B.M.S.

College Engineering, Bengaluru 560 019

(Autonomous institute, affiliated to VTU)
Department of Chemistry

Unit 1: Electrochemistry - Electrode Systems and Corrosion 08 hours

Electrodes and Cells – Introduction- Classification of cells - primary, secondary and concentration
cells; Reference electrodes - Calomel electrode; Ion-selective electrodes - Glass electrode.
Determination of pH using glass electrode, numerical on concentration cells.

Corrosion - Definition, Electrochemical theory of corrosion, Types of corrosion - differential metal,

differential aeration and stress corrosion; Factors affecting the rate of corrosion; Corrosion
Penetration Rate (CPR), numerical. Corrosion control: Cathodic protection – Sacrificial anode,
Impressed current method. Metal finishing - Introduction, technological importance; Electroless
plating: Introduction, Electroless plating of copper (PCB). Inorganic coatings – anodizing and
phosphating.

Self-study: Galvanic series and its importance, Electroplating of chromium

Electrodes and Cells

Introduction
Electrochemistry is a branch of chemistry dealing with the interconversion of electrical and
chemical energy. Principles of electrochemistry are vital in understanding of batteries, corrosion,
metallurgy, electrolysis, etc.
Origin of electrode potential
When a metal is dipped in a solution containing its own ions, the metal may undergo oxidation
and goes into solution as metal ions or the metal ions from solution may undergo reduction and get
deposited as metal atoms on the metal surface. Consider a metal ‘M’ is dipped in a solution
containing its own ions (Mn+). The above two opposite tendencies will results in equilibrium as
follows,
M n+ + ne M
When a metal undergoes oxidation, it loses positive ions into solution leaving behind a layer of
negative charges on its surfaces (Figure 1b). This layer attracts positive changes and forms an
electric double layer (EDL). When metal ions in solution are reduced to metal atoms and get
deposited over on the metallic surfaces, the metal surface becomes positively charged. The
accumulated positive charge on the metal surface attracts a layer of –ve charges from solution
(Figure 1a) and it also forms an electrical double layer.

1 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Thus, an electrical double layer (EDL) or Helmotz EDL is established at the metal-solution interface.
This resultant potential difference between the metal and the solution is termed as the electrode
potential.
Electrode potential (E) is the potential developed at the inter-phase of the metal and the solution,
when a metal is dipped in a solution containing its own ions.
When the concentrations of all the species involved in a half-cell is unity then the electrode potential
is known as standard electrode potential (E°). According to IUPAC convention, always standard
reduction potentials are called standard electrode potentials.
When a net reaction proceeds in an electrochemical cell, oxidation occurs at the anode and
reduction takes place at the cathode. The cell consists of two half-cells joined together by an
external circuit through which electrons flow and an internal pathway (salt bridge) that allows ions
to migrate between them so as to preserve electroneutrality.
Nernst equation gives a quantitative relationship between the electrode potential and the
concentrations of metal ions are involved.
Nernst equation for the following reaction, M n+ + ne M

is E = E° – ln[ ]
.
E = E° – log
[ ]

and it can be simplified to (at 298 K),

.
E = E° – log[ ]

Where,
E = Electrode potential
E° = Standard electrode potential
n = Number of electrons involved in the reaction
[Mn+] = Concentration of metal ions in solution
R = Universal gas constant
T = Temperature (K)
F = Faraday constant
The potential difference between the two electrodes of a galvanic cell is called the cell potential
(Ecell) and is measured in volts (V). The Ecell is the difference between the electrode potentials
(reduction potentials) of the cathode and anode. It is called as the electromotive force (emf) of the
cell when no current is drawn through the cell. By convention the anode is kept on the left and the

2 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
cathode on the right while representing the galvanic cell. A galvanic cell is generally represented by
putting a vertical line between metal and electrolyte solution and putting a double vertical line
between the two electrolytes connected by a salt bridge. Under this convention the emf of the cell is
positive and is given by,
i.e., Ecell = Eright – Eleft = Ecathode – Eanode
This is shown in the following example,
+
For the cell reaction: Fe(s) + 2Ag (aq) Fe2+(aq) + 2Ag(s)
Half-cell reactions are,
+ 2Ag(s)
Cathode (reduction): 2Ag (aq) + 2e

Anode (oxidation): Fe(s) Fe2+(aq) + 2e

The cell can be represented as,

Fe(s)|Fe2+(aq)||Ag+(aq)|Ag(s)
and we have Ecell = Eright – Eleft = EAg+|Ag – EFe2+|Fe
The potential of individual half-cell cannot be measured. We can measure only the difference
between the two half-cell potentials that gives the emf of the cell.
Numerical problems
1. Calculate the electrode potential when a copper rod is in contact with 0.01 M Cu +2 ions.
2. For the cell Fe|Fe+2(0.01M)||Ag+(0.1M)|Ag, Write cell reactions and calculate EMF of the cell at
298K.
3. A zinc electrode in contact with 0.01M ZnSO4 is coupled with copper electrode in contact with
1M CuSO4 to form a galvanic cell. Calculate Ecell.
4. If [Cu+2] = 0.03M, what [Fe+2] is needed so that Ecell = 0.76 V.
5. A cell constructed by coupling Zn electrode dipped in 0.5M ZnSO4 and a Ni electrode dipped in
0.05M NiSO4. Write cell representation, cell reaction and calculate the EMF of the cell.
Electrochemical cells
The electrochemical cell is a device, which converts chemical energy into electrical energy and vice
versa. The electrochemical cells are classified into two types,
1) Electrolytic cells: These are the electrochemical cells, which are used to convert electrical energy
into chemical energy.
Ex: Lithium-ion cell (while charging), electro-refining, electro-plating, etc
2) Galvanic or Voltaic cells: These are the electrochemical cells, which convert chemical energy
into electrical energy.
Ex: Daniel cell, Dry cell, etc
Galvanic or Voltaic cells are further classified into three types as follows,

