• The coating applied must be chemically inert to the environment under particular conditions
of temperatures and pressure. Moreover, coatings must prevent the penetration of the
environment to the material which they protect.
• A brief description of two important protective coatings (metallic coatings) is given below.
(i) anodic coatings
(ii) cathodic coatings:

80. Dr. Suresh Siliveri
Anodic Coatings:
 These are produced from coating metals which are anodic to the base metal (i.e. which is to
be protected).
 Coating of Zn on steel are anodic, because their electrode potentials are lower than that of
the base metal iron.
 If any pores, breaks and discontinuities occur in such an anodic coating, a galvanic cell is
formed between the coating metal and the exposed part of the base metal.
Ex; In case of galvanized steel Zn, the coating metal (Zn) being anodic is attacked, leaving the
underlying cathodic metal (iron) un attacked.
 Galvanic Cell is formed between Zn and the exposed Iron. Zn being anodic to Iron dissolves
anodically, whereas the Iron being cathodic is protected.
 Thus no attack on the iron occurs. Until particularly all the "Zn" has first corroded in the
vicinity of the exposed iron spot. So 'Zn' coating protects iron 'Sacrificially'.

81. Dr. Suresh Siliveri
(ii) Cathodic Coatings ;
• These are obtained by coating a more noble metal (i.e. having higher electrode potential) than the base
metal.
• They protect the base metal, by their higher corrosion resistance than the coat metal.
• Cathodic coating provides effective protection to the base metal only when they are completely
continuous and free from pores, breaks or discontinuities. If such coatings are punctured, much more
corrosion damages can be done to the base metal than to the coating metal.
Ex: A Tin (Sn) coating on a sheet of Iron provides protection only as long as the surface of the metal is
completely covered, since tin is lower than iron in electrochemical series.
• However if the surface coating is punctured, then tin becomes the cathode; while the exposed iron
(which is above tin in the electrochemical series) acts as anode.
• A galvanic cell is set up and an intense localized attack act the small exposed parts occurs, resulting in
severe corrosion of the base metal, iron, such combination of a small anode and large cathode area is
always very dangerous.

82. Dr. Suresh Siliveri
Hot dipping galvanization or galvanizing:
“The process of coating molten zinc on the base metal (Iron) surface by hot dipping is known as
galvanization”.
 In the process of galvanization, the base metal iron / steel is coated with a thin film of zinc metal.
 Since zinc is higher up in the galvanic series the Zn coating acts as anodic with respect to the base
metal (cathodic).
 With repeated exposure of the galvanized metal to environment, only the coated active zinc gets
corroded. Thus, the base metal iron / steel is protected.

83. Dr. Suresh Siliveri
The galvanization process is carried as follows;
1. The metal is cleaned and degreased using organic solvents.
2. The metal is treated with dilute H2SO4 (pickling process) for 10 minutes at about 60-90oC to
remove any rust or scales, if there is any.
3. The metal is further treated with flux materials, ZnCl2 and NH4Cl for the best adhesion property.
4. Finally, the metal is dipped in hot molten zinc at 425 -430oC
5. The excess Zn is removed from the surface of the coated metal by rolling, wiping or air blow
techniques.
Uses: It is most widely used for protection of iron from atmospheric corrosion in the form of roofing
sheets, wires, pipes, nails, bolts, screws, buckets, tubes, etc.
It may be pointed here that 'Zn' gets dissolved in dilute acids to form highly toxic (or poisonous)
compounds. Hence, galvanized utensils cannot be used for preparing and storing food stuffs,
especially acidic ones.

84. Dr. Suresh Siliveri
Tinning:
 It is a coating of tin over the iron or steel articles.
 This process consists of first treating steel sheet in dilute H2SO4 acid to remove any oxide
film. After this it is passed through a bath of ZnCl2 flux. The flux helps the molten tin to
adhere to the metal sheet. Next, the sheet passes through a tank of molten tin and finally
through a series of rollers from underneath the hot tin coated surface of a layer of palm oil.
 The palm oil protects the hot tin coated surface against oxidation.
 The rollers remove any excess of "Sn" and produce a thin film of uniform thickness on the
steel sheet.

85. Dr. Suresh Siliveri
Uses:
 Tin possesses considerable resistance against atmospheric corrosion, moreover, because of
nontoxic nature of tin, tinning is widely used for coating steel, copper and brass sheets,
used for manufacturing containers for storing food stuffs, ghee, oils, kerosene and packing
food materials.
 Tinned copper sheets are employed for making cooking utensils and refrigeration
equipment.

86. Dr. Suresh Siliveri
 Electroplating is the application of electrolytic cells in which a thin layer of metal is deposited onto
an electrically conductive surface.
 A simple example of the electroplating process is the electroplating of copper in which the metal to
be plated (copper) is used as the anode and the electrolyte solution contains the ion of the metal to
be plated (Cu2+ in this example).
 Copper goes into solution at the anode as it is plated at the cathode.
 A constant concentration of Cu2+ is maintained in the electrolyte solution surrounding the
electrodes
Electroplating
anode: Cu(s) Cu2+(aq) + 2e-
cathode: Cu2+(aq) + 2 e- Cu(s)

87.  The method of deposition of a metal from its salt solution on a catalytically active surface by a suitable
reducing agent without using electrical energy is called electroless plating.
 This process is also called chemical plating or autocatalytic plating.
 The metallic ions (M+) are reduced to the metal with the help of reducing agents (R-1). When the
metal(M) is formed, it gets plated over a catalytic surface.
Electro less plating (Ni platting)
Metal ions + Reducing agent metal + oxidized product(S)
Dr. Suresh Siliveri
(Cathode ) Ni2++2e- Ni
(Anode) H2PO2
- +H2O H2PO3
-+2H+ +2e-
Net redox reaction Ni2+ + H2PO2
-+H2O Ni+ H2PO3
-+2H+
a) Coating solution: NiCl2 solution (20 g/L)
b) Reducing agent: Sodium hypo phosphate (20 g/L)
c) Buffer: Sodium acetate (10 g/L)
d) Complexing agent: Sodium succinate (15 g/L)
e) Optimum PH : 4-6.
f) Optimum temp: 85-95oC.

88. Dr. Suresh Siliveri
Advantages of Electroless plating:
• Electrical energy is not required.
• Even intricate parts ( of irregular shapes) can be plated uniformly
• There is flexibility in plating volume and thickness.
• The process can plate recesses and blind holes with stable thickness.
• Chemical replenishment can be monitored automatically.
• Bright finishes can be obtained.
• Plating on articles made of insulators (like plastics) and semiconductors can easily be carried out.
• Electroless plated Ni objects has better corrosion resistance, deposits are pore free, hard and wear
resistant.
Applications:
 They are used in electronic industry for fabricating printed circuits and diodes.
 It is used in domestic as well as automotive fields (eg. jewellery, tops of perfume bottles).
 Its polymers are used in decorative and functional works.
 Its plastic cabinets are used in digital as well as electronic instruments.

89. Dr. Suresh Siliveri
Disadvantages
1. Life span of chemicals is limited
2. Waste treatment cost is high due to the speedy chemical renewal.
Advantages
1. Does not use electrical power.
2. Even coating on parts surface can be achieved.
3. No sophisticated jigs or racks are required.
4. There is flexibility in plating volume and thickness.
5. The process can plate recesses and blind holes with stable thickness.
6. Chemical replenishment can be monitored automatically.
7. Complex filtration method is not required
8. Matte, Semi Bright or Bright finishes can be obtained.

90. Dr. Suresh Siliveri