Electrochemistry

Sherif M.A.S. Keshk; M. Sc.; D.Eng.

Electroplating
Corrosion Batteries Fuel Cell
1 Sherif Keshk of metals
 Contents
• Introduction
• Oxidation numbers
• Galvanic cells
• Standard reduction potentials
• Cell potential, electrical work, and free energy
• Dependence of cell potential on concentration
Types of Electrodes
• Application of electromotive forces
• Batteries
• Corrosion
• Electrolysis
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• Homework will be taken in randomly
• Quiz will be given 40%

• Electrochem Unit Exam at the end 60%

• Lab Report will be expected

• This course will cover Electrochemistry and its

applications

Sherif Keshk 3
 Electrochemistry

3 Sherif Keshk
 Electrochemistry
ELECTROCHEMISTRY

Types of Process

Reversible process Irreversible process

In this process the In this process the
chemical reaction is used electrical energy is used
to produce electrical to produce chemical
energy reaction.

Example:Galvanic cell. Example:Electrolytic

cell.

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 Electrochemistry
What happen if a bar of Zinc is immersed onto copper sulfate
solution?

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 Direct redox reaction
Oxidizing and reducing agents are mixed together

2+
Zn → Zn + 2e
2+
Cu + 2e → Cu

Oxidation reaction produces electrons

Reduction reaction consumes electrons

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 Oxidation-Reduction Reactions

Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)

Zn Zn2+ + 2e- Zn is oxidized

Zn is the reducing agent

Cu2+ + 2e- Cu Cu2+ is reduced

Cu2+ is the oxidizing agent

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 Oxidation-Reduction Reactions

2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e- Oxidation half-reaction (lose e-)

O2 + 4e- 2O2- Reduction half-reaction (gain e-)

2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-

2Mg + O2 2MgO
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 Oxidation-Reduction Reactions

• Oxidation-reduction reactions always involve a transfer of

electrons from one species to another.

• Oxidation : Lose electron - higher ox state

Fe2+ → Fe3+ + e-

• Reduction : Gain electrons - lower ox state

Fe3+ + e- → Fe2+

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 Oxidation-Reduction Reactions
Oxidation-Reduction Reactions
• Oxidizing Agent- a substance that accepts
electrons from another substance, causing
the other substance to be oxidized.

• Reducing Agent- a substance that donates

electrons to another substance, causing it
to become reduced.

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 Oxidation-Reduction Reactions
Oxidation Number
• Oxidation Number- the number of charges in
the atom would have in a molecule (or an
ionic compound) if electrons were transferred
completely.

• H2(g) + Cl2(g) → 2HCl(g)

• 0 0 +1 -1

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 Assigning Oxidation Numbers
1. Free elements (uncombined state) have an
oxidation number of zero.

Na, Be, K, Pb, H2, O2, P4 = 0

2. In monatomic ions, the oxidation number is equal

to the charge on the ion.

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2

3. The oxidation number of oxygen is usually –2. In

H2O2 and O22- it is –1.

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 Assigning Oxidation Numbers

4. The oxidation number of hydrogen is +1 except

when it is bonded to metals in binary compounds.
In these cases, its oxidation number is –1.

5. Group IA metals are +1, IIA metals are +2 and

fluorine is always –1.

6. The sum of the oxidation numbers of all the atoms

in a molecule or ion is equal to the charge on the
molecule or ion.

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 Assigning Oxidation Numbers

Oxidation numbers of
all the elements in
HCO3- ?

HCO3-
O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4
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 Redox Reactions
• Direct redox reaction
– Oxidizing and reducing agents are mixed together
• Indirect redox reaction
– Oxidizing and reducing agents are separated but
connected electrically
• Example
– Zn and Cu2+ can be reacted indirectly
– Basis for electrochemistry
– Electrochemical cell

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 Electrochemical cell
• An electrochemical cell is a system consisting of electrodes that dip
into an electrolyte in which a chemical reaction either uses or
generates an electric current.

• I- A voltaic, or galvanic, cell is an electrochemical cell in which a

spontaneous reaction generates an electric current (+V, rxn
occurs on own).

• II- An electrolytic cell is an electrochemical cell in which an electric

current drives an otherwise nonspontaneous reaction (-V, rxn
does not occur on own, need outside source).
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 Components of Electrochemical cell
Electrochemical Cells
• Voltaic Cell
– Cell in which a spontaneous redox reaction
generates electricity
– Chemical energy → electrical energy

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 Components of Electrochemical
Electrochemical Cells cell

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 Electrochemical Cells

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Generally, salts like KCl, KNO3, etc. are used. The seturated
solutions of these electrolytes are prepared in agar agar jelly or
gelatin. The jelly keeps the electrolyte in semi-solid phase and
thus prevents mixing.
The important functions of the salt bridge are:
a) Salt bridge completes the electrical circuit.
b) Salt bridge maintains electrical neutrality of two half cell
solution.
The accumulation of charges in the two half cells (accumulation of
extra positive charge in the solution around the anode according
to the realizing of Zn2+ in excess and accumulation of extra
negative charge in the solution around the catode due to excess
of SO42- ) is prevented by using salt bridge, which provides a
passage for the flow of the charge in the internal circuit.
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 Example of galvanic cell :
Danniel cell:
The Danniel cell consists of Zn electrode
immersed in Zn Sulphate (ZnSO4) → (anode)
and Cu electrode immersed in CuSO4
(cathode). The two electrode are immersed in
beaker glass and separated by porous
diphgram. The porousdiphgram prevents the
mixed between solution and permits to pass
the electron or ions.

The Daniel cell can be represented as :

Zn /ZnSO4 // Cu /CuSO4
E  = −0.76 V E  = 0.34 V

From the values of E (standard electrode

potential).

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 Danniel Cell

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23 Sherif Keshk
 Zn electrode is anode.
Cu electrode is cathode.

1) At node Zn → Zn 2+ + 2e E  = −0.76 V

2) At Cathode Cu 2+ + 2e → Cu E  = 0.34 V

Zn + Cu2+ Zn2+ + Cu

E ocell = Ecathode
o
− Eanode
o

= 0.34 – (– 0.76) = 1.1 V

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 Example: Write the electrochemical reaction and calculate the electrode
potential of this cell.

Zn/Zn+2 // Ag + Ag
E = - 0.76 V E = 1.7 V

(Solution)
From the values of E :
The lowest value is the anode (Zn) and the highest value is the
cathode (Ag).

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 1) Al anode Zn → Zn 2+ + 2e E  = 0.34 V

2) At cathode Ag + + e → Ag E  = 1.7 V

by multiplying equation (3)  2

 equation 2 become
2 Ag + + 2e → 2 Ag E  = 1.7  2 = 3.4 V

 The total reaction:

Zn + 2Ag + → Zn+2 +2Ag

o
E cell = E cathode
o
− E anode
o

= 3.4 – (-0.34) = 3.74 V

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 Another example galvanic cell (gas cell).
It is consists of two pt electrode immersed in
HCl.

