CHAPTER 2
CORROSION PRINCIPLES
Chapter Outlines
2.1
2.2
2.3
2.4
2.5
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Oxidation and Reduction Reactions
Standard Electrode Half- Cell Potentials
Standard EMF Series
Galvanic Cells With 1 Molar Electrolytes
Galvanic Cells Not 1Molar Electrolytes
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2.1 Oxidation and Reduction Reactions

In metal, corrosion process is normally
electrochemical @ electrochemistry
(a chemical reaction in which there is transfer
of electrons from one chemical species to
another)

2 reactions that occur during corrosion process:
i. Oxidation reaction
ii. Reduction reaction
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i.
Oxidation reaction @ anodic reaction

Definition:
the removal of one or more electrons from an
atom, ion or molecule
Equation:

M
Mn+
+
ne-
(in which M becomes an n+ positively charged ion and in the
process loses its n valence electrons; e- is used to symbolize
an electron)
Zn  Zn2  2e

Example:

Anode is the side at which oxidation takes place.
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ii.
Reduction reaction

Definition:
the addition of one or more electrons to an atom, ion or
molecule (because the electrons generated from each metal
atom that is oxidized must be transferred to and become a part
of another chemical species = reduction reaction)

Equation:
M+ + e-
M(n-1)+
(some metals undergo corrosion in acid solutions, which have a high
concentration of hydrogen (H+) and hydrogen gas (H2) is evolved)

Cathode is the side at which reduction occurs
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 There are 3 possibilities of reaction that can occur at
cathode (reduction):
 First possibilities
Cathodic half- cell reaction
Condition
Reaction
: 2H  2e  H2 (gas )
: if the electrolyte is an acid
solution
: hydrogen ions in the acid solution
will be reduced to hydrogen atom to
form diatomic hydrogen gas
 Second possibilities
Cathodic half- cell reaction
Condition
Reaction
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

O

4H

4e
 2H2O
:
2
: if the electrolyte also contain
oxidizing agent
: oxygen will combine with hydrogen
ions to form water molecules
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 Third possibilities


O

2H
O

4e

4(OH)
Cathodic half- cell reaction : 2
2
Condition
: if the electrolyte is basic or
neutral and oxygen is
present
Reaction
: oxygen and water molecules
will react to form hydroxyl
ions
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iii.