3 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
a) Primary cells: These are the cells which serve as a source of energy only as long as the active
chemical species are present in the cell. The cell reactions are irreversible. These are designed for
only single discharge and cannot be charged again.
Ex: Dry Cell, Zn–HgO cell, Zn–Ag2O cell etc.
b) Secondary cells: These cells are chargeable and can be reused. The cell reactions are reversible
and are often called as reversible cells. During discharging the cells acts like voltaic cell converting
chemical energy into electrical energy. During charging the cell acts like electrolytic cell by
converting electric energy into chemical energy.
Ex: Ni–MH cells, Lithium-ion cells etc.
c) Concentration cells: These are the electrochemical cells consisting of same electrodes dipped in
same ionic solution in both the half cells but are different in the concentration of the ions. In
concentration cells, the emf arises due to the change in the concentration of either the electrolytes or
the electrodes. This is in contrast to galvanic cell where the emf arises from the change in the free
energy of the chemical reaction taking place in the cell. However, in a concentration cell, there is no
net chemical reaction. A concentration cell is made up of two half cells having identical electrodes,
except that the concentration of the reactive ions at the two electrodes are different. The half cells
may be joined by a salt bridge.
Ex: Copper concentration cell, Zinc concentration cell, O2 concentration cell, etc
The electrode, which is dipped in less concentrated solution (C1) act as anode and undergoes
oxidation. The electrode, which is dipped in more concentrated solution (C2) act as cathode and
undergoes reduction.

At anode: M(S) → Mn+ (C1) + ne-

At cathode: Mn+ (C2) + ne- → M(S)
Overall: Mn+ (C2) → Mn+ (C1)
Ecell = Ecathode – Eanode
2.303RT 1 2.303RT 1
Ecell = E   log n – [ E  log n ]
nF M (C 2 ) nF M ( C1 )

2.303RT C
Ecell = log 2 where C2 > C1
nF C1

0.0592 C
Ecell = log 2 at 298 K
n C1

4 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Problems
1. A cell contains two hydrogen electrodes. The negative electrode is in contact with a solution of
H+ ions at pH 6. The emf of the cell is 0.118 V at 25 °C. Calculate the pH at the positive
electrode. Ans: pH = 4
2. Calculate the emf of a concentration cell at 25 °C which contains two zinc electrodes dipped in
two solutions of 0.05 M and 0.5 M concentrations. Ans: 0.0296 V
3. A spontaneous galvanic cell, Sn|Snn+(0.024M)||Snn+(0.064M)|Sn develops an EMF of 0.0126 V
at 298K. Calculate the valency of tin.
4. Calculate the EMF of the following concentrations cells at 298K, (a)
Ni(s)|Ni+2(0.01M)||Ni+2(0.1M)|Ni(s), (b) Cu(s)|Cu+2(0.05M)||Cu+2(5M)|Cu(s).
5. EMF of the cell Cd|CdSO4(0.03M)||CdSO4(xM)|Cd is 0.086V at 298K. Find out the value of x.
6. A cell membrane at 37 °C is found to be permeable to Ca2+ but not to anions, and analysis shows
the inside concentration to be 0.1 M and the outside concentration to be 0.001 M in Ca 2+. What
potential difference would have to exist across the membrane for Ca2+ to be in equilibrium at the
stated concentrations?
Types of electrodes
i) Metal-metal ion electrodes
These electrodes consist of a metal dipped in a solution of its own ions.
Ex: Zn/Zn2+, Cu/Cu2+, Ag/Ag+
ii) Metal-metal sparingly soluble salt electrodes
These electrodes consist of a metal in contact with its sparingly soluble salt
Ex: Calomel electrode (Hg/Hg2Cl2/Cl-), Silver/Silver chloride electrode (Ag/AgCl/Cl -)
Lead – lead sulfate electrode (Pb/PbSO4/SO42-)
iii) Gas electrodes
These electrodes consist of a gas bubbling about an inert metal immersed in a solution containing
ions to which gas is reversible.
Ex: H2 electrode (Pt/ H2/H+), Chlorine electrode (Pt/Cl2/Cl-)
iv) Amalgam electrodes
These electrodes are similar to metal/metal ion electrode. Here, instead of metal a metal amalgam is
used. Ex: Lead amalgam electrode (Pb-Hg/Pb2+)
v) Oxidation – Reduction electrodes
An Oxidation – Reduction electrode is the one which the electrode potentials arises from the
presence of Oxidation – Reduction forms of same species in solution. The potential is measured by
using an inert metal like Pt.
Ex: Pt/Fe3+, Fe2+
Pt/Ce4+, Ce3+