H2 gas is passed on one electrode and Cl2 gas

is passed on the another electrode.

Reaction at anode:

1/2 H2(gas) → H(adsorbed) → H+ + e

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 Reaction at anode:

1/2 H2(gas) → H(adsorbed) → H+ + e

1
Cl 2 → Cl ( adsorbed ) + e → Cl −
2
The total reaction:

H + Cl → H + + Cl −
The pt electrode not reacted but facilitate the
adsorption of gases on its surface.

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Voltaic Cells (aka Galvanic Cell)
• A voltaic cell consists of two half-cells that are electrically connected.
• Each half-cell is a portion of the electrochemical cell in which a half-
reaction takes place.
No reaction between species involved just a transfer of electrons
A simple half-cell can be made from a metal strip dipped into a
solution of its metal ion.

For example,
The silver-silver ion half cell consists of a silver
strip dipped into a solution of a silver salt.

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• Voltaic Cells (aka Galvanic Cell)
• Composed of two half-cells; which each consist of a metal rod or strip
immersed in a solution of its own ions or an inert electrolyte.
• Electrodes: solid conductors connecting the cell to an external circuit
• Anode: electrode where oxidation occurs (-)
• Cathode: electrode where reduction occurs (+)
• The electrons flow from the anode to the cathode (“a before c”)
through an electrical circuit rather than passing directly from one
substance to another
• A porous boundary separates the two electrolytes while still allowing
ions to flow to maintain cell neutrality

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• The positive electrode is defined as the cathode and
the negative electrode is defined as the anode

• The electrons flow through the external circuit from

the anode to the cathode.

• To test the voltage of a battery, the red (+) lead

is connected to the cathode (+ electrode),
and the black (-) lead is connected to
the anode (- electrode)

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• Often the porous boundary is a salt bridge,
containing an inert aqueous electrolyte

such as Na2SO4(aq) or KNO3(aq)

• Or you can use a porous cup containing one

electrolyte which sits in a container of a second
electrolyte.

Sherif Keshk 34
 Cathode & reduction

• If the electronic conductor gives up electrons to

the ionic conductor, then the electrode is called as
a cathode and the reaction is called as a cathodic
reaction. The flow of electrons and ions at a
cathode is shown in the figure below.

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 Cathode & reduction
• The copper strip is the electronic conductor that
gives up electrons to the ionic conductor which
is the copper (II) sulfate solution.
• The chemistry that takes place in the cathode is reduction.
Cu2+ (aq) +2e- → Cu(s)

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• During reduction copper is deposited on the metal strip.
Solid metal appears as a product.
• This process of depositing a metal on a conducive surface is
used in the electroplating of silver into jewelry and flatware.

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 Anode & oxidation

• An anode is an electrode at which electrons are

generated. A reaction occurring at the anode is
anodic reaction. Electrons flow in the opposite
direction to the positive electric current .

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 Anode & oxidation

• The figure below shows the reaction of zinc

atoms to form zinc ions.

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 Anode & oxidation

• A reaction occurring at an anode is an anodic reaction, and

the resulting chemical change is called oxidation. Zinc atoms
are oxidized as they lose electrons at the anode. The metal
strip dissolves as the reaction proceeds.

• Reduction reactions always occur with oxidation reactions.

The electrons used in reduction must come from an
oxidation reaction. The overall reaction is called an oxidation
reduction reaction, or redox reaction.

40 Sherif Keshk
 Nernst Equation
Nernst Equation

Dependence of electrode potential on the activity of

ions.

Firstly: G= Gproduct - Greactant

G0 = G 0 product- - G0reactant
Where : G → free energy
G → standard free energy
G = - Z F E
G 0 = - Z F E0
Where : E → electrode potential
E0 → Standard electrode potential.

G = G0 + R T ln a
A is the activity of ions.

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 Nernst Equation

Nernst equation:
Suppose the following equation:
A+B C+D
G = Gproduct – Greactant

= (GC + GD) – (GA + GB)

( ) (
= GCo + RT ln aC + GDo + RT ln aD − )
(G o
A ) (
+ RT ln a A + GBo + RT ln a B )

( ) ( )
G = GCo + GDo − G Ao + GBo + RT ln
aC . a D
a A . aB

a product
G = G o + RT ln
areac tan t

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 Nernst Equation
a product
− ZFE = − ZFE o + RT ln
areac tan t

RT a product
E = Eo − ln
ZF areac tan t

Where R is ideal gas constant and equal to 8.314

T is the temperature and equal to 25+273 =298
F is the Faraday and equal to 96500 Coloumb
Z is the number of electron
ln = 2.303 log
RT
 ln → 0.059 log
F

0.059 a product
 E = E − log
a areac tan t
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 Nernst Equation
Problems:
1- Write the electrochemical reaction and calculate the electrode potential
of this cell
2+
Zn Zn Cu 2+ Cu
(0.1 M ) (0.01 M )
E  = −0.76 V E  = 0.34 V

Answer
From the values of E
Zn is anode and Cu is cathode.

Zn → Z n2+ + 2e E  = −0.76 V

Cu 2+ + 2e → Cu E  = 0.34 V

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 Nernst Equation
Zn + Cu 2+ → Z n2+ + Cu

 The electrochemical reaction:

Zn + Cu 2+ Z n2+ + Cu

E of the cell = E cathode − E anode

o o

= 0.34 − (− 0.76) = 1.1 V

From the Nernst equation:

E=E −
0.059 
Zn 2+ Cu  
 
o
log
Z Cu 2+ Zn 

Cu  = Z n  = 1
solid solid

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 Nernst Equation

0.059 [Zn 2+ ]
 E = 1.1 − log 2+
Z [Cu ]
0.059 0.1
= 1.1 − log =?
2 0.01

46 Sherif Keshk
 Nernst Equation
Problem 2
Write the electrochemical reaction and calculate the concentration
of Zn2+.
2+
Zn Zn Ag + Ag
(?) (1 M )
E  = −0.76 V E  = 1.7 V

Where the electrode potential of the cell = 2 V.

Problem 3
Write the electrochemical reaction and calculate the concentration
2+
of Sn in the following cell.

Zn Zn 2+ Sn 2+ Sn
(1M ) (?)

Where the electrode potential of the cell equal to 1.0V.

