1. In the name of GOD

2. What is Corrosion? Reaction of a metal with its environment
Aqueous corrosion
reaction with water (usually containing dissolved ions)
High temperature oxidation
reaction with oxygen at high temperature
High temperature corrosion
reaction with other gases

3. Examples of Corrosion Rusting of steel
corrosion product (rust) is solid but not protective
Reaction of aluminium with water
corrosion product is insoluble in water, so may be protective
Burning of magnesium in air
high temperature oxidation

4. Acids and Bases
An acid is a substance that produces excess hydrogen ions (H+) when dissolved in water
examples are HCl, H2SO4
A base is a substance that produces excess hydroxyl ions (OH-) when dissolved in water
examples are NaOH, KOH

5. Corrosion of Zinc in Acid
Zinc dissolves with hydrogen evolution
Zn + 2HCl  ZnCl2 + H2
Zinc known as a base or active metal
One atom of zinc metal
plus two molecules of hydrogen chloride (hydrochloric acid)
reacts to form goes to
one molecule of zinc chloride
plus one molecule of hydrogen gas

6. Corrosion of Platinum in Acid
Platinum does not react with acids
Platinum is known as a noble metal

7. Connection of Platinum to Zinc
Zinc and platinum not connected, no reaction on platinum
Zinc and platinum connected, current flows and hydrogen is evolved on platinum
electrons Zn Pt
HCl
Zn  Zn2+ + 2e-
metal  metal ions + electrons
Zn + 2HCl  ZnCl2 + H2
metal + acid  salt + hydrogen
2H+ + 2e-  H2
hydrogen ions + electrons  hydrogen gas

8. Connection of Platinum to Zinc
Zn + 2HCl  ZnCl2 + H2
But we can separate metal dissolution and hydrogen evolution
Zn  Zn2+ + 2e-
2H+ + 2e-  H2
These are known as electrochemical reactions
One atom of zinc metal
two electrons in the metal
one zinc ion in solution
Reactions that involve both chemical change and the transfer of charge

9. External Current Applied to Platinum in Acid
Hydrogen evolved on
negative electrode
2H+ + 2e-  H2
Oxygen evolved on
positive electrode
2H2O  O2 + 4H+ + 4e- Pt
HCl + -
Acid - chemical species that produces hydrogen ions in water
Overall reaction
2H2O  2H2 + O2

10. External Current Applied to Platinum in Alkali
Hydrogen evolved on
negative electrode
2H2O + 2e-  H2 + 2OH-
Oxygen evolved on
positive electrode
4OH-  O2 + 2H2O + 4e-
The product of [H+] times [OH-] is 10-14, so in pure water both [H+] and [OH-] are This leads to the concept of pH, which is defined as -log[H+]
Hence pH = 0 is strong acid, 7 is neutral, and 14 is strong alkali Pt
NaOH + -
Alkali - chemical species that produces hydroxyl ions (OH-) in water
Note that H+ and OH- are in equilibrium in water:
H2O  H+ + OH-
Overall reaction
2H2O  2H2 + O2

11. External Current Applied to Platinum
A piece of metal in the solution
Hydrogen evolution at one electrode
2H+ + 2e-  H2 (acids)
or H2O + 2e- H2 + 2OH- (alkalis)
Oxygen evolution at the other electrode
2H2O  O2 + 4H+ + 4e- (acids)
or OH- O2 + 2H2O + 4e- (alkalis)

12. Faraday’s Law
Charge is related to mass of material reacted in and electrochemical reaction:
2H+ + 2e-  H2
Two hydrogen ions
To produce one molecule of hydrogen gas
React with two electrons

13. Faraday’s Constant
One mole of hydrogen ions (1 g) contains Avogadro’s number (6 1023) ions
Hence electrons will react with each mole of hydrogen ions
Charge on the electron is 1.6  C
Hence one mole of ions requires C
This is known as Faraday’s constant

14. Faraday’s Law

15. Electrodes
Electrodes are pieces of metal on which an electrochemical reaction is occurring
An anode is an electrode on which an anodic or oxidation reaction is occurring
A cathode is an electrode on which a cathodic or reduction reaction is occurring