5 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Pt/Sn4+, Sn2+
vi) Ion selective electrodes – in which a membrane is in contact with a solution, with which it can
exchange ions
Ex: Glass electrode

Reference electrodes: are the standard electrodes with reference to these, the electrode potentials of
any other electrode can be determined.
Any electrode of know potential cannot be treated as a reference electrode. The criteria for an
electrode to act as a reference electrode are:
a. The potential of such electrode should be known, under the conditions of utility.
b. The potential of such electrode should show minimum variation with respect to temperature. In
other words, the potential gradient with respect to temperature should be minimum.
The reference electrodes can be classified in to two types,
i) Primary reference electrode - Standard hydrogen electrode
ii) Secondary reference electrodes, Ex: Calomel electrode, Ag/AgCl electrode
The potential of all the electrodes are measured with respect to Standard Hydrogen Electrode
(SHE). Hence, SHE is the primary reference electrode.

Secondary Reference Electrodes

Practical application of the SHE is limited due the difficulty in preparing and maintaining the
electrode, primarily due to the requirement for H2(g) in the half-cell. Therefore, one of two common
secondary reference electrodes, saturated calomel electrode (SCE) or silver-silver chloride electrode
(Ag/AgCl) are used frequently.
Calomel Electrode
Calomel electrode is a secondary reference electrode belongs to the category of metal-sparingly
soluble salt electrode. Mercurous chloride (Hg2Cl2) is commonly called as calomel and it is
sparingly soluble in water. Calomel electrode consisting of a glass container having pure mercury at
the bottom, above which a layer of mixture of Hg and Hg2Cl2 is placed with 3/4th of tube is filled
with KCl solution of known concentration. Electrode potential of the cell depends on the
concentration of KCl used. The platinum wire is used for electrical connections. Salt bridge is used
to couple with other half cell.

6 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
 Or
The calomel electrode is represented as (cell representation): Pt, Hg(l), Hg2Cl2(s)|Cl-.
The reversible cell reactions are,
2+
Hg2 + 2e 2Hg(l)
2+
Hg2Cl 2 (s) Hg2 + 2Cl-
Hg2Cl2 + 2e 2Hg(l) + 2Cl-
The electrode reactions are reversible wrt Cl- ions. In other words the potential of calomel electrode
depends on the concentration of chloride ions.
The Nernst equation for calomel electrode is,
2.303RT
E  E  log[Cl  ]2
nF
At 298 K, it can be simplified to, E = E° – 0.0592*log[Cl-]
The potential of calomel electrode at different concentrations of KCl (at 298 K) are given below,
0.1 N KCl E = 0.3338 V
1 N KCl E = 0.282 V
Saturated KCl E = 0.2422 V
Advantages of calomel electrode are, it is simple to construct, the electrode potential is reproducible
and stable.

Ion Selective Electrode (ISE)

These are the electrodes, which responds to specific ion and develops a potential against that ion
while ignoring the other ions present in the solution.
Ion-selective electrodes are membrane-based electrodes that measure a specific ion (e.g., Ca2+) in
solutions. When the membrane of the electrode is in contact with a solution containing the specific
ion, a voltage, dependent on the concentration level of that ion in solution, develops at the
membrane. The ISEs measure for the specific ion concentration directly. The voltage developed
between the sensing and the reference electrodes is a measure of the concentration of the ion being
measured. As the concentration of the ion at the sensing electrodes varies, so does the voltage
measured between two electrodes.
7 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Types of ISEs
1. Glass membrane electrodes. Ex: Glass electrode
2. Polymer membrane electrodes. Ex: K+, Ca2+ and NO3- ion sensitive electrodes
3. Solid state electrodes. Ex: Cl-, F- ion sensitive electrodes
4. Gas sensing electrodes. Ex: SO2, CO2 sensing electrodes
Application of ISEs
 Determination of hardness of water, Calcium in milk (Ca2+ ISE)
 Nitrate monitoring in waste water (NO3- ISE)
 Salinity of sea water (Cl- ISE)
 In the determination of H+ ion concentration (pH electrode or glass electrode)

Glass (pH) electrode

The glass electrode is perhaps the most successful and ubiquitous electrochemical sensor. It
provides information about the activity of H+ ions. The H+ activity in terms of the pH is the most
commonly monitored and recorded parameter for liquid samples.
The glass electrode consists of a glass membrane, which gives it
its name and separates an internal solution (0.1 M HCl) and a
silver/silver chloride electrode from the solution to be studied. The
Ag/AgCl electrode is connected to a special voltmeter, the pH meter.
The pH meter measures the potential difference and its changes
across the glass membrane. The potential difference must be
obtained between two points; one is the electrode contacting the
internal solution. A second point is obtained by connecting to a
reference electrode, immersed in the solution to be studied. The
complete glass electrode with a reference electrode cell is represented as,
Ag-AgCl | HCl (0.1 M) | glass || solution of unknown pH | reference electrode
Working
When the surface of glass is exposed to water, some Si–O- groups become protonated,

Si-O- + H3O+ Si-O-H + H2O

The exchange of proton, (H+) between the solid membrane and the surrounding solution, and the
equilibrium nature of this exchange, is the key principle of H+ sensing. As with any interface
separating two phases between which ionic exchange equilibrium is established, the glass
membrane/solution interface becomes the site of a potential difference. The glass membrane has two
wall/solution interfaces and there is potential buildup on each of them, with opposite polarity. But
the pH inside the bulb is constant, because the internal solution is sealed. Therefore, the inner
surface potential is constant.