E of Zn/Zn+2 =-0.76 V
E of Sn/Sn+2 = -0.14V

47 Sherif Keshk
 Electrochemical Reactions

There are two system for writing the

electrochemical reaction.
1) Eureopean system:
In this system, all the reaction are written as
cathodic reaction.
Cu 2 + + 2e → Cu

Z n2+ + 2e → Zn
+
Ag + e → Ag
Example:
Cu 2 + + 2e → Cu

G = G p − Gr

= GCu − GCu 2 +

(
= GCu
o
+ RT ln a Cu ) − (G o
Cu 2 +
+ RT ln a Cu 2 + )
(
= GCu
o
− GCu
o
2+ )+ RT ln
a Cu
aCu 2 +

48 Sherif Keshk
 Electrochemical Reactions
ap
 G =  G + RT ln
o

ar
ap
− Z F E = − Z F E + RT ln
o
ar
RT ap
E=E − o
ln
ZF ar

0.059 aP
E=E − o
log
Z ar

49 Sherif Keshk
 Electromotive Force

Calculation of the electromotive force (E.M.F) from the

electrode potential
1) Eureopesn system (Reduction Table):
E = Ecathode – Eanode
Example Daniell cell:
 RT   RT 
E =  0.34 + ln aCu 2 +  −  − 0.76 + ln a Zn2 + 
 ZF   ZF 

RT aCu 2 +
E = 1.1V + ln
ZF a Zn2 +

 E o = 1.1 V

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 Spontaneity

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 Spontaneity

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53 Sherif Keshk
 Spontaneity

54 Sherif Keshk
 Spontaneity

Spontaneous reactions:
The reactions which accompany by decrease in free
energy (G = - ve) to obtain positive electrode potential
(E = t).

Example: Zn + Cu 2 + → Zn + Cu 2 +

This reaction is not spontaneous

G = +

Where : E = −
 G = − ZFE

= −  −
= + ve
 When electrode potential is +ve.
The reaction is spontaneous.
G = − ZFE = −  + = −

 When electrode potential is –ve

The reaction is nonspontaneous
G = −  − = +

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 Standard Cell
Standard Cells:
In Standard cells (Electromotive force).
- The electrode potential of this cell is
constant for long period.
Not affected by temperature
56 Sherif Keshk
 Standard Cell
There are two examples:
(1) Standard western cell:
CdSO4 , 3CdSO4 . 8 H 2 O Cd ama lg am
Hg / Hg 2 SO4
Saturated Solid 12.5%
Solid
Solution
It is consists of :
- Positive (+) electrode consists of Hg coated with solid Hg2SO4
(Anode).
- Negative (-) electrode consists of amalgam Cd (12.5%)(cathode).
- The electrolyte containing saturated solution of Cd SO4. Also solid
CdSO4. 8H2O is added amalgam Cd is dissolved (anode) and the
Hg is precipitated.

Cd + Hg 2 SO4 = CdSO 4 + 2 Hg

The electromotive force of this cell in constant = 1.0183 Volt

 0.0004V

57 Sherif Keshk
 Standard Cell
2- Standard clark cell:
It is consists of
ZnSO4 , ZnSO4 . 7 H 2 O Zn ama lg am
Hg / Hg 2 SO4
Saturated Solid 10%
Solid
Solution

- The positive (+) electrode (cathode)

consists of Hg coated with solid Hg2SO4.
- The negative (-) electrode (anode) consists
of 10% amalgam Zn.
- The electrolyte contains saturated solution
of ZnSO4 and solid ZnSO4 to keep the
saturation of solution.

Reaction:
Zn + Hg 2 SO4 → ZnSO4 + 2 Hg

The electromotive force of this cell =

1.432 Volt.

58 Sherif Keshk
 Types of Electrodes
The electrochemical reaction.
M 2+ + Ze → M

From Nerns't equation.

RT a M 2+
EM M 2+
= E o
M M 2+
+ ln
ZF aM

0.059
 EM M 2+
= EMo M 2+
+ log aM 2+ = aM = 1
Z

 The electrode potential of the first type is dependent on

the activity of metal ions.

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 Types of Electrodes

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 I- Standard Hydrogen Electrode

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 I- Standard Hydrogen Electrode
Hydrogen electrode:
+
pt H , H2

+
2H + 2e → H2

EH = EH
o
+
0.059
log
a H + 
2 H 2 
H 2  = PH 2
=1

EH = EH
o
+ 0.059 log a H +

E H = 0.00 + 0.059 log a H +

Since ( pH = − log a H + )
E H = −0.059 pH

 The electrode potential of hydrogen electrode

depending on the activity of Hydrogen ion or pH of
solution.
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 II- Metal Electrode
Electrode of the second type

It is consist. of a met at covered by one of

its slightly soluble salts and immersed in
electrolyte which have the same anion of its
slightly soluble salts example
Hg Hg 2 Cl 2 KCl → calomel electrode
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 II- Metal Electrode
or Ag AgCl KCl
+
Since Ag AgCl Ag + e → Ag
0.059
 E Ag = E Ag
o
+ log a Ag +
Z
E Ag = E Ag
o
+ 0.059 log a Ag +

AgCl is slightly soluble salt.

+
AgCl → Ag + Cl −

The solubility product Ks = a Ag + . a Cl −

Ks
a Ag + =
aCl −

Ks
E Ag = E Ag
o
+ 0.059 log
a Cl −
o
E Ag + 0.059 log K s = E Ag
o

o
E Ag is constant. Ks is constant

o
 E Ag = E Ag − 0.059 log a Cl −

 The electrode potential of Ag is dependent on the concentration of

anion (Cl −) not cation (Ag+).

64 Sherif Keshk
 III- Solid Electrode
Third type of electrode:
It is consisted of solid electrode immersed in two slightly soluble
salts. One of them is slightly soluble salt of this solid electrode and the
other is slightly soluble salt of another active metal.
Zn Zn oxalate Ca oxalate Ca
Sparingly Sparingly
Soluble soluble
Salts salt.

Consider Zn Zn Ox
RT
E Zn = E Zn
o
+ ln a Zn 2 +
ZF
2+ 2−
Since: Zn Ox → Zn + OX

Ks = a Zn 2 + . a 6 X 2 −
Ks
aOX 2 − =
a Zn 2 +

65 Sherif Keshk
 IV- Organic Electrode
Quinohydrone Electrode:
It is consisted of pt electrode immersed in solutions
of quinione and hydroquinone (H2q).

pt q , H 2 q

q + 2 H + + 2e → H 2q

a q . (a H + ) 2
RT
E q + E qo + ln aq = a H 2q
ZF a H 2q

RT
E q = E qo + ln a H +
F
66 Sherif Keshk
 Quinone hydron Electrodes

Eq = E + 0.059 log aH +
o
q

E q = E − 0.059 pH
o
q

Eq = 0.699 − 0.059 pH
The electrode potential of quinonehydrone electrode
depending on the pH of solution.