16. Anodic Reactions Examples Zn  Zn2+ + 2e- zinc corrosion
Fe  Fe2+ + 2e- iron corrosion
Al  Al3+ + 3e- aluminium corrosion
Fe2+  Fe3+ + e- ferrous ion oxidation
H2  2H+ + 2e- hydrogen oxidation
2H2O  O2 + 4H+ + 4e- oxygen evolution
Oxidation reactions
Produce electrons

17. Cathodic Reactions Examples O2 + 2H2O + 4e-  4OH- oxygen reduction
2H2O + 2e-  H2 + 2OH- hydrogen evolution
Cu2+ + 2e-  Cu copper plating
Fe3+ + e-  Fe2+ ferric ion reduction
Reduction reactions
Consume electrons

18. Metal Ion Hydrolysis
Note that metal ions may react with water (a hydrolysis reaction)
e.g. Al3+ + 3H2O  Al(OH)3 + 3H+
or 2Al3+ + 3H2O  Al2O3 + 6H+
Note that in an electrochemical reaction, we have the same number of each atom on each side of the equation, and the same overall charge

19. Effect of Potential
Electrochemical reactions involve transfer of charge
Hence, we expect that the voltage of the metal with respect to the solution will affect electrochemical reactions
Voltage of metal with respect to solution is known as the electrochemical potential

20. Corrosion of zinc in acid
When zinc is placed in acid the metal will start to dissolve and hydrogen will start to be liberated according to the potential of the metal
Consider the anodic zinc dissolution reaction Zn  Zn2+ + 2e-

21. Corrosion of zinc in acid
If the potential is above the Corrosion Potential, then it will fall due to
production of electrons
Then the corrosion rate may be expressed as the corrosion current density, icorr
If the potential is below the Corrosion Potential, then it will rise, due to
consumption of electrons
Zn  Zn2+ + 2e-
Rate of Reaction
At the Corrosion Potential, Ecorr, we have a stable mixed equilibrium
As the reaction involves transfer of charge, the rate of reaction may be expressed as a current per unit area, or current density
Electrochemical Potential
icorr
Ecorr
2H+ + 2e-  H2
Current density
Rate of Reaction

22. How Fast will Corrosion Occur?
Corrosion kinetics
Concerned with the rates of corrosion reactions
Mixed potential theory:
The corrosion potential will be that potential at which the sum of all anodic (positive) and cathodic (negative) currents on the electrode is zero
Polarization
The change in potential that is caused by the passage of a current

23. Types of Polarization Activation Polarization
The polarization necessary for the electrochemical reaction to go at the given rate
Given by Tafel’s Law:
E = potential at current i
Eo = potential at current io
b = Tafel slope

24. E-log i and Evans Diagrams
Plot E against log |i|, then activation polarization gives a straight line
Eo and io for the cathodic reaction
Anodic reaction, Tafel slope is positive
Tafel slope expressed as mV per decade of current mV
log (-i2) - log (-i1)
Mixed equilibrium occurs when sum of all currents is zero
log |current|
Electrode Potential
Ecorr and icorr for the corrosion reaction
Cathodic reaction, Tafel slope is negative
Eo and io for the anodic reaction

25. Concentration Polarization
Additional polarization caused by drop in concentration of a reactant at the electrode surface
As concentration falls, more polarization is needed to make the current flow
Eventually, no more current can flow because no more reactant can reach the metal, and a limiting current is reached

26. Concentration Polarization
Oxygen reduction is often affected by concentration polarization
Rate of cathodic oxygen reduction without concentration polarization
log |current density|
Electrode Potential
Rate of cathodic oxygen reduction with concentration polarization
Limiting current density - rate of reaction limited by availability of oxygen at
the metal surface

27. Resistance Polarization
If there is a resistance between the anode and the cathode in a cell, then the current flowing through that resistance will cause a potential drop given by Ohm’s Law:
V = IR
This is important for paint films and for high resistance solutions

28. Resistance Polarization
Resistance Polarization causes potential of anode and cathode to differ due to potential drop across solution, hence corrosion current is reduced
log |current density|
Electrode Potential

29. Passivation
When a passive film is formed, this causes a marked drop in current density due to the resistance of the film and its effect as a barrier to diffusion
This effect is seen on the anodic curve

30. Passivation
The rate of corrosion will be critically affected by the cathodic curve
When a stable