8 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
 When a thin-walled glass bulb containing 0.1 M HCl (C 1) solution is immersed in an acidic
solution of different concentration (Cx) a boundary potential, Eb is developed across glass layers of
the glass membrane. This potential Eb arises due to the difference in the H+ ion concentration inside
and outside the glass bulb i.e.
Eb = Ex – E1 = 2.303RT/nF log Cx/ C1
Eb = 2.303RT/nF log Cx - 2.303RT/nF log C1
Where Cx is the concentration of H+ ions of the solution into which the glass membrane is immersed.
The concentration of H+ ions (C1) inside the bulb is constant.
Additional contribution to the potential of the device comes from potentials of the internal
reference electrode (Ag-AgCl), which is also constant. A small asymmetric potential contribution
exists even if C1 = Cx. The changes in the device potential are therefore due entirely to the pH
changes of the outside solution.
The potential of the glass electrode is given by,

EG = Eb + EAg/AgCl + Easymmetric

EG = 2.303RT/nF log Cx - 2.303RT/nF log C1 + EAg/AgCl + Easymmetric

EG = EG° + 2.303RT/nF log Cx

Where EG° = - 2.303RT/nF log C1 + EAg/AgCl + Easymmetric, is a constant.

EG = EG° – 0.0592pH

At 30 °C the potential of the glass membrane changes by about 60 mV for each one unit of pH.
Asymmetric potential: A small potential that exists even when C1 = Cx. It is due to differences in
the strain of inner and outer surfaces of the glass membrane. Therefore, it is necessary to standardize
each glass electrode before every use by a standard buffer of known pH.
In order to determine the pH of a solution the glass electrode is coupled with a saturated calomel
electrode (SCE).
The cell is represented as, Hg|Hg2Cl2|Cl||unknown solution|glass electrode
SCE and glass electrodes are connected to electronic potentiometers to determine the emf of the
cell. By knowing the EGo of the glass and potential of the saturated calomel electrode, the pH of
unknown solution can be calculated as,
Ecell = Ecathode – Eanode
Ecell = EG – ESCE
Ecell = E°G – 0.0592pH – ESCE
pH = EGo – Ecell – ESCE
0.0592

9 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Advantages
 It is simple to operate.
 It is not poisoned easily.
 It can be used in the presence of oxidizing or reducing agents.
 The equilibrium between the two solutions is reached quickly.
 Only a few drops of the liquid are enough to determine the pH.
 It provides accurate results.
Limitations
 Glass electrode cannot be used in presence of fluoride ions as they can physically damage glass.
 The bulb is too fragile, so the glass electrode has to be used with care.
 A common interference while sensing H+ comes from ions of similar size, notably the alkali metals
ions. The interference is somewhat alleviated by using special lithium glass, with sites too small to
fit Na+ or K+.

10 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Corrosion
Introduction
Corrosion is a natural phenomenon in which the metals are oxidized to their oxides or salts
results in loss of material. The metals lose their metallic strength resulting in the damage to
machineries in which they are used. Thus, corrosion causes damage to metals. The estimate of loss
due to corrosion is approximately 2 lakh crores/annum in India alone. Hence, it is necessary to
understand the process of corrosion in detail.
Corrosion is defined as the destruction or deterioration of a metal or its alloy and consequent loss of
metal, caused due to direct chemical or electrochemical reactions with its environment.
Electrochemical theory of corrosion
According to electrochemical theory, corrosion of metals occurs due to the following reasons,
when they are exposed to the environment.
1) Formation of anodic and cathodic areas on metal surface leading a large number of tiny galvanic
cells.
2) At anode the metal undergoes oxidation and electrons are liberated which migrate towards
cathodic region.
3) Oxygen of the atmosphere undergoes reduction at cathodic area in the presence of moisture
forming hydroxyl ions (OH-).
Most of the corrosion can be explained on the basis of electrochemical reactions on the surface of
metal such a type of corrosion is known as wet corrosion.
Electrochemical theory of corrosion (by taking iron as an example)
When a metal like iron is exposed to the corrosive environment according to electrochemical
theory, corrosion of metal takes place due to the formation of anodic and cathodic regions on the
same metal surface in a corrosive medium. The anodes and cathodes are formed due to the
heterogeneities at the interfaces of the metal and environment. The heterogeneities on a metal
surface could develop due to several factors like,
1. On a metal surface, if the concentration of the oxygen is different.
2. Due to contact of two different metals.
3. If metal surface subjected to stress.
Thus, anodic and cathodic area are formed, in corrosive media (like moisture)
At anode oxidation takes place so that the metal is converted into metal ions with the liberation of
electrons.
M → Mn+ + ne-
Ex: when iron is exposed to the corrosive environment it undergoes oxidation as
Fe → Fe2+ + 2e-

11 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
At the cathodic regions, reduction takes place, since the metal at cathodic region cannot be reduced
further, so some constituents of the corrosive medium take part in the cathodic reaction. Since, in the
cathodic reaction as the constituents of the corrosion medium are involved, they are dependent on
the nature of environment. Most common types of cathodic reactions are,
1. Liberation of hydrogen and 2. Absorption of oxygen.