67 Sherif Keshk
 Quinone hydron Electrodes
The scheme of quinhydrone cell with one
electrolyte
Pt, Н2 | quinhydr, H+ | KCl | KCl,Hg2Cl2| Hg

Ecell = E quinhydr - Ecalomel

68 Sherif Keshk
 Concentration Cell
Concentration Cell: both compartments contain same
components but at different concentrations

Half cell potential are not

identical
Because the Ag+ Conc.
On both sides are not
same

Eright > Eleft

• To make them equal, [Ag+]
On both sides should same
• Electrons move from left to
right

69 Sherif Keshk
 Application of E.M.F.
Applications of E.M.F.

1- Det. of Activity and Activity Coefficient

In dilute solution a=C
Concentration
activity
a=f.C
Where f → activity coefficient
RT
E = E +
ZF
ln a  a can be

determined
RT
E = E + ln f . C
ZF

 f can be determine
70 Sherif Keshk
 Application of E.M.F.
2- Det. of Equilibrium constant
Consider the reaction aA → bB
[ B ]b
K =
[ A] a

0.059 [ B ]b
E = E − log
Z [ A] A

at Eqb ( E = 0)

0.059
 E = log K
Z
71 Sherif Keshk
 Application of E.M.F.
Problem:
Calculate the eqb constant for the reaction
Sn + pb 2 + S n2 + + pb

Where e of Sn Sn 2 + = − 0.141 V

pb / bp 2 + = −0.121 V

solution
o
E Cell = E Cathode
o
− E anode
o

= (− 0.126 ) − (− 0.141 ) = 0.015 V

0.059
 E = log K
Z

0.059
0.015 = log K
2

K =?

72 Sherif Keshk
 Application of E.M.F.
3- Det of solubility product (KS):
Consider the electrode Ag/AgCl /KCl
For Ag/ AgCl Ag + + e → Ag

0.059 g +
E Ag = E Ag
o
+ log aAg
Z

For AgCl Ks =a Ag+. aCl-

Ks
 aAg + =
aCl −

 E Ag = E Ag
o
+
0.059
Z
log
Ks
aCl −
consider aCl- = 1

 E Ag = E Ag
o
+
0.059
Z
log Ks 

Ks can be determine
73 Sherif Keshk
 Application of E.M.F.

4- Det. of volency of ions (Z)

If metal is dipped in its solution the electrode potential
M M 2+
RT
 E = Eo + ln a M + z
ZF
Z can be calculated.

74 Sherif Keshk
 Application of E.M.F.
For concentration cell:
Consider the cell
Hg 0.05 N Hg 2 ( NO3 ) 2 0.5 N Hg 2 ( NO3 ) 2 Hg
I II

RT
E I = E Io + ln a Hg ( I )
ZF
RT
E II = E IIo + ln a Hg ( II )
ZF
E Cell = E II − E I

RT 0.5
E cell = ln Since E I = E II
o o

ZF 0.05

0.059 0.059
E cell = log 10 Z =
Z E cell

It was found that E of cell = 0.029

0.059
Z = =2
0.029
75 Sherif Keshk
 Application of E.M.F.
5- Determination of pH of solution:
By combining H – electrode with reference electrode in cell

Hg Hg 2 Cl2 Kcl H + ( H 2 ) pt
( Ref.electrode) ( H − electrode)

E cell = E ref − EH

 RT 
 E cell = E ref − ln a H t 
 F 

RT
Where EH = EH
o
− ln a H + o
EH = o
F

E cell = E ref − 0.059 log a H +

Ecell = Eref + 0.059 pH

By Knowing Ecell and Eref

pH can be determined
76 Sherif Keshk
 Application of E.M.F.
6- Determination of transition point:
The transition point is defined at the temperature at which the two
all tropic form is in eqb E=0

For example:
Sn grey SnCl 2 Sn white
- By measuring temperature against E
- There is a linear relationship between E and T as shown in Figure.
- At transition point 18C
- The electrode potential E = 0
- The temperature at which E = 0 is known as transition point.

77 Sherif Keshk
 Application of E.M.F.
Problems
1- Consider the cell:
pt H2 HCl Hg 2 Cl 2 Hg
10tm a =1
e o = 0.0 V e = 0.789 V

The electrode potential E = 0.268 at 25C

Calculate KS of Hg2Cl2

78 Sherif Keshk
 Application of E.M.F.
2- E.M.F. of the cell
Hg Hg 2 Cl 2 KCl quinohydrone / pt

is 0.219 V at 25C. calculate the pH of solution.

E of calomel electrode = 0.2415 V and

Eo of quinohydrone is 0.699 at the same temperature.

79 Sherif Keshk
 Application of E.M.F.
o
3- Write the electrochemical reaction and calculate E of the cell and
the eqb constant for the following cell

2+ 4+ 3+ 2+
pt Sn Sn Fe Fe pt
E = 0.15 V
o
E = 0.77 V

80 Sherif Keshk
 Application of E.M.F.
4- Consider the cell
Hg Hg 2Cl2 KCl Fe3+ Fe2+ pt
calome H − electrode

The electrode potential of the cell is

0.516 V. Claculate the pH of the solution
if ecalomel is 0.283 V.

81 Sherif Keshk
 Application of Galvanic Cells

FUEL CELLS BATTERY

Sherif Keshk 82
 Classification (or) Types of batteries

Primary battery (Primary cells)

In which the cell reaction is not reversible. When all the
reactants have been converted to product, no more
electricity is produced, and the battery is dead.
Ex: Dry cell, Leclanche Cell

Secondary battery (secondary cells)

In which cell reactions can be reversed by passing electric
current in the opposite direction. Thus, it can be used for
many cycles
Ex: Lead-Acid Batteries, Ni-Cd, Li-ion

Sherif Keshk 83
Dry-Cell Battery This is a common galvanic cell,
known as voltaic cell or Daniell cell.
which contains a moist ammonium chloride electrolyte.

Sherif Keshk 84
 Consequently, the zinc eventually corrodes galvanically
since it provides the electrons to the electrolyte for
generating reduction reactions. The electrolyte (moist
paste) carries the current from the zinc anode to the carbon
cathode . This particular electrochemical cell is a well-
sealed battery (dry-cell battery) useful for flashlights,
portable radios, and the like. Nonetheless, zinc casing
oxidizes, Zn →Zn + 2e, at the anode surface

, while MnO2 is reduced at the carbon cathode surface, and

ammonia is then produced, decreasing the cell current by
forming a gaseous film around the carbon electrode. Logically,
the dry cell battery does not last long when used continuously.
However, if this battery is partially used, its design life is
slightly extended.

