# Short Course _Thermodynmics

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of 18 Short Course _Thermodynmics

## Content

To be a partner of choice in corrosion
research
5/27/2014 1
Corrosion Electrochemistry :
Thermodynamics
 Thermodynamics deals with energy and its
changes in reactions.
 Reactions are viewed in terms of changes in free
energy.
 According to the first law of thermodynamics
energy can be neither created nor destroyed.
 The second law states that the free energy is
released from the system to surroundings in all
spontaneous changes.
 Corrosion reactions are spontaneous and are
governed by the laws of thermodynamics.
2
Tendency to corrode and rate of corrosion
Tendency to corrode is determined by the free energy difference
between a metal and its corrosion product.
Consider the reaction:

Rate of corrosion is determined by the size of the energy barrier, the free energy of
activation,
G 
For the forward reaction the temperature dependence is given by the
Arrhenius equation.

A= undefined constant; R= gas constant, T= absolute temperature

The relationship between free energy and equilibrium constant for a
reaction:
For a spontaneous transition from high energy (initial
state, G
i
) to low energy (final state, G
f
), the free energy
change

G = G
f
- G
i
< 0 (negative)

For a spontaneous reaction to occur, G must be
negative !

.

Mg + H2O + ½ O2 = Mg(OH)2 (G
o
= -596 kJ/mol)
Cu + H2O + ½ O2 = Cu(OH)2 (G
o
= -119 kJ/mol)
Au + 3/2H2O + ¾ O2 = Au(OH)3 (G
o
= +66 kJ/mol)
Nernst equation
The free energy change is related to electrical potential by
G=–nFE and G
Ø
=–nFE
Ø

Where, F=Faraday constant, 96,500 coulombs/mole; n
number of electrons transferred in the corrosion reaction and
E is the measured potential in volts.
Under standard conditions, G
Ø
= -nFE
o

From G = G
Ø
+ (RT)ln{[C][D]/{[A][B]},
we have
- nFE= -nFE
o
+ (RT)ln{[C][D]}/{[A][B],
the Nernst equation:
E= E
o
- (RT/nF)ln{[[C][D]}/{[A][B]}
This is one of the most fundamental equations in
corrosion science and engineering.

Under standard conditions: T=298k, R=8.3143 J (mol k)
-1
,
ln X = 2.3 logX
Nernst equation can be written as
E= E
o
- (0.059/z)log{[Products]/[Reactants]}

E is the non equilibrium potential generated by the corrosion
reaction;
[reactants] = concentration of reactants and
[products] = concentration of products.

Standard Electrode Potentials
The potential difference between the anode and cathode can be
measured by voltage measuring device. The absolute potential of the
anode and cathode cannot be measured directly.
To define a standard electrode, all other potential measurements are
made against the standard electrode. If the standard electrode
potential is arbitrarily set to zero, the potential difference measured
can be considered as the absolute potential.

Standard Hydrogen Electrode
The half-cell in which the hydrogen reaction takes place is called the
standard hydrogen electrode, often abbreviated SHE
Electrode Standard Electrode Potential,E
o
V
(SHE)
Au
3+
+ 3e
-
= Au + 1.50
Cl
2
+ 2e
-
= 2Cl
-
+ 1.360
O
2
+ 2H
+
+ 2e = H
2
O + 1.228
Br
2
+ 2e = 2Br
_
+ 1.065
Ag
+
+ e = Ag + 0.799
Hg
2
2+
+ 2e = 2Hg + 0.789
Fe
2+
+ e
_
= Fe
3+
+ 0.771
I
2
+ 2e = 2I
-
+ 0.536
Cu
+
+ e = Cu + 0.520
Cu
2+
+ 2e = Cu + 0.337
2H
+
+ 2e = H
2
0.000 (by definition)
Pb
2+
+ 2e = Pb – 0.126
Sn
2+
+ 2e = Sn – 0.136
Ni
2+
+ 2e = Ni – 0.250
Fe
2+
+ 2e = Fe – 0.440
Cr
3+
+ 3e = Cr – 0.740
Zn
2+
+ 2e = Zn – 0.763
Al
3+
+ 3e = Al – 1.663
Mg
2+
+ 2e = Mg – 2.370
Na
+
+ e = Na – 2.714
The logical choice
is for all reactants
in their standard
states and
potentials for this
condition are
described as
standard
electrode
potentials, E
o
.
Noble
or
Cathodic
Active
or
Anodic
Standard Electrode Potential
The potential difference measured between metal, M, and the
hydrogen electrode, under well defined conditions.
Example:
Iron corrodes in a solution of H+
(a) iron dissolves: half-cell reaction: Fe = Fe
2+
+ 2e-
(b) hydrogen gas formed: half-cell reaction: 2H
+
+ 2e- = H
2

