MODULE-3 CORROSION

1. Define metallic corrosion. Describe the electrochemical theory of corrosion taking

iron as an example.
Corrosion is defined as “the destruction and consequent loss of metals or alloys through
chemical or electrochemical attack by the surrounding environment”.
Ex: Rusting of iron, green scales are formed on copper vessels
When iron is exposed to air in the presence of moisture, hydrated ferric oxide (rust) is
formed.
Electrochemical theory of corrosion:
According to electrochemical theory, when a metal such as iron is exposed to
corrosive environment, following electrochemical changes occur.
a) Formation of galvanic cells: Anodic and Cathodic areas are formed resulting in a
minute galvanic cells.
b) Anodic reaction: At the anodic area, oxidation takes place resulting in the corrosion
of iron. Fe Fe2+ + 2e-
c) Cathodic reactions: The electrons flows from anodic to cathodic area and cause
reduction of species present in environment around the metal. There are three
possible ways in which the reduction can take place.
 If the environment is aerated and neutral, oxygen is reduced in presence of water to
OH- ions according to the equation.
O2 + 2H2O + 4e- 4OH-
 If the environment is deaerated and neutral, water is reduced to H2 and OH-.
2H2O + 2e- H2 + 2OH-
 If the environment is deaerated and acidic, the H+ ions are reduced to H2.
2H+ + 2e- H2
The metal ions (formed at anode) combine with the OH - ions to form the metal
hydroxide
Fe2++ 2OH- Fe (OH) 2.

Ferrous hydroxide further reacts with O 2 and H2O forming hydrated ferric oxide which
is the familiar brownish-red colored corrosion product (Rust).
2 Fe (OH) 2 + O2 + (n-2) H2O Fe2O3.nH2O
Entire surface of iron will be covered with rust
 Rusting of iron

2.Discuss the following types of Corrosion:

a. Differential Metal Corrosion b. Differential aeration Corrosion
Galvanic corrosion or differential metal corrosion:
 This type of corrosion occurs when two dissimilar metals are in contact with each
other and are exposed to a corrosive environment.
 The two metals differ in their electrode potentials. The metal with lower
electrode potential acts as anode and the other metal with higher electrode
potential acts as cathode.
 The anodic metal undergoes oxidation and gets corroded. A reduction reaction
occurs at the cathodic metal. The cathodic metal does not undergo corrosion.
Example: When iron is in contact with copper, iron has lower electrode potential acts
as anode and undergoes oxidation as, Fe Fe 2+ + 2e-
Copper having higher electrode potential acts as cathode & is unaffected. Depending on
the corrosive environment near the cathode either hydrogen evolved or oxygen
absorbed resulting in hydroxide ion formation.
At cathode O2 + 2 H2O + 4 e- 4 OH-
Fe2 ++ 2OH - Fe (OH)2.
Fe Anode Cu metal
2Fe (OH)2 .+ O2 + (n-2)H2O Fe2O3.nH2O
Anode Cathode
The rate of corrosion depends mainly on
 Potential difference between two metals. Higher the difference, higher is the rate of
corrosion.
 Smaller anodic area and larger cathodic area increases the rate of galvanic corrosion.
Thus iron corrodes faster when it is in contact with copper than with tin.
Differential aeration corrosion:
 This type of corrosion occurs when two different parts of the same metal are
exposed to different oxygen concentrations.
 The part of the metal which is exposed to less oxygen concentration acts as
anode. The part which is exposed to more oxygen concentration acts as cathode.
 The anodic region undergoes corrosion and the cathodic region is unaffected.
Example: When a metal strip of iron is partially
immersed in an aerated solution of NaCl the Iron

concentration of O2 is higher at the top surface More O ,

(Cathode)
2

than inside the solution. Since cathodic reaction

requires oxygen, the cathodic area tends to
concentrate near the water line so that bottom
portion of the specimen acts as anode where Less O , (Anode)
2

corrosion starts.
Wate
At anode: r

Fe Fe2+ + 2e-
At cathode (near water line):
O2 + 2H20 + 4e- 4OH-
Water line corrosion:
This is a case of differential aeration corrosion commonly observed in steel water
tanks, ocean going ships etc. The part of the metal below the water line is exposed only
to dissolved oxygen while the part above the water line is exposed to higher
concentration of atmospheric oxygen.