Sherif Keshk 85
 Lead-Acid Battery The basic operation of a lead-
acid .Pb-H2SO4/ battery is based on groups of
positive and negative plates immersed in an electrolyte
that consists of diluted sulfuric .H2SO4/ acid and water.
Hence, the mechanism of this type of battery is based on
the electron-balanced anodic (-) and cathodic (+)
reactions. Hence, the ideal reactions are

Sherif Keshk 86
 Hydrogen cation (HCl) in the anode half-cell crosses the Pt
catalyst-coated porous membrane to react with oxygen
within the cathode half-cell to form water, which is carried
away through design (not shown) channels. This simple
mechanism is powerful enough to produce energy for
implantable devices, portable electronic devices,

Sherif Keshk 87
 Batteries are Galvanic Cells

Batteries
Lead-Storage Battery
◼ A 12 V car battery consists of 6 cathode/anode
pairs each producing 2 V.
◼ Cathode: PbO2 on a metal grid in sulfuric acid:

PbO2(s) + SO42-(aq) + 4H+(aq) + 2e- →

PbSO4(s) + 2H2O(l)
◼ Anode: Pb:

Pb(s) + SO42-(aq) → PbSO4(s) + 2e-

88 Sherif Keshk
 Lead storage battery
Batteries are Galvanic Cells

Anode: Pb (s) + SO2- (aq)4 PbSO4 (s) + 2e-

Cathode: PbO2 (s) + 4H+ (aq) + SO2-4 (aq) + 2e- PbSO4 (s) + 2H2O (l)

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) 2PbSO4 (s) + 2H2O (l)

89 Sherif Keshk
 Batteries
Dry cell are Galvanic Cells
Batteries

Anode: Zn (s) Zn2+ (aq) + 2e-

Cathode:2NH4+(aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)

Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) +
Mn2O3 (s)
90 Sherif Keshk
 Batteries are Galvanic Cells
Dry Cell Battery
◼ Anode: Zn cap:
Zn(s) → Zn2+(aq) + 2e-
◼ Cathode: MnO2, NH4Cl and C paste:

2NH4+(aq) + 2MnO2(s) + 2e- → Mn2O3(s) + 2NH3(aq)

+ 2H2O(l)
◼ Total reaction:

Zn + NH4+ +MnO2 → Zn2+ + NH3 + H2O

◼ This cell produces a potential of about 1.5 V.

◼ The graphite rod in the center is an inert

cathode.
91 Sherif Keshk
 Batteries are
Alkaline Cell Galvanic Cells
Battery

◼For an alkaline battery, NH4Cl is replaced with

KOH.
◼Anode: oxidation of Zn

Zn(s) + 2OH- → ZnO + H2O + 2e-

◼Cathode: reduction of MnO2.

2MnO2 + H2O + 2e- → Mn2O3 + 2OH-

• Total reaction
Zn(s) + 2 MnO2(s) ---> ZnO(s) + Mn2O3(s)
◼It lasts longer because Zn anode corrodes

less rapidly than under acidic conditions.

92 Sherif Keshk
 Batteries are Galvanic Cells
Alkaline Battery

93 Sherif Keshk
 Fuel Cells
Batteries are Galvanic Cells
A fuel cell is a galvanic cell that requires a continuous
supply of reactants to keep functioning

Anode: 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-

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

94
2H2 (g) + O2 (g) Sherif Keshk
2H2O (l)
 Corrosion
Corrosion

• Rusting - spontaneous oxidation of metals.

• Most metals used for structural purposes have reduction
potentials that are less positive than O2 . (They are readily
oxidized by O2)
• Fe+2 +2e- → Fe Eº= - 0.44 V
• O2 + 2H2O + 4e- → 4OH- Eº= 0.40 V
• When a cell is formed from these two half reactions a cell
with +ve potential will be obtained
• Au, Pt, Cu, Ag are difficult to be oxidized (noble metals)
• Most metals are readily oxidized by O2 however, this
process develops a thin oxide coating that protect the
internal atoms from being further oxidized.
• Al that has Eo = -1,7V is easily oxidized. Thus, it is used for
making the body of the airplane.

98 Sherif Keshk
 Corrosion
Electrochemical corrosion of iron
Salt speeds up process by increasing conductivity

Water Cathodic area

Rust

Anodic area

Iron dissolves forming a pit e-

Fe Fe+2 + 2e- O2 + 2H2O + 4e- 4OH-
Anodic reaction cathodic reaction
Fe2+ (aq) + O2(g) + (4-2n) H2O(l) →
2F2O3(s).nH2O (s)+ 8H+(aq)
99 Sherif Keshk
 Corrosion
• Fe on the steel surface is oxidized (anodic regions)
• Fe → Fe+2 +2e- Eº=- 0.44 V
• e-’s released flow through the steel to the areas that have
O2 and moisture (cathodic regions). Oxygen is reduced
• O2 + 2H2O + 4e- → 4OH- Eº= 0.40 V
• Thus, in the cathodic region Fe+2 will react with O2
• The total reaction is:
Fe2+ (aq) + O2(g) + (4-2n) H2O(l) →
2F2O3(s).nH2O (s)+ 8H+(aq)
• Thus, iron is dissolved to form pits in steel
• Moisture must be present to act as the salt bridge
• Steel does not rust in the dry air
• Salts accelerates the process due to the increase in
conductivity on the surface
100 Sherif Keshk
 Preventing of Corrosion

• Coating to keep out air and water.

• Galvanizing - Putting on a zinc coat
• Fe Fe2+ + 2e- Eoox = 0.44V
• Zn Zn2+ + 2e- Eoox = 0.76 V
• Zn has a more positive oxidation potential than Fe, so it is
more easily oxidized.
• Any oxidation dissolves Zn rather than Fe
• Alloying is also used to prevent corrosion. stainless steel
contains Cr and Ni that make make steel as a noble metal
• Cathodic Protection - Attaching large pieces of an active
metal like magnesium by wire to the pipeline that get
oxidized instead. By time Mg must be replaced since it
dissolves by time

101 Sherif Keshk

 Preventing of Corrosion
Cathodic Protection of an Underground Pipe

102 Sherif Keshk

 Preventing of Corrosion
Cathodic Protection of an Iron Storage Tank

103 Sherif Keshk

 Faraday’s Laws of Electrolysis
Preventing of Corrosion

If you wished to electroplate an object, you might ask yourself the

following questions:
How can I determine how much metal is being plated?
How long should I leave the object being plated in the
electroplating cell?
What size electric current should be used?

104 Sherif Keshk

 Faraday’s Laws of Electrolysis
Preventing of Corrosion

Michael Faraday was a 19th century English chemist.

His studies yielded two laws relating amount of electric current
and the chemicals produced by the current or used to produce it.
These laws enable us to:
Work out the amount of energy required to discharge a metal ion
and place it out on an object or
To calculate the amount of metal produced in an electrolytic cell.

105 Sherif Keshk

 Faraday’s Laws of Electrolysis – First Law
Preventing of Corrosion
The mass of metal produced at the cathode is directly
proportional to the quantity of electricity passed through the cell

mQ

Electric charge, Q, is measured using the unit coulomb.

The electric charge passing through a cell may be calculated

from measurements of the current, I, through the cell and the
time, t, for which the current flows.

Charge (coulombs) = current (amps) x time (seconds)

Q = It
106 Sherif Keshk
 Faraday’s Laws of Electrolysis – Second Law
Preventing of Corrosion
In order to produce one mole of metal, one two or three moles
of electrons must be consumed.

Faraday found that there was a certain charge associated with

one mole of electrons.