(c) overall reaction: corrosion reaction: Fe + 2H
+
= Fe
2+
+ H
2
Substituting into Nernst equation:
E=E
o
- (0.059/2)lg{[Fe
2+
][H
2
]/[Fe][H
+
]
2
}

The terms [H
+
] and [H
2
] have been made to 1, [Fe] can be
approximated as unity, so under standard conditions, [Fe
2+
] = 1 M,
E= E
o

The measured potential difference is the electrode potential of the
iron under standard conditions.

Points to note: if the measured potential is positive
dG=-nFE < 0, spontaneous reaction.
E
o
= + 0.44 V, dG is negative, iron corrodes spontaneously in acid.

The Basic Wet Corrosion Cell

Four essential components of a corrosion cell
The Anode
The Cathode
The Ionic Conductor (electrolyte)
The Metallic Conductor (electrical connection)

(1) The anode
The anode corrodes by loss of electrons: M= M
z+
+ ze-
Anodic reaction
Oxidation reaction
Electron generation
(2) The cathode
The cathode does not corrode. Most important cathodic reactions:
(i) pH < 7 2H
+
+ 2e- =H
2

(ii) pH > = 7 2H
2
O + O
2
+ 4e- = 4OH
-

Other cathodic reactions are possible Depending on the environment.
Cathode
Cathodic reaction
Reduction reaction
Electron consumption
3) An electrolyte (ionic conductor)
a solution conducting electricity
(4) Electrical connection
the anode and cathode in a corrosion cell must be in electrical contact.
Difference in free energies between the anode and the cathode
produces electrical potential which is the driving force for corrosion
reaction.
Current: flow of electrons; Corrosion current: corrosion rate
Points to note: all aqueous corrosion reactions can be thought of in
terms of the simple corrosion cell.
Separation of anode and cathode
permanent separation
differential aeration
random distribution
Cell Potential
Dry batteries and Daniell cell
Symbolism of a corrosion cell: Zn | Zn
2+
|| Cu
2+
| Cu

Zn electrode on the left, immersed in Zn
2+
ions; Cu electrode on
the right, immersed in Cu
2+
ions. Ionic species are separated by || .
Zinc is the anode, copper is the cathode.

The two half-cell reactions are:
Zn = Zn
2+
+ 2e-
Cu
2+
+ 2e- = Cu
The overall reaction Zn + Cu
2+
= Zn
2+
+ Cu
To find the theoretical cell potential, using Nernst equation for each
of the half-cell reactions:
E(Zn/Zn
2+
)=Eo(Zn/Zn
2+
) - (0.059/2)lg[Zn
2+
], and
E(Cu
2+
/Cu)=Eo(Cu
2+
/Cu) - (0.059/2)lg{1/[Cu
2+
]}.
So the cell potential
E(cell)=Eo(cell) - (0.059/2)lg{[Zn
2+
]/[Cu
2+
]}

Convention:
do not use oxidation potential
do use reduction potential
E
o
(cell) = E
o
(cathode) + E
o
(anode)

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