Thus the metal part below the water line acts as anode and undergoes corrosion.
Whereas the metal part above the waterline, which is more oxygenated acts as cathode
and unaffected. More intense corrosion is observed just below the water line, hence it is
called water line corrosion. This type of corrosion is commonly observed in ships
floating in seawater for a long period of time.
3.Define galvanizing. Describe galvanizing of Iron and mention its applications.
Galvanizing is the process of Coating a layer of zinc on iron by hot dipping process.
Galvanizing of iron is an example of anodic metal coating.
Process: Galvanization involves the following steps
1. Solvent Cleaning: The metal surface is washed with organic solvents to remove
organic impurities on the surface.
2. Alkali Cleaning: Residual organic impurities are removed by treating the object with
alkali such as NaOH.
3. Picking: Rust and Scale is removed by washing the object with dilute sulphuric acid
H2SO4
4. Finally, the article is washed with water and air-dried.
5. Then it is treated with the mixture of aqueous solution of ZnCl 2 and NH4Cl which acts as flux
and then dried.
6.The article is then dipped in a bath of molten zinc at 450˚C (Molten zinc is covered
with a flux of ammonium chloride to prevent the oxidation of molten zinc.)
7. The excess zinc on the surface is removed by passing through a pair of hot rollers.
Application
Galvanization of iron is carried out to produce roofing sheets, fencing wire, buckets,
bolts, nuts, pipes etc.

4.What is anodizing of aluminum? Describe anodizing of Aluminum and mention its

applications.
Anodizing is a type of inorganic coating.
Anodizing is the process of conversion of surface atoms of metal object into its highly
protective metal oxide film through electrolysis.
Eg. Passive metals like aluminum, tantalum, zirconium and titanium etc. which are
capable of forming non-porous, non-conducting oxide layer can be protected from
corrosion by subjecting to anodizing.

Following steps involves in Anodizing of Aluminum

 Aluminum article is degreased, cleaned and polished.
 Article is made it as anode and copper as cathode and placed in an electrolytic cell
containing 5-10% chromic acid as electrolyte.
 Electrolysis is carried out at 350C and suitable voltage is applied to get desired
thickness of the oxide film. Outer aluminum layer is oxidized to form alumina which
gets deposited on the metal.
 Then article is treated with boiling water to convert porous alumina into nonporous
hydrated alumina.
 Then article is dipped in aq. Dye solution at 50 - 600C.
 The Al2O3.nH2O layer on the surface acts as a protective coating, hence corrosion is
prevented.
 Reactions during electrolysis
At anode: 2Al + 3H2O  Al2O3 + 6H+ + 6e
At cathode: 6H+ + 6e  3H2
Overall reactions: 2Al + 3H2O  Al2O3 + 3H2

Application: Anodized articles are used in Tiffin carriers, household utensils, window
frames etc.
5.What is cathodic protection? Describe sacrificial anode technique and mention
the advantages and disadvantages.
Cathodic protection: The process of protecting a metal against corrosion by making the
entire metal as cathode by providing electrons from external source is called as
Cathodic protection.
 A metallic structure is converted to cathode by connecting it to another highly
reactive metal.
 A metallic structure is converted to cathode by supplying electrons DC source.
Sacrificial anode method:
In sacrificial anode method, the metal to be
protected is electrically connected to a more
active metal using insulated copper wire.
For example, when steel is to be protected,
it may be connected to a block of Mg or Zn.
In such a situation, steel acts as cathode
(high electrode potential) and is unaffected
as shown Fig. Mg and Zn act as anode (low
electrode potential) and undergo sacrificial
corrosion. When the sacrificial anode gets
exhausted, it is replaced with new ones.
Ex: 1. Al, Mg or Zn block connected to a buried
oil storage tanks or pipe lines.
2. Mg bars are connected to ocean going ships
Advantages: Sacrificial anode methods are
simple with low installation cost and do not require power supply.
Disadvantages; When the sacrificial anode gets exhausted, it must be replaced with
new ones otherwise specimen undergoes corrosion without inhibition.
6.What is CPR? A thick brass sheet of area 400 inch 2 is exposed to moist air. After 2
years of period, it was found to experience a weight loss 375 g due to corrosion. If the
density of brass is 8.73 g/cm3. Calculate CPR in mpy and mmpy.