This amount of charge is now called the Faraday and is

equivalent to 96 487 (96 500) Coulomb.

1 Faraday = 96 500 Coulomb

107 Sherif Keshk

 Faraday’s Laws of Electrolysis – Second Law
Preventing of Corrosion
In order to produce one mole of metal, one two or three moles
of electrons must be consumed.

Ag+ (aq), Cu2+ (aq) and Cr3+ (aq) require 1, 2 and 3 moles of
electrons for discharge.

The quantity of electricity required will be:

Amount of charge = no. of moles of metal ions x charge

on an ion x 1 Faraday

Q = n z F and Q = I t

It = nzF
108 Sherif Keshk
 Faraday’s Laws of Electrolysis – Questions
Preventing of Corrosion
1. Calculate the number of mole of copper produced in an electrolytic
cell if a current of 5.0 A at a voltage of 6.0 V flows through a
solution of copper ions for 10 minutes.
2. Calculate the time taken to deposit 1.00 g of copper onto an object
that is placed in a solution of copper nitrate, Cu(NO3)2, and has a
current of 2.50 A flowing through it.
3. In an operating Hall-Héroult cell, a current of 150 000 A is used at
5.0 V. Calculate the mass of aluminium that would be produced if
this cell operates continuously for 1 day.
4. The electrolysis of a solution of chromium ions using a current of
2.2 A for 25 minutes produced 0.60 g of chromium. Calculate the
charge on the chromium ion.
5. Calculate the masses of metal produced when 600 Faraday of
109
charge is used to reduce theSherif
ionsKeshk
of aluminium, silver and zinc.
 Faraday’s Laws of Electrolysis – Solutions
Preventing of Corrosion
1. n (Cu) = I t / z F
n (Cu) = 5.0 x 10 x 60 / 2 x 96 500
n (Cu) = 1.6 x 10-2 mol
That is, 1.6 x 10-2 mole of copper would be produced

2. t = n (Cu) z F / I
t = (1.00 / 63.5) x 2 x 96 500 / 2.5
t = 1216 seconds
That is, it takes 20 minutes 15 seconds to deposit 1.00 g of copper

3. n (Al) = I t / z F
n (Al) = 150 000 x 24 x 60 x 60 / 3 x 96 500
n (Al) = 44 767 mol
m (Al) = 44 767 x 27 = 1 208 705 g
That is, 1.2 tonne of aluminium is produced per day
110 Sherif Keshk
 Faraday’s Laws of Electrolysis – Solutions
Preventing of Corrosion
4. z (Cr) = I t / n F
z (Cr) = 2.2 x 25 x 60 / (0.60 / 52) x 96 500
z (Cr) = 2.96
As the charge on an ion is a small integer, it must be 3+.
The ion is Cr3+.

5. Q = m z F / M
As Q = 600 F, then m = 600 M / z
m (Al) = 600 x 27 / 3 = 5400 g
m (Ag) = 600 x 107.9 / 1 = 64 740 g
m (Zn) = 600 x 65.4 / 2 = 19 620 g
The same 600 F of charge would produce different masses of these
metals; 5.4 kg of Al, nearly 20 kg of Zn and almost 65 kg of Ag.

111 Sherif Keshk

 Corrosion kinetics
Corrosion Kinetics
• M → Mn+ + ne-
• Up to this point, we have dealt with
the thermodynamics of corrosion, i.e.
which combinations of conditions w=
ItM
results in anodic and cathodic regions nF
under equilibrium w = weight loss during corrosion (or weight
• Corrosion does not occur under gain during electroplating)
equilibrium conditions I = current in amps = iA
• Of interest is the corrosion kinetics, i = current density
i.e. the rate at which a metal A = area of corroding surface
corrodes t = time in seconds
• For each atom of a metal that M = atomic mass g/mole
participates in the oxidation reaction, n = number of electrons involved in the
n electrons need to get transported corrosion reaction
away
F = Faraday’s constant
• The weight of a metal that is lost due = 96,500 C/mol
to corrosion is given by Faraday’s law

112 Sherif Keshk

 Polarization
Corrosion Kinetics
• In an electrochemical half cell the
metal atoms are in a state of
equilibrium with its ions in solution
– There is an equilibrium exchange
current density i0 associated with
the transfer of electrons at the
standard emf potential E0 (or
V0) of the half cell
• There is an i0 and E0 associated with
the anodic and cathodic reactions
• However, potential differences
cannot be maintained in a conductive
metal, such as Zn
• There is a displacement of the
electrode potentials and currents Point A: -0.763V, 10-7A/cm2
from points A and B to C Point B: 0V, 10-10A/cm2
Point C: ~-0.5V, 10-4A/cm2
• This displacement of electrode
potentials is called polarization
icorr = 10-4A/cm2 is used in
Sherif Keshk CPR calculations
113
 Polarization

• Activation polarization: In a • Concentration polarization:

multistep electrochemical Corrosion reaction may result in a
reaction the rate is controlled by build up or depletion of the ions or
the slowest step. atoms that are required for a
corrosion reaction.

114 Sherif Keshk

 Passivity
Passivation
• Passivation is the loss of chemical reactivity in
presence of a environmental condition.
– The formation of surface layer of reaction products that inhibit further
reaction
• Oxide film theory: A passive film of reaction
products acts as a diffusion barrier.
• Adsorption theory: Passive metals are covered by
chemisorbed films of oxygen.
• Examples: Stainless steel, nickel alloys, titanium
and aluminum alloys passivate in certain
environments

115 Sherif Keshk

 Passivity
Polarization Curve for Passivation
• Initially, the potential of the metal increases with current density, i.e. the metal
undergoes active corrosion
• When potential reaches Epp the primary passive potential, current density
decreases, i.e., the corrosion rate also decreases
• In order to make the metal active again, there may need to be an externally
applied potential

116 Sherif Keshk

 Passivity
Galvanic Series
• The Standard emf series gives the
relative oxidation or reduction
behavior under standard
conditions.
• The Galvanic series ranks
materials on the basis of
corrosion behavior in sea water

Sherif Keshk 117

 Passivity
Types of Corrosion
• Though the basic corrosion mechanism is the same, i.e. the oxidation of a
metal due to transfer of electrons, conditions under which the process
occurs can be vastly different, and has lead to the classification of
corrosion into the following types:
– Uniform
– Galvanic
– Crevice
– Pitting
– Inter-granular
– Selective leaching
– Erosion/cavitation
– Stress corrosion
– Fretting

118 Sherif Keshk

 Types of Corrosion

• Uniform or general corrosion – Tin coating of steel in “tin” cans

– Most common in terms of • If the coating is scratched, the
weight loss and destruction, underlying steel rusts quickly
but also the easiest to control – Zinc coating of steel by galvanizing
• If the coating is scratched, the
• Galvanic or two metal corrosion
galvanic cell results in corrosion
– When dissimilar metals are in of the coating, but due to large
contact surface area of Zn, the CPR of
– Very sensitive to the relative the Zn coating is relatively low
areas of the anodic and
cathodic regions
– If the anodic region is small,
then corrosion can be rapid
and deep

119 Sherif Keshk

 Types of Corrosion
Crevice Corrosion
• Localized electrochemical corrosion
in crevices and under shielded
surfaces where stagnant solutions
can exist.
• Occurs under valve gaskets, rivets
and bolts in alloy systems like steel,
titanium and copper alloys.

