Thermodynamics in Corrosion Engineering

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Thermodynamics in Corrosion Engineering

• Lecture03-04

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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

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Objectives

• To relate your thermodynamic knowledge with the
  thermodynamics of corrosion-related
  electrochemistry
• 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)

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Free Energy Driving Force of a Chemical Reaction
Spontaneous
Spontaneous
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Relation of ?G and emf

• ?G is in Joules
• E is emf in volts
• n is the number of electrons involved in the
  reaction
• F is the Faraday (96500 C/equivalent)

The larger the value of E for any cell more is
the tendency for the overall cell reaction to
proceed Ecell Ecathode - Eanode
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The Nernst Equation
General Reaction for a Galvanic Cell
Nernst Equation
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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 f0
• Potential of each electrode can be calculated
  using Nernst equation

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Example Zinc Electrode
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Hydrogen Electrode

• It is assumed arbitrarily that the standard
  potential for the following reaction is equal to
  zero at all temperatures
• So

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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 fH or f ( S.H.E. )

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Convention of Signs and Calculation of EMF

• 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

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Pt H2, H, Zn2 Zn Cell

• Ecell Ecathode Eanode ??

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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

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Calomel Reference Electrode
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Ag-AgCl Reference Electrode
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Cu-CuSO4 Half Cell
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Number Line for Potential Conversion Among
Different Reference Electrode Scales
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Oxygen Electrode
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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

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Oxygen Electrode and Differential Aeration Cell

• The reaction is not thermodynamically possible as
  written
• Thus, the electrode 1 is cathode electrode 2 the
  anode.
• 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

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EMF Series

• All metals have been arranged in a series
  according to their standard potential (f0)
  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

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EMF Series
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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.

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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.

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Galvanic Series in Seawater
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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.

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Simplified Pourbaix Diagram for Iron
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Pourbaix Diagram for Iron
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Pourbaix Diagram for Iron at 25C
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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

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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