Definition: The Corrosion penetration rate is the speed at which any metal or alloy
deteriorates in a specific corrosive environment through chemical or electrochemical
reactions.
It is also defined as the amount of weight loss per year in the thickness of metal or
alloy due to corrosion. The Corrosion penetration rate also referred as corrosion rate.
CPR = (K x W) / (D x A x T)

CPR in mpy CPR in mmpy

K 534 87.6
W =375 g 375 X 1000mg 375 X 1000mg
D 8.73 g/cm3 8.73 g/cm3
A= 400 inch 2 400 Inch2 400X 6.45 cm2
T=2 years 2X 365X 24hrs 2X 365X 24hrs
CPR in terms of mpy :
CPR= 534 X 375 X 1000 / 8.73 X 400 X 2 X 365 X 24
= 3.273 mpy
CPR in terms of mmpy :
CPR= 87.6 X 375 X 1000 / 8.73 X 400 X 6.45X2 X3 65 X 24
= 0.08325 mmpy

7.What are concentration cell? Explain the construction and working of concentration
cell.

Concentration cell is an electrochemical cell where similar electrodes are in

contact with solutions of the same electrolytes but of different concentrations.
As a result the potential difference arises due to transfer of a substance from
the solution of higher concentration to the solution lower concentration.

Electrolyte concentration cell is formed when electrodes of same substances are dipped
in same kind of electrolyte but of different concentration and are coupled through a
salt bridge. This type of cell is also known as Concentration cell with transference.

Example: Two Zn electrodes are in contact with ZnSO 4 solution of C1 and C2 molar
concentration.

Cell Representation: Zn / Zn2+(C1) // Zn2+(C2) / Zn

 The half cell reactions
At anode (Negative electrode)(Oxidation) Zn Zn2+(C1 ) + 2e
At cathode (Positive electrode)(Reduction) Zn (C2 ) + 2e
2+ Zn

Net Cell reaction Zn2+(C2) Zn2+(C1)

The cell reaction is a change in concentration as a result of which current flows. This
takes place till the concentrations in the two half- cells become equal.
2.303RT C
E conc cell = log 2
nF C1
1)When the two solutions are of the same concentrations, E cell = 0 and hence no current
flows.
2) When C2 > C1, Ecell is positive.
3) Higher the ratio C2 / C1, higher is the value of cell potential

Electrode concentration cell is formed when electrodes of different concentrations are

dissolved in mercury coupled by immersion in a common electrolyte without a salt
bridge. Such type of cell is also known as Amalgam cell or Concentration cell without
transference
Example: Cd-Hg ([Cd]= C1)  CdSO4 (aq)  Cd-Hg ([Cd]= C2) provided, C1 > C2
2.303RT C
Ecell = log 1 cell is operative when C1 > C2
nF C2

8.What are reference electrodes? Explain the construction, working and application
of Calomel electrode

An electrode of known potential, relative to which, the electrode potential of test

electrode can be measured is known as reference electrode.
Calomel electrode: Calomel electrode is a metal-metal insoluble salt electrode and
secondary reference electrode.
Electrode representation: Pt, Hg(l), Hg2Cl2(s)KCl (aq, [KCl] = C)
Construction:
The calomel electrode consists of a glass tube with a side tube. Mercury is placed at the
bottom of the tube. A paste of calomel (Hg2Cl2) and Hg is placed over a pool of Hg. The
remaining part of the tube is filled with saturated KCl solution. Electrical connection is
made through a Pt wire fused into a glass tube dipped in the Hg at the bottom of the
tube. The side tube is filled with saturated solution of KCl with agar-agar acts like a
salt bridge.