Overall Electrochemical Reaction
Consist of at least one oxidation (half reaction) and
one reduction (half reaction), and will be the sum of
them
Example:
(Zinc metal immersed in an acid solution)
At some regions on the metal
surface, zinc will experience
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oxidation or corrosion
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 Oxidation half reaction:
Zn  Zn2  2e
Since Zn is a metal and good electrical conductor, these electrons
may be transferred to an adjacent region at which the H+ ions are
reduced.
 Reduction half reaction:
2H  2e  H2 (gas )
 Total electrochemical reaction:
Zn  Zn2  2e
2H  2e  H2 (gas )
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Zn + 2H+Asyadi Zn2+ + H2 (gas)
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Zinc metal
Chemical reaction:
Zn + 2HCl
ZnCl2 + H2
Ionic form:
Zn + 2H+
Zn2+ + H2
Half- cell reaction:
Zn
Zn2+ + 2e- (oxidation)
2H+ + 2eH2 (reduction)
hydrochloric acid
Fig. Reaction of hydrochloric acid with zinc to produce hydrogen gas
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2.2 Standard Electrode Half- Cell Potentials
 Every metal has a different tendency to corrode in a
particular environment
 E.g. ‘zinc is chemically attacked or corroded by dilute
hydrochloric acid, whereas gold is not’
 Method for comparing the tendency for metals to form
ions in aqueous solution is to compare their half- cell
oxidation or reduction potentials (voltages) to a
standard hydrogen- hydrogen ion half- cell potential.
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Experimental Setup for the Determination of
Half- cell Standard Electrode Potentials
Experimental setup for the determination of the standard emf of zinc. In a beaker a
Zn Electrode is placed in a solution of 1 M Zn2+ ions. In the other there is a
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standard hydrogen reference electrode
consisting of a platinum electrode
immersed in a solution of 1 MH+ ions which contains H2 gas at 1 atm.
Standard Hydrogen Electrode
 Represent only differences in electrical potential and thus it is
convenient to establish a reference point/ reference cell to which
other cell halves may be compared.
 It consist of an inert platinum electrode in a 1M solution of
H+ ions, saturated with hydrogen gas that is bubbled through
the solution at a pressure of 1 atm and temperature of 25°C.
 The platinum itself does not take part in the electrochemical
reaction: it acts only as a surface on which hydrogen atoms
may be oxidized or hydrogen ions may be reduced.
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2.2 Standard EMF Series
 Electromotive force (EMF) series:
is generated by coupling to the standard
hydrogen electrode, std half- cells for
various metals and ranking them according to
measured voltage.
 Table 17.1- show the list of the standard halfcell potentials of some selected metals which
represents the corrosion tendencies for the
several metals
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Increasingly inert
(cathodic)
Increasingly active
(anodic)
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 those metals at the top (Au & Pt) --- are noble or
chemically inert
 Moving down --- metals become increasingly more
active (more susceptible to oxidation) (sodium &
potassium)
 The voltages --- are for the half- reactions
oxidation reaction: electron on the right hand side
M1
Mn+
+ ne-
V1º
reduction reaction: electron on the left hand side
(sign of the voltage changed)
M+ + e4/12/2015
M(n-1)+
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V2º
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 Overall cell potential ΔV° is:
ΔVcell° = V°
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1
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+ V°
2
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GALVANIC CELLS
Galvanic couple:
Two metals electrically connected in a liquid
electrolyte wherein one metal becomes anode and
corrodes, while the other acts as a cathode
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2.4 Galvanic Cells With 1 Molar Electrolytes
 Can be constructed with two dissimilar metal electrodes each
immersed in a 1M solution of their own ions
 The two solutions are separated by a porous wall to prevent
their mechanical mixing, and an external wire in series with a
switch and a voltmeter connects the two electrodes
 E.g.:
zinc electrode immersed in a 1 M solution of Zn2+ ions and
another of copper immersed in a 1 M solution of Cu2+ ions with
the solutions at 25°C
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A macroscopic galvanic cell with zinc and copper electrodes. When the switch is closed
and the electrons flow, the voltage difference between the zinc and copper electrodes
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is -1.10V. The zinc electrode is the anode of the cell and corrodes.
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Calculation of electrochemical potential of Zn- Cu galvanic cell
 From the Standard emf Series:
Zn
Zn2+ + 2eCu
Cu2+ + 2e Oxidation half- cell reaction:
Zn
(ANODE)
Zn2+ + 2e-
E° = -0.763
 Reduction half- cell reaction:
Cu2+ +
2e-
E° = -0.763 V
E° = +0.340 V
Cu
V°1
(CATHODE)
E° = -0.340
V°2
 Overall reaction (by adding):
Zn + Cu2+
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Zn2+ + Cu
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E°cell = V°1 + V°2
= -0.763 + (-0.340)
= -1.103 V
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Problem 1:
A galvanic cell consist of an electrode of zinc in a 1M ZnSO4
solution and another of nickel in a 1 M NiSO4 solution. The two
electrodes are separated by a porous wall so that mixing of the
solutions is prevented. An external wire with a switch connects the
two electrodes. When the switch is just closed:
(a)
(b)
(c)
(d)
(e)
(f)
At which electrode does oxidation occur
Which electrode is the anode of the cell?
Which electrode corrodes?
Write the equation for the half- cell reaction at the anode?
Write the equation for the half- cell reaction at the cathode?
What is the emf of this galvanic cell when the switch is just
closed?
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Answer:
(a)
(b)
(c)
(d)
(e)
(f)
Oxidation occurs at the zinc electrode since the zinc half- cell
reaction has a more negative E° potential of -0.763 V as
compared to -0.250 V for the nickel half- cell reaction.
The zinc electrode is the anode since oxidation occurs at the
anode
The zinc electrode corrodes since the anode in a galvanic cell
corrodes.
Zn
Zn2+ + 2eE° = -0.763V
Ni2+ + 2eNi
E° =+0.250V
The emf of the cell is obtained by adding the half- cell reactions
together:
Anode reaction:
Cathode reaction:
Zn
Zn2+
Ni2+ + 2e-
Overall reaction:Zn + Ni2+
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+
2eNi
Zn2+ + Ni
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E° = -0.763 V
E° = +0.250 V
E°cell = -0.513 V
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2.5 Galvanic Cells Not 1 Molar Electrolytes
 Most electrolytes for real corrosion galvanic cells are
not 1 M, but are usually dilute solutions that are
much lower than 1 M.
 If the concentration of the ions in an electrolyte
surrounding an anodic electrode is less than 1 M, the
driving force for the reaction to dissolve or corrode
the anode is greater since there is a lower
concentration of ions to cause the reverse reaction
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Nernst equation:
E = E° + 0.0592 log Cion
n
Where: E = new emf of half- cell
E° = standard emf of half- cell
n = number of electrons transferred
(for example, M
Mn+ + ne-)
Cion = molar concentration of ions
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Problem 2:
A galvanic cell at 25ºC consist of an electrode of zinc in a 0.10 M
ZnSO4 solution and another of nickel in a 0.05 M NiSO4 solution.
The two electrodes are separated by a porous wall and connected
by an external wire. What is the emf of the cell when a switch
between the two electrodes is just closed?
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Answer:

Zn
Ni

Half cell reactions:
Zn2+ + 2eNi2+ + 2e-
E° = -0.763V (ANODE)
E° = -0.250V (CATHODE)
Apply Nernst Equation:
Ecell = E° + 0.0592 log Cion
n
Anode reaction: EA = -0.763 V + 0.0592 log 0.10 = -0.793 V
2
Cathode reaction: Ec = - (- 0.250 V + 0.0592 log 0.05) = +0.289 V
2
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Emf of the cell (Ecell) = EA + EC
= -0.793V + 0.289 V
= -0.505 V
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