03.Thermodynamics in Corrosion Engineering

Thermodynamics in
Corrosion Engineering
Utility of Thermodynamics in Electrochemistry
Thermodynamic considerations allow the
determination of whether a reaction can
occur spontaneously
If metal dissolution is unfavorable
thermodynamically in a given set of
circumstances – the job of the corrosion
engineer is done
Example: Copper in pure deoxygenated water
To relate your thermodynamic knowledge
with the thermodynamics of corrosion-related
To describe the need for and characteristics
of reference electrodes
To describe the origin, use, and limitations of
electrochemical phase diagrams (such as
Pourbaix diagram)
Free Energy: Driving Force of a Chemical Reaction
Relation of ΔG and emf
G  EnF
ΔG is in Joules
E is emf in volts
n is the number of electrons
involved in the reaction
F is the Faraday (96500
The larger the value of E for any cell – more is the
tendency for the overall cell reaction to proceed
Ecell = Ecathode - Eanode
The Nernst Equation
General Reaction for a Galvanic Cell
lL  mM        qQ  rR       
Nernst Equation:
RT aQ .aR ...
EE 
ln l m
nF aL .aM ...
Half Cell Potential
When a metal M is immersed in an aqueous
electrolyte, it acquires a certain potential. If
the activity of the metal ions M++ in the
aqueous environment is unity, then the
acquired potential is known as standard
potential φ0
Potential of each electrode can be calculated
using Nernst equation
Example: Zinc Electrode
 Zn
 2e  Zn
  Zn 
2 F Zn
Hydrogen Electrode
It is assumed arbitrarily
that the standard
potential for the
following reaction is
equal to zero at all
Standard Hydrogen Electrode (SHE)
The potential of the electrode equals zero if the
hydrogen ion activity and the pressure of hydrogen
gas in atmospheres are both unity. This is the
standard hydrogen potential
The half - cell potential for any electrode is equal to
the emf of a cell with the standard hydrogen
electrode as the other electrode.
The half - cell potential for any electrode expressed
on this basis is said to be on the normal hydrogen
scale or on the standard hydrogen scale ,
sometimes expressed as φH or φ ( S.H.E. )
Convention of Signs and Calculation of EMF
Zn 2  2e   Zn
 0  0.763 V
Zn - 2e   Zn 2
 0  0.763 V
It was agreed at the 1953 meeting of the
International Union of Pure and Applied Chemistry
that the reduction potential for any half - cell
electrode reaction would be called the potential
Pt; H2,
H ,
Zn ;
Zn Cell
Ecell = Ecathode – Eanode= ??
Reference Half Cells
It is not always convenient to have a
hydrogen electrode in the laboratory
Other reference half-cells (reference
electrodes) have been introduced.
Calomel reference electrode
Ag-AgCl half cell
The Saturated Copper-Copper Sulfate half cell
Calomel Reference Electrode
Ag-AgCl Reference Electrode
Cu-CuSO4 Half Cell
Number Line for Potential Conversion Among
Different Reference Electrode Scales
Oxygen Electrode
Oxygen Electrode and Differential Aeration Cell
Consider two O2 electrodes:
one in contact with O2 at 1 atm
other in contact with O2 at 0.2 atm
Oxygen Electrode and Differential Aeration Cell
The reaction is not thermodynamically possible as
Thus, the electrode 1 is cathode electrode 2 the
In a differential aeration cell, the electrode in lower
O2 pressure acts as the anode and the one in higher
O2 pressure acts as the cathode
EMF Series
All metals have been arranged in a series according
to their standard potential (φ0) values.
The more positive value corresponds to noble
metals and the more negative value corresponds to
more reactive metals (when arranged according to
reduction potential)
Of the EMF series – if two metals make up a cell,
the more active metal acts as the anode and the
more noble metal of the two will act as cathode
EMF Series
Problems with EMF Series
In real situation, the activities of the metal ions in equilibrium with
the respective metals usually do not equal unity
The position of a metal in the EMF series with respect to another
metal may change because of complex formation as is the case
with tin (Sn) and steel (Fe)
Alloys are not included in the EMF series
In oxidizing environment, some metals undergo passivation and
are known as active-passive metals. Transition metals usually
show passive behaviour in aerated aqueous environment. This
dual position of some metals is not reflected in the EMF series.
Galvanic Series
Galvanic series is an arrangement of both
metals and alloys according to their actual
measured potentials in a particular
environment. There would be one Galvanic
series for each environment
Metals and alloys showing active-passive
behaviour are listed in both active and
passive states.
Galvanic Series in Seawater
Pourbaix Diagram
Marcel Pourbaix developed potential-pH diagrams to
show the thermodynamic state of most metals in
dilute aqueous solutions
With pH as abscissa and potential as ordinate,
these diagrams have curves representing chemical
and electrochemical equilibria between metal and
aqueous environment
These diagrams ultimately show the conditions for
immunity, corrosion or passivation.
Simplified Pourbaix Diagram for Iron
Pourbaix Diagram for Iron
Pourbaix Diagram for Iron at 25°C
Benefits of Pourbaix Diagram
Pourbaix diagrams offer a large volume of thermodynamic
information in a very efficient and compact format.
The information in the diagrams can be beneficially used to
control corrosion of pure metals in the aqueous environment
 By altering the pH and potential to the regions of immunity and
passivation, corrosion can be controlled. For example, on
increasing the pH of environment in moving to slightly alkaline
regions, the corrosion of iron can be controlled
 Changing the potential of iron to more negative values eliminate
corrosion, this technique is called cathodic protection.
 Raising the potentials to more positive values reduces the
corrosion by formation of stable films of oxides on the surface of
transition metals
Limitations of Pourbaix Diagrams
These diagrams are purely based on
thermodynamic data and do not provide any
information on the reaction rates
Consideration is given only to equilibrium conditions
in specified environment and factors, such as
temperature and velocity are not considered which
may seriously affect the corrosion rate
Pourbaix diagrams deal with pure metals which are
not of much interest to the engineers
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