Working: Calomel electrode behaves as anode or cathode depending on the nature of

the electrode with which it is connected.
The half-cell reaction when it acts as an anode is 𝐻𝑔 + 𝐶𝑙− → 1 𝐻𝑔 2𝐶𝑙2 + 𝑒−
2
1
The half-cell reaction when it acts as a cathode is 𝐻𝑔2𝐶𝑙2 + 𝑒− → 𝐻𝑔 + 𝐶𝑙−
2
1
Electrode reaction is Hg2Cl2 + e Hg + Cl-
2

2.303 RT
Electrode potential: E = Eo  log
F
[Cl¯]
Electrode potential depends on concentration of KCl solution and potential decreases
with increase in KCl concentration.
Concentration of Potential of calomel Name of the electrode
KCl electrode at 298K
0.1N 0.3358V Deci normal calomel electrode
1.0N 0.2824V Normal calomel electrode
Saturated KCl 0.2444V Saturated calomel electrode

Application: Calomel electrode is used as reference electrode in potentiometric

determinations and also as reference electrode in Ion selective electrodes.

9.Explain the construction and working of ion selective electrode and how it can be
used for the determination of pH of a solution. List out applications of glass electrode
'' Ion selective electrodes are membrane electrodes, which are selectively
sensitive to specific ions in a mixture and potential developed across the
membrane is a function of the concentration of specific ion.''.

Glass electrode: A glass electrode is an ion selective electrode where potential depends
upon the pH of the medium
Electrode representation: Pt, Ag , AgCl(s) |0.1N HCl solution | glass membrane

Construction:
Glass electrode is a long glass tube with thin walled (0.001mm) glass membrane
bulb at the bottom. Glass membrane is made up of corning glass of low MP and
high electrical conductivity with composition 6% CaO, 22% Na 2O and 72% SiO2.
Glass bulb contains 0.1M HCl in which Ag/AgCl electrode is dipped acts as
internal reference electrode and also serves for the external electrical contact.
This electrode is sensitive to H + ions up to a pH of about 9. The electrode dipped
in a solution containing H+ ions (Fig:)
Principle: If a thin walled glass bulb containing an acid is immersed in another solution
containing H+ ions, a potential is developed across the glass membrane. This is due the
ion exchange reaction taking place at the gel layers of glass membrane.
This is called boundary potential Eb .The boundary potential is due to the difference in
potential (E1-E2) developed across the gel layer of the membrane between the two
liquids.

Glass

0.1M HCl
Internal electrolyte

Reference
electrode
Glass membrane

Working:
When a glass electrode is placed in analyte solution containing H + ions. Inner and
outer surface of the glass membrane undergoes ion exchange reaction with H + ions of
solutions and develops potential on both sides of membrane will be different. Since
concentrations of solution inside and outside the membrane are different and hence
difference in potential is developed across the membrane is referred as boundary
potential . If the concentration of H + ( or pH ) of inner solution is kept constant, then
boundary potential becomes proportional to the concentration of H+( or pH) of the
analyte solution. This is the basis of the glass electrode.

The resulting difference in potential between the two surfaces of the glass is the
boundary potential, which is related to the concentrations of hydrogen ion in each of
the solutions by the Nernst-like equation:
( C1 and C2 are concentrations of H+ ions in test solution and internal standard
solution respectively)
RT C
Eb = E1  E2 = ln 1
nF C2

= 0.0591 log C1 /C2 for univalent ion, H+ ion, n = 1 and at 298 K

Eb = L  0.0591 ( log C1 )

= L  0.0591 ( log[H+]) {For a glass electrode, the hydrogen ion concentration of the internal
solution is held constant Hence ( 0.0591 log C2) becomes

constant L}
Eb = L  0.0591 pHtest solution
Glass membrane electrode potential is given by,
Eb = Boundary potential
EG = Eb + EIntRef + Easy EIntRef = Internal reference electrode potential
EG = (L  0.0591 pHanalyte) + EAg/AgCl + Easy Easy = Asymmetric potential when C1 = C2

EG = ( L + EAg/AgCl + Easy)  0.0591 pHanalyte

EG = E oG  0.0591 pHtest
{ E oG = L + EAg/AgCl + Easy = constant}
solution

Determination of PH of solution using glass electrode

To determine the pH of a given solution, the glass electrode is dipped in a test solution and is combined
with reference electrode such as calomel electrode through salt bridge. The cell assembly is represented as
Calomel electrode │Test solution(C1, pH =?)│ Glass electrode
Hg(s),Hg2Cl2(s) KCl (sat) test solution glass membrane 0.1M HCl solution AgCl(s), Ag(s)
The potential of the above cell is measured using an electronic voltmeter and p H of the test solution is
calculated as
Ecell = ER  EL = EG  EExtRef
= E oG  0.0591 pHtest solution  ESCE

H E oG  E SCE  E cell
p test solution =
0.0591
K'  Ecell
 pHtest solution =
0.0591
 o
{K’ = ( E G – EExtRef), glass electrode assembly constant}
The value of constant K' can be calculated by measuring EMF of the above cell dipped in standard buffer
solution of known pH.