Anode: M → M+ + e-
Cathode:O2 + 2H2O + 4e- → 4OH-
Rivet holes
• As the solution is stagnant, oxygen is
used up and not replaced.
• Chloride ions migrate to the crevice
to balance positive charge and form
metal hydroxide and free acid that
causes corrosion

Sherif Keshk 120

 Types of Corrosion
• Pitting
– Initiates at non-uniformities in
composition
– Growth of pit involves
dissolution of metal in pit
maintaining high acidity at the
bottom.
– Anodic reaction at the bottom
and cathodic reaction at the
metal surface.
– At the bottom of the pit

MCl + H2O → M(OH) + H+ + Cl-

– The presence of H+ at the

bottom of the pit pulls more M
into solution
– Some metals (stainless steel)
have better resistance than
others (titanium).

Sherif Keshk 121

 Types of Corrosion
• Intergranular corrosion
– Localized corrosion at and/or adjacent to highly reactive grain
boundaries resulting in disintegration.
– When stainless steels are heated to or cooled through sensitizing
temperature range (500-800C) chromium carbide precipitate along
grain boundaries.
– When exposed to corrosive environment, the region next to grain
boundaries become anodic and corrode.

Sherif Keshk 122

 Types of Corrosion
Stress Corrosion Cracking (SCC)
• Cracking caused by combined
effect of tensile stress and
corrosive environment
• Stress might be residual and
applied
• Only certain combination of alloy
and environment causes SCC
• Crack initiates at pit or other
discontinuity
• Crack propagates perpendicular
to stress
• Crack growth stops if either stress
or corrosive environment is
removed.

Sherif Keshk 123

 Types of Corrosion

Erosion corrosion
• Acceleration in rate of corrosion due
to relative motion between corrosive
fluid and surface.
• Pits, grooves, valleys appear on
surface in direction of flow.
• Corrosion is due to abrasive action
and removal of protective film.
Cavitation damage
• Caused by collapse of air bubbles or
vapor filled cavities in a liquid near
metal surface.
• Rapidly collapsing air bubbles
produce very high pressure (60,000
PSI) and damage the surface.
• Occurs at metal surface when high
velocity flow and pressure are
present.

Sherif Keshk 124

 Types
Types of Corrosion
of Corrosion
Selective leaching:
• Selective removal of one element of alloy by corrosion
• Example: Dezincification or selective removal of zinc from
copper and brasses
• Weakens the alloy as single metal might not have same
strength as the alloy
Fretting corrosion
• Occurs at interface between materials under load subjected
to vibration and sliding
• Metal fragments get oxidized and act as abrasives between
the surfaces

125 Sherif Keshk

 Corrosion Control

126 Sherif Keshk

 Corrosion Control
Corrosion
• It is the degradation of a material due to a reaction with its environment.
OR
• Process of Distruction of the material through chemical or electrochemical
attack by its environment.
• Slow process
• Measured in weight loss per unit time.
Classification:
1. Dry or Chemical Corrosion
2. Wet or Electrochemical corrosion
1.Dry or Chemical Corrosion
- Occurs due to chemical attack of by the environment such as dry gas.
- Occurs due to high temperature and without liquid phase.

127 Sherif Keshk

 Corrosion Control
• It is of two types:
a) Oxidation corrosion b) Corrosion by gases
(a)Oxidation Corrosion:
- It is due to direct attack of oxygen on metals.
- Oxygen molecules are attracted to the surface by Vander Wall Force
Mechanism:-
1. When temp increases the metal undergoes oxidation and losses e-

2M → 2M+n + 2ne-
Metal Ion
2. Electron are gained by the oxygen molecules forms oxide ions

nO2 + 4ne- → 2n O2-

Oxide Ion
3. Scale of metal oxide formed

2M + nO2 → 2M + 2n O2-
Metal Oxide

128 Sherif Keshk

 Corrosion Control
• Stable Corrosion: -Aluminium, Tin, Lead, Copper
• Non-stable corrosion:- Silver, Gold, Platinum
• Pilling – Bed Worth Ratio
Ratio of volume of oxide formed to the volume of metal
consumed.
(b)Corrosion by Gases
Carbon di-oxide, Chlorine, Hydrogen Sulphide, Sulphur di-oxide, Flourine
- Depends on chemical affinity b/w metal and the gas.

2. Wet or Electrochemical Corrosion

• Occurs when aqueous solution or liquid electrolytes are present
• Wet corrosion takes place in environments where the relative humidity exceeds 60
%.
• Wet corrosion is most efficient in waters containing salts, such as NaCl (e.g. marine
conditions), due to the high conductivity of the solution.
129 Sherif Keshk
 Corrosion
Mechanism OfControl
Electrochemical Corrosion

Anodic Reaction:
Dissolution of metal takes place.
• As result metal ions are formed with the liberation of free electrons.
M ↔ M+n + e-
Metal Ion
130
Sherif Keshk
Cathodic Reaction Corrosion Control
(i) Hydrogen Evolution :- Occurs usually in acidic medium
2H+ + 2e- ↔ H2 (g)
(ii) Oxygen Absorption :- occurs when solution is aerated sufficiently.
O2+ 4H+ + 4e- ↔ 2H2O (In acidic medium)
O2+ 4H+ + 4e- ↔ 4OH- (In basic medium)
Forms of Corrosion:
(a)Galvanic Corrosion:- When two different metals are present in
contact with each other in conducting medium e.g. Electrolyte

131 Sherif Keshk

(b) Concentration Cell Corrosion:-
• Same as Galvanic corrosion Contents
• Occurs when two different metals are exposed to different air conc.

(c) Pitting Corrosion:-

• Formed as a result of pit and cavities
• Localized attack and formed by cracking protective coating

132 Sherif Keshk

 Corrosion Control

(d) Stress Corrosion:

• Occurs in the presence of tensile stress and corrosive environment
• E.g. brass get corrode in traces of ammonia.

133 Sherif Keshk

 Factors Affecting Corrosion
1. Nature of the Metal Corrosion Control
2. Nature of the environment.

1. Nature of Metal
(i) Position in Galvanic Series:
If two metals are present in in electrolyte,
the metal with less reduction potential undergoes corrosion.
- Greater the difference faster the corrosion.

(ii) Over Voltage:

Due to high evolution of hydrogen, the rate is slow.