Cathodic reactions: The electrons released at anode migrate to the cathodic area.
a) In acidic medium in the absence of O2: In acidic medium and in the absence of oxygen,
hydrogen ions are reduced to hydrogen gas (H2 liberation).
2H+ + 2e- → H2
b) In neutral or alkaline medium and in the absence of O2: If the solution is alkaline and in the
absence of oxygen the cathodic reaction is (H2 liberation),
2 H2O + 2e- → 2OH- + H2
c) In neutral or alkaline medium in presence of O2: when the solution is neutral and aerated,
hydroxyl ions are formed as follows (O2 consumption),
2 H2O + O2 + 4e- → 4OH-
d) In acidic medium in the presence of O2: In acidic medium and in the presence of oxygen (O2
consumption),
4H+ + O2 + 4e- → 2H2O
Formation of corrosion product: If the medium is neutral or mildly alkaline, the hydroxyl ions
formed migrate towards anode and the metal ions (eg: Fe+2 ions) migrate towards cathode and react
with each other and forms corrosion product. In the case of iron, OH- reacts with Fe2+ ions and forms
ferrous hydroxide (Fe(OH)2).
2Fe2+ + 4OH- → 2Fe(OH)2
Ferrous hydroxide further reacts with O2 and H2O forming hydrated ferric oxide or rust.
4 Fe(OH)2 + O2 + 2 H2O → 2 Fe2O3·3H2O (rust)

12 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
As there are large numbers of tiny galvanic cell, entire surface of iron will be covered with rust over
a period of time. Since rust is porous, air and moisture enters in side and the inner layer of metal also
corrodes slowly.

Types of corrosion
1. Differential metal corrosion (Galvanic corrosion)
This type corrosion occurs when two different metals are in contact with each other. The metal
having less standard reduction potential (SRP) value undergoes oxidation and liberates electrons,
which migrates to the cathode. The other metal having higher SRP value acts as cathode and
reduction reaction takes places on its surface forming OH- ions or any one kind of reduction
reactions. The rate of corrosion depends on the potential difference between the two metals. If the
difference is more corrosion occurs faster and vice
versa. The anodic metal undergoes corrosion and
cathodic metal is unaffected.
The reactions are,
At anode: M → Mn++ ne-
At cathode: Depending on the nature of corrosion
environment the cathodic reactions are as follows,
i) 2H+ + 2e- → H2
ii) 2 H2O + 2e- → 2OH- + H2
iii) 4 H+ + O2 + 4e- → 2H2O
iv) 2H2O + O2 + 4e- → 4 OH-
Eg: Iron metal (anode) in contact with copper metal (cathode), iron pipe in contact with brass tap,
etc.
2. Differential aeration corrosion
This type of corrosion occurs when a metal is exposed to different concentrations of oxygen
(O2). The part of metal which is more exposed to air acts as cathode and unaffected. The other part
of the metal, which is less, exposed to air acts as anode and undergoes corrosion.
Ex: Iron rod partially immersed in water, ship in sea etc.
Water line and pitting corrosion are two typical forms of differential
aeration corrosion.
i) Water line corrosion: It is differential aeration type of corrosion
observed in water storage tanks, ships, etc. During water line corrosion,
the part of the metal below water line is exposed to less oxygen
concentration act as anode and undergoes corrosion than the other part
which is more exposed to atmospheric oxygen which acts as cathode.

13 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
More intense corrosion is observed just below the water line, hence called as water line corrosion.
At anode: Fe → Fe2+ + 2e-
At cathode: 2 H2O + O2 + 4e- → 4 OH-
Then, 2 Fe2+ + 4OH- → 2 Fe(OH)2
2 Fe(OH)2+ O2 + 2H2O → Fe2O3·3H2O rust
ii) Pitting corrosion: Pitting corrosion occurs when small particle like dust, mud etc get deposited
on metals surface. The portion of metal covered by the dust or other particles is less aerated and acts
as anode. The other portion of the metal exposed to more oxygen of the environment act as cathodic
region. Corrosion takes place at the portion below dust and a small pit is formed. Then the rate of
corrosion increases due to small anodic area and large cathodic area.
The anodic and cathodic reactions are same as water line corrosion.