K’ = 0.0591 pHStd buffer + E’cell

Substitute the values of cell potential E’cell and pH of standard buffer solution in the above equation.
Applications of Glass electrode:
 pH measurement.
 Food Analysis.
 Cosmetic analysis.
 Environmental regulations.
 Soil acidity determination.
10.Define concentration cell. Emf of the cell Ag/AgNO3(0.001M) // AgNO3(xM) /Ag
is 0.0659 V at 298K. Write the cell representation, cell reactions and calculate the value
of x.

11. Briefly explain the principle, instrumentation and working of potentiometry

taking estimation of Iron as example. Mention the advantages of potentiometry.

Potentiometry is an electro analytical technique. It deals with the determination of

concentration of a solution by measuring the e.m.f. between two electrodes that are
dipped in the solution.

Potentiometric titration:
 Potentiometric titrations involve the measurement of the potential of a suitable
indicator electrode with respect to a reference electrode as a function of titrant
volume. Potentiometry is usually employed to find the concentration of a solute
in solution.
 The equivalence point of the titration is indicated by a rapid change in the
potential at that point. The indicator electrode must respond rapidly to the
change in ion concentration and the change in potential should be large enough
to be measured.
 The equivalence point can be determined more precisely by plotting first
derivative ∆E/∆V against volume of the titrant. The equivalence point is the
volume corresponding to the a maximum.
Principle: Redox titrations can be carried out potentiometrically using platinum
and calomel electrode combination .
For the reaction; Reduced form → Oxidized form + n electrons
The potential is given by Nernst equation
0.0591 [Oxidisedform]
E  Eo  log
n [Reducedform]

0.0591 [Fe3]
E  Eo  log
n [Fe2]

Where, E0 is the standard electrode potential of the system. The potential of the
system is controlled by the ratio of concentration of the oxidized to that of the
reduced species. A plot of change in potential against volume is characterized by
a sudden change of potential at the equivalent point. At the end point, potential
is determined by large jump in the potential value. Hence there is large increase in
potential at the end point.
Instrumentation: Potentiometer consists of a reference electrode, an indicator
electrode and a device for measuring the potential.
Platinum electrode: It is an indicator electrode, used to measure the potential of
the analyte solution comparing with that of reference electrode.
Calomel electrode: It is a reference electrode and used for the determination of
the analyte by maintaining the fixed potential.
Burette: Standard K2Cr2O7 solution.
Beaker: Test solution (FAS) and 5ml of dilute Sulphuric acid ( K2Cr2O7 is a
strong oxidizing agent in acidic medium)

Estimation of Iron present in FAS solution using standard potassium dichromate

solution.
Titration FAS solution with potassium dichromate solution, calomel electrode is used
as the reference electrode and platinum electrode as indicator electrode, which is
actually the oxidation reduction electrode involving redox species in the solution.
Procedure:
 Pipette out 25ml of the given ferrous ammonium sulphate solution into a clean
beaker. Add one test tube of dilute Sulphuric acid. Dip a platinum electrode and
calomel electrode into the solution and connect the electrode assembly to
potentiometer. Calomel electrode acts as reference electrode and platinum
electrode is the indicator electrode.
 The potential is measured. A solution of standard potassium dichromate from a
burette is added in small increments say 0.5cm 3 at a time . Stir the solution
carefully and measure the potential after 30 seconds. Continue the addition of
potassium dichromate solution till there is a sudden rise in the potential (At this
stage wait for 5 minutes ).
 The potential initially increases gradually. Then there is a steep rise in the
potential, later it remains constant. A few readings are taken after the steep rise.
 A differential graph of change in potential against volume of dichromate is
plotted. The peak gives the equivalence point. From the equivalence point,
concentration of FAS in the analyte is calculated. Finally iron present in FAS
solution is calculated.