(iii) Area and Distance:

When anodic metal area is smaller than cathodic
area, rate of corrosion at anode is higher because of demand of
134
electron by cathodic area. Sherif Keshk
 Corrosion Control
(iv) Physical and Mechanical properties of Metal:
(a) Pure metals are more corrosion resistant.
(b) Smaller grain size metal have high solubility and corrosion.
(c) Uniform distribution of stress on metal reduces rate of corrosion.
(d) Passive metals shows higher corrosion resistance because of formation of
protective oxide film on their surface.
(e) Polycrystalline forms are more sensitive.

2. Nature of Environment
(i) Temperature: directly proportional
(ii) Humidity: faster in humid conditions
(iii) pH : If less than 7 rate is high. Al, Zn, Sn, Pb, and Fe are affected by both acid
and bases.
(iv) Impurities and Suspended Particles: When these will get dissolved in
moisture, provides electrolyte for conductivity and hence corrosion increases.

135 Sherif Keshk

 Corrosion Control
Corrosion Control:
1. Selection of metal and alloy:
- Using pure and noble metals
- Practically not possible because of low strength of pure metal
- Use of metal alloys which are homogeneous
2. Proper design of metal:
(i) Minimal contact with medium
(ii) Prevention from moisture
(iii) Adequate ventilation and drainage
(iv) Welding
(v) Avoid cervices b/w adjacent parts
(vi) Bend should be smooth
(vii) Bimetallic contacts should be avoided
(viii) Paint cathodic portion
(ix) Prevent uneven stress
136 Sherif Keshk
 Corrosion Control
3. Cathodic Protection:
Force the metal to be protected to behave like cathode.
(i) Sacrificial anodic protection:
- Metal to be protected from corrosion connected to more anodic
metal
- Commonly used metals Mg, Zn, Al and their alloys
(ii) Impressed current method:
- Direct current is applied in opposite direction to nullify the
corrosion current
- Converts the corroding metal from anode to cathode.

137 Sherif Keshk

 Corrosion Control
4. Modifying Environment
(i) Eliminating dissolved oxygen:
- De-aeration
- By using chemical substances like sodium sulphite and hydrazine. Also
called Deactivation.
(ii) Reducing Moisture:
- Dehumidification by using silica gels
(iii) Reducing Acidity:
- Neutralizing the acidic environment by adding lime, NaOH, Ammonia
- Commonly used in refineries
5. Protective coating:
- Application of coating
- Coating material should be chemically inert under particular temp and
pressure.
138 Sherif Keshk
 Corrosion Control
6. Use of corrosion Inhibitor

(i) Anodic Inhibitor:

- These are oxygen and oxidizing agent.
- They combine the anodic metal forming an oxide film which reduce
corrosion

(ii) Cathodic Protection:

- Organic inhibitors like amines, mercaptans, urea and thiourea
reduces the H ion diffusion by adsorption
- Mercury, arsenic and antimony deposits films at cathodic area
which raise the hydrogen over volume.
- Eliminating Oxygen from the medium by adding sodium sulphate
and hydrazine.
139 Sherif Keshk
 Corrosion Control

140 Sherif Keshk

 Corrosion Control
Protective Coating
Surface preparation for Coating:
1. Cleaning:
- To prepare for suitable condition
- Removing contaminants to prevent detrimental reaction product
- E.g. de-greasing, sand blasting, vapour degreasing, pickling and alkaline
cleaning.
2. Solvent Cleaning:
- Must be non-inflammable and nontoxic
- Trichloro trifluoroethane which has low toxicity are costlier
- Vapour de-greasing is economical and advantageous because of
continuous cleaning with small quantities of solvent.
3. Electrolyte Pickling:
- Provides better and rapid cleaning by increasing hydrogen evolution
resulting in agitation and blasting action
- Sand blasting is mechanical cleaning.
141 Sherif Keshk
 Corrosion Control
4. Alkaline Cleaning:
- Cheaper and less hazardous
- Used in conjunction with surface active (wetting) agent
- Ability depends on pH, rapidly decreases below 8.5
- Other abilities are rinsability, detergent properties, sequestering,
wetting etc.

5. Acid Cleaning
- Acid such as HCl, H2SO4, H3PO4 is very effective.
- 5-10% H2SO4 and HCl used to remove inorganic contaminants.
- Pickling are performed at high temp. (60 ̊C)
- Is effective for removal of grease, oil , dirt and rust.
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Methods of Application of Metallic Coating

1. Hot Dipping:
- Metal is kept in molten state and base metal is dipped into it.
- Used for producing a coating of low M.P
- E.G. Tinning (Tin coating on Iron)
- Process is followed by cooling the coating through a palm oil to
prevent oxidation of tin plate to its oxide.
- Palm oil layer is removed by alkaline cleansing agent.

2. Metal Cladding:
- The surface to be protected is sandwiched between two layers of
the coating metals and pressed between rollers.
- E.g. Alclad Sheeting– Plate of duralumin is sandwiched between
99.5%pure aluminum
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3. Electro Plating:
- Pure metal is made as cathode and base metal as anode.
- Electrochemically coat metal is deposited on base metal.
- This metal gives smooth, fine and uniform coating
- It depends on
(i) Temperature (ii) Current density (iii) Electrolyte Concentration
(iv) Nature of base metal (v) Time

4. Electroless Plating:
- Nobel metal is deposited catalytically on less noble metal by using reducing
agent without using electrical energy.
- Advantage over Electro plating
(i) More economical since no electricity required
(ii) Irregular shape can be plated uniformly
(iii) Plating on plastics can also be done
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5. Metal Spraying:
- Coating is applied by means of spraying device
- E.g. Aluminum is plated in this way on Aircrafts.

Chemical Conversion Coating

• These are formed on metal surface by chemical reaction b/w metal surface
and inorganic salt solution
• Coating base metal is converted into one of the resultant protective film.
• These films are insoluble, adherent, crystalline or amorphous in nature.
• Can be done in 3 ways
1. Phosphate coating
2. Chromate coating
3. Anodized coating
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1. Phosphate Coating
Corrosion Control
- Produced by chemical reaction b/w base metal and aq. H3PO4, Zn or
Fe or Mn Phosphate
- Phosphate coating are applied Iron, Steel, and Zinc
- Film formed on base metal after coating consist of Zn-Fe, Mn-Fe
Phosphates.

2. Chromate Coating

- Produced by dipping the base metal in Potassium chromate (acidic)

followed by immersion in neutral chromate bath.
- Resulting film consist of trivalent and hexavalent chromium.
- Used as base for paints, lacquers and enamels.

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3.Anodized Coating
- Formed by anodic oxidation process
- This is produced on non-ferrous metals like Al, Zn, Mg
- In this method base metal is made as anode
- Process is carried out by passing moderate direct current through a
bath in which the metal is suspended as anode.
- Coating are formed as a result of Progressive oxidation starting at
surface of base metal.

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