Pitting corrosion is a more dangerous, localized form of corrosion. Because, once pit is formed the
corrosion rate increases due to further decrease in concentration of O2 inside the pit. Hence in a short
period of time a narrow deep pit is formed, results in sudden failure of structure.
3. Stress corrosion
Stress corrosion is seen in metals suffering from stress which may result from mechanical
operations such as design riveting, cold working, welding, bending, pressing, quenching etc. In a
corrosive environment the stressed portion act as anode and undergoes corrosion. The other
unstressed part of the metal acts as cathode. The atoms in the stressed part are slightly displaced
from their original position thus weakening the cohesive force
between surface metal atoms. So, the metal atoms at stressed part are
more reactive and act as anode.
Ex: Corrosion of head and point portions of a nail indicates that they
have been acting as anode to the middle portion. Generally, the head
and the point portions were put under stress during their manufacture.
In the case of iron-wire hammered at the middle, corrosion takes place
at the hammered part and results in breaking of the wire into two pieces. Caustic embrittlement takes
place in stressed parts such as bends, joints and rivets in boilers.
Caustic embrittlement in boilers: It is a form of stress corrosion that takes places in boilers
operating at high pressure between 10 to 20 atm. The stressed portion contains fine hairline cracks.
14 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Boiler water contains of alkali, which enters into the cracks. The water evaporates leaving behind
caustic soda. An electrochemical cell is set up between the iron under stress and iron in the main
body. The iron surrounded by dilute NaOH (main body) act as cathode and iron under stress (crack)
act as anode and gets corroded resulting in failure of boiler.
The Na2CO3 present in boiler water is hydrolyzed at high temperature to give NaOH.
Na2CO3 + H2O → 2 NaOH + CO2
The NaOH formed concentrates on hairline cracks, and set up a concentration cell. The Iron in crack
corrodes to produce Sodium Ferrate, which decomposes to give Fe3O4 and NaOH.
2 NaOH + Fe → Na2FeO2 + H2
3 Na2FeO2 + 4 H2O → Fe3O4 + H2 + 6 NaOH
So, the NaOH is regenerated and reacts with more iron to convert it into Fe3O4

Factors affecting the rate of corrosion

i) Nature of the metal (electrode potential of anode and cathode): The rate of corrosion increases
as the potential difference between anode and cathode increases. In general, lower the reduction
potential higher will be the tendency to get oxidized (corroded).
Ex: The potential difference between copper and iron is 0.78 V. The potential difference between
iron and silver is 1.24 V. In these two cases, iron undergoes fast corrosion when it is in contact with
silver than when it is in contact with copper.
ii) Nature of corrosion product: If the corrosion product forms a protective layer on the metal
surface, it prevents the further corrosion of metal (passivation of metal). If the corrosion product
layer in non-porous, adherent, insoluble, non-conducting and stoichiometric then it act as a perfect
barrier between metal and its environment. Such a layer passivates the metal from further corrosion.
If the corrosion product is porous, non- adherent, soluble and non- stoichiometric then it cannot form
a protective layer and corrosion of metal continuous.
It is observed that the passivation is related specific volume (volume/unit weight) of metal and its
oxide (corrosion product).
Specific volume ratio = Volume of metal oxide / volume of metal
Ex: In oxidizing environment, metals like Al, Cr and Ti (specific volume ratio > 1) form protective
metal oxide films on their surfaces which prevent further corrosion. Metals like Zn, Fe and Mg
(specific volume ratio < 1) do not form protective layer and are readily under goes corrosion.
Therefore, the rate of corrosion depends on the nature of corrosion product.
iii) Anodic to cathodic area ratio: If a metal has small anodic and large cathodic area the rate of
corrosion increases and vice versa. This is because when small anode is contact with large cathode,
the large cathode puts high demand for electron. The electrons liberated during oxidation at anode

15 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
are quickly consumed on large cathodic surface for the reduction reactions and rate of corrosion is
high.
iv) pH: In general, at lower pH value the rate of corrosion is more at higher pH value (pH more than
10) the rate of corrosion ceases due to the formation of protective coating of oxides on the metal.
When, pH < 4 high corrosion.
pH between 4 and 10 the rate corrosion depends on availability of oxygen.
pH > 10 no corrosion.
v) Temperature: Increase in temperature results in an increase in the conductance of the aqueous
medium and rate of corrosion also increases. Increase in temperature also reduces polarization that in
turn increases the rate of corrosion.
vi) Polarization: The variation of electrode potential due to the inadequate supply of species from
the bulk of the solution to the electrode is known as polarization.
Anodic polarization: Increase in anodic polarization, due to accumulation of M n+ ions around anode
decreases the rate of corrosion. Due to the anodic polarization the tendency of oxidation decreases.
Cathodic polarization: Increase in cathodic polarization, due to accumulation of OH- ions around
cathode, decreases the rate of corrosion. The cathodic polarization reduces the tendency for
reduction. Due to inadequate consumption of electrons at cathode, oxidation reaction (corrosion)
cannot continue. Therefore, the rate corrosion decreases.
vii) Hydrogen over voltage
In acidic medium, the competing cathodic reaction is the reduction of H+ to liberate H2 gas.
Hydrogen over voltage is the measure of the tendency of the electrode to liberate H 2 gas. A metal
with low hydrogen over voltage on its surface is more susceptible for corrosion. When the cathodic
reaction is hydrogen evolution type with low hydrogen over voltage, liberation of H 2 gas is easier so
that cathodic reaction is fast, that makes anodic reaction faster, hence overall corrosion process is
fast. If the H2 over voltage is high then cathodic reaction will be slow and hence anodic reaction
(corrosion) will also be slower.
viii) Humidity: Atmospheric corrosion of iron is slow in dry air but increases rapidly in the presence
of moisture. This is due to the fact that moisture acts as the solvent for the oxygen in the air to
furnish the electrolyte essential for setting up a corrosion cell. Rusting of iron increases when the
relative humidity of air reaches from 60 to 80%.