Advantages
 Potentiometric titrations can be carried out for colored solutions.
 The end point in the titration of very weak acids or very weak bases can be
obtained.
 It is applicable for turbid, fluorescent and opaque solution

12.Explain the principle, instrumentation and working of conductometry taking

estimation of weak acid using a strong base as an example
The electrochemical method of analysis used for the determination of the electrical
conductance of an electrolyte solution by means of a conductometer.

"Conductometry deals with the determination of concentration of the

electrolytic solution by measuring its conductance by using conductivity
meter".
Theory : The conductance of the solution is explained by considering ohm’s law.
Electrolyte solution conducts current by the migration of ions under the influence of an
electric field and like a metallic conductor, they obey Ohm’s law.
The current i (amperes) flowing in a conductor is directly proportional to the applied
electromotive force, E (volts), and inversely proportional to the resistance R (ohms) of the
E
conductor. i
R

The reciprocal of the resistance is called the conductance, C = 1/R

The resistance of the conductor is directly proportional to the length(l) and inversely
proportional to the area of cross section(a) of the conductor.
R = ρ (l/a) (ρ is specific resistance or resistivity of the conductor)
C = 1/R = (1/ ρ) (a/l) = κ(a/l) κ = specific conductance or conductivity of the
electrolyte solution.
κ = C (l/a)
Or Specific conductance = Conductance × cell constant
When a = 1 and l =1, C = κ
" Specific conductance is the conductance of an electrolyte solution kept
between two electrodes of 1m2 cross sectional area at 1 m apart".
Conductometric titration:
 Conductomertic titrations are based on the measurement of conductance of
solution, which is mainly depends on number of ions, charge on the ions and
mobility of the ions.
 The replacement of ions of a particular conductance by ions of different conductance
during the titration.
 Hence the end points of varies titrations determined accurately .
 The variation of electrical conductance of a solution during the course of the
titration is followed. The conductance is measured after each addition of small
volume of the reagent.
  The conductance values thus obtained are plotted against the volume of the
reagent added. The plot obtained is consists of two straight lines intersecting at
equivalence point.

Instrumentation: The basic units used in measuring conductance of a solution are 1.

Conductivity cell 2.Conductivity meter. Conductivity cell consists of two platinum
electrodes have unit area of cross section and are placed unit distance apart. A simple
arrangement of conductometric titration is depicted in figure. The solution to be
titrated is taken in the beaker.

Conductivity cell

Estimation of weak acid by using a strong base: (CH 3COOH vs NaOH)

 In the conductometric titration of a weak acid with a strong base, the

conductance of the acid will be initially low due to poor dissociation of acetic
acid.
 When a strong base is added to the acid, the salt formed is highly ionized and the
conductance increases.
 On complete neutralization of the acid, further addition of base leads to an
increase in the number of more mobile OH - ions. And hence conductance
increases sharply.
 Thus a plot of conductance against volume of base added gives two straight lines
as shown in figure.
 The point of intersection of the two lines gives the neutralization point.
 In the above graph, point (B) represents the equivalent point or the
neutralization point

CH 3COOH + NaOH → CH 3COONa + H2O

Advantages:
 Mixture of acids can be titrated more accurately by conductometric titration.
 Conductometric titrations may be applied where potentiometric methods fail.
 Accurate in dilute solution as well as in more concentrated solution.
 It can be employed with colored solutions.
 Very weak acids which cannot be titrated potentiometrically in aqueous solutions
can be titrated conductometrically with relative ease.
 Applied Chemistry for Computer Science & Engineering stream Dr S K REVATHI

Specific conductance is also called conductivity.

A R
Further, 
l
1 l 1
  
 A R
1
 C
A
Or Specific conductance = Conductance × cell constant
Specific conductance is defined as the conductance of a solution enclosed in a cell
having two electrodes of unit area separated by one centimeter apart. The
conductance of the solution is depends on mobility of the ions and number of the
ions present in the solution. Unit of specific conductance: Ohm–1 cm–1

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Applied Chemistry for Computer Science & Engineering stream Dr S K REVATHI

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Applied Chemistry for Computer Science & Engineering stream Dr S K REVATHI

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Applied Chemistry for Computer Science & Engineering stream Dr S K REVATHI

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