Corrosion control methods

The important methods used to control corrosion are,
1. Design and selection of the materials
2. Protective coatings
a) Organic coatings – paints and enamels

16 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
 b) Inorganic coatings
i) Metal coatings – anodic (galvanization) and cathodic (Tinning)
ii) Surface conversion coatings – anodizing, phosphating
3) Corrosion inhibitors
4) Cathodic protection

Inorganic coatings
Surface conversion coatings
Surface conversion coatings are chemical conversion coatings. The surface layer of the base
metal is converted into a compound by chemical or electrochemical reactions, which prevents the
base metal form corrosion. The coating can be done by chemical dip, spray or by electrolytic
method. The coating helps in the increased electrical insulation, enhanced adherence for paints and
prevention of corrosion.
Ex: anodizing, phosphating
Anodizing
Anodizing is a process of converting outer
layer of base metal in to a protective passive oxide
film. Anodizing can be carried out on the surface
of metals such Al, Zn, Mg, Ti, Zr, Ta, Cr, etc by
electrochemical oxidation.
In anodizing of aluminum, it is cleaned, degreased
and polished then made as anode in an electrolytic cell. It is immersed in an electrolyte consisting 5 -
10% chromic acid. Steel or copper is taken as cathode. The temperature of the bath is maintained at
35 ᵒC. A current density of 100 A/m2 or more is applied which oxidizes outer layer of Al into Al2O3.
Reaction
2 Al + 3 H2O → Al2O3 + 3 H2
Uses: tiffin carriers, soapboxes, household utensils, window frames, etc.
Phosphating
Phosphate coating is a process of converting the surface layer of the base metal into their
phosphates by chemical or electrochemical reaction between base metal and aqueous phosphoric
acid. Phosphate coating is generally obtained on steel, Al and Zn surfaces.
-The metal to be coated (base metal) is degreased, polished washed and dried.
-It is dipped in a solution containing mixture of phosphoric acid, metal phosphates such as Fe, Mn,
Zn phosphates
-Accelerators such as copper salts, ClO3-, nitrates etc. are added to increase the rate of reaction.
-pH of the bath is maintained between 1.8 - 3.2

17 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
-Temperature is maintained at 35 °C
-The metal reacts with phosphoric acid forming metal phosphate
-The metal phosphates deposits on the surface of base metal
Reaction: 3 Fe + 2 H3PO4 → Fe3(PO4)2 + 3 H2
Phosphate coatings are porous and they do not provide appreciable corrosion resistance. They are
mainly used as base coating for paints and enamels, as they increase adhesion quality of paints.
Uses: Phosphate coated metal widely used in automobile industry, bolts, nuts, and refrigerators,
washing machining, cars bodies etc.
Cathodic protection
Cathodic protection is a method in which the base metal to be protected from corrosion is made
to act as cathodic by attaching a more active anodic metal to it. The active anodic metal undergoes
corrosion and base metal is protected from corrosion the following methods are used to protect metal
from corrosion. a) Sacrificial anodic protection method, b) Impressed voltage method
a) Sacrificial anodic protection method
In this method the more active metals like Zn or Mg are attached to base metal (Fe). The anodic
metals being more reactive undergo corrosion but base metal remains unaffected. The sacrificial
anodes have to be replaced from time to time after complete corrosion. The method is used for
protecting buried pipeline, ship hulls, industrial water tank and steel rods in RCC columns. Several
hundred kilometers long zinc wire is buried along oil pipe line in Alaska is an example for sacrificial
anodic protection method.

a) Impressed voltage method

The arrangement for protecting a buried pipeline is illustrated below Figure The buried pipe
(object to be protected) receives current from a DC power source via an auxiliary inert electrode
(like graphite) buried in the ground. The pipe becomes the cathode (connected to –ve terminal of DC
battery) and the auxiliary electrode the anode (connected to +ve terminal of DC battery).

18 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Metal finishing
Metal finishing is a surface process carried out for modifying surface properties of metals by
deposition of a layer of another metal, or an alloy or a polymer film or by the formation of an oxide
film on its surface. e.g.: electroplating of metals, electro-less plating, chemical coating, etc.
The technological importance of metal finishing
Metal finishing imparts desirable surface characteristics such as,
1. Higher corrosion resistance.
2. Improved wear resistance.
3. Providing electrical and thermal conducting surface.
4. Thermal resistance and hardness.
5. Providing optical and thermal reflectivity.
6. In the manufacture of electrical and electronic components such as PCB’s, capacitors contacts etc
7. In electro-framing of articles, electrochemical machining, electro-polishing and electro-chemical
etching.
8. To increase the decorativeness of metal surface.
Electroless plating
Electroless plating is the controlled autocatalytic deposition of a continuous film of a metal from
its salt solution on a catalytically active surface by a suitable reducing agent without using electrical
energy.
Electroless plating is a method of depositing a metal or alloy over a substrate (conductor or non-
conductor) by controlled chemical reduction of the metal ions by a suitable reducing agent without
using electrical energy. The reduction of metal ions by the reducing agent is catalyzed by the metal
atoms being plated. Therefore, electro-less plating is also termed as autocatalytic plating.
The electro less plating process can be represented as,
Metal ions + reducing agent Metal + oxidized product
Advantages of electroless plating
1. Does not require electrical power source
2. It is applicable to conductors, semiconductors and non-conductors like plastics

19 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
3. Electroless plating is less porous than electroplates and possess unique characteristic chemical,
mechanical and magnetic properties
4. Electroless plating baths have better throwing power and deposit a more uniform metal coating
over an article irrespective of its shape or size
Disadvantages of electro-less plating
1. Coatings generally contaminated with oxidized products
2. Costlier process compared electroplating
3. Rate of deposition is slow
Comparison of electro-less and electroplating
Property Electroplating Electroless plating
Driving force Power supply Auto-catalytic redox reaction
Cathodic reaction Mn+ + ne- M Mn+ + ne- M
Anodic reaction M Mn+ + ne- or Red Ox + ne-
n/2 H2O n/4 O2 + nH+ +ne-
Site of cathodic article to be plated article to be plated (catalytic
reaction surface)
Site of anodic separate anode article to be plated
reaction
Anode reactant M or H2O reducing agent
pure metal or definite alloy metal with reducing agent and
oxidized products as impurities
Deposition can’t be made on non- Deposition can be made on
conductors such as plastics, non-conductors such as
Nature of deposit
ceramics etc. plastics, ceramics etc.
Requires levelers Does not require levelers
Plating baths don’t have excellent Plating baths have excellent
throwing power throwing power
Electro-less plating of copper
Pretreatment and activation of the surface: The surface to be coated is first degreased by organic
solvents or alkali followed by acid treatment.
1. Metals like Fe, CO, Ni etc. do not need pretreatment.
2. Non-metallic materials (e.g., glass, plastics, printed circuit boards, PCB) are activated by first
dipping in SnCl2 and HCl solution, followed by dipping in PdCl2 solution and dried.
SnCl2 + PdCl2 SnCl4 + Pd

20 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Composition of bath
Coating solution: CuSO4 solution (12 g/L)
Reducing agent: Formaldehyde (8 g/L)
Buffer: NaOH (15 g/L) + Rochelle salt (14 g/L)
Complexing agent: Di sodium salt of EDTA (20 g/L)
pH: 11.0
Temperature: 25 °C
Reactions
Cu2+ + 2e- Cu
2HCHO + 4OH- 2HCOO- +2H2O + H2 +2e-

Net redox reaction: Cu2+ + 2HCHO + 4OH- Cu +2HCOO- +2H2O + H2

Applications
1. Mainly used in printed circuit boards (PCB) manufacturing
2. For plating on non-conductors
3. Applied on wave-guides and for decorative plating on plastics

Application of electro-less plating in printed circuit board (PCB) manufacture

Preparation of printed circuit boards (PCB): A PCB is made of glass-reinforced rubber (GR-P)
like epoxy or phenolic polymers. A thin layer of copper is electroplated (5 – 100 mm thickness) over
the PCB. Then, selected areas are protected by employing electroplated image and the remaining
part of the plated copper is removed by etching so as to get required type of circuit pattern (or track).
Connection between the two sides of PCB is made by drilling holes, followed by electro-less copper
plating through holes.

Introduction to corrosion rate Corrosion rate is the speed at which any metal in a specific
environment disintegrates. It also can be described as the amount of corrosion loss per year in
thickness typically computed using mils per year (mpy). The rate of deterioration depends on the
environmental conditions and the nature of metal under consideration generally termed as
“Corrosion penetration rate” (CPR).
21 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Corrosion penetration rate - The speed at which any metal in a specific environment deteriorates
due to a chemical reaction (corrosion). It indicates the amount of corrosion loss per year in terms of
thickness and the speed at which corrosion spreads to the inner portions of a material. It is expressed
as mils penetration per year (mpy) or millimeter per year (mmpy) The following are the data
required to estimate the corrosion rate of any metal
1. Weight loss of metal (decrease in weight during a said period)
2. Density of metal
3. Total surface area present initially
4. Length of time taken
Methods to measure corrosion rate
1. Weight loss method- Weight loss of the metal before and after exposure to the corrosive
environment.
2. Microscopic method – The nature of corroded surface (number of pits and their depth) is
determined
3. Measurement of corrosion potential – Corrosion tendency is measured by coupling the test metal
with standard half-cell and determining the corrosion current.
4. Measurement of electrical resistance – As electrical resistance increases with corrosion; a periodic
measurement of resistance can determine the corrosion.
5. Corrosometer – The ratio of resistance of the corroding test metal to the resistance of its non-
corroding standard probe covered with highly corrosion resistant coating is measured
6. Electrochemical tests – The amount of externally applied current required to change the corrosion
current of corroding test metal is measured.
Weight loss method
It is commonly used method for the measurement of uniform corrosion. It involves the exposure of a
weighed piece of test metal or alloy to a specific environment for a precise time. This is
accompanied by thorough cleaning to remove the corrosion products and then determining the
weight of the lost metal as a result of corrosion. The corrosion rate is best described in terms of the
thickness or weight loss where the surface of the metal disintegrates uniformly across the area that
has been exposed. Corrosion penetration rate is calculated using the following equation

22 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)
Summary

23 Dr. AN, Dept. of Chemistry, BMS CE (2023-24 - II, for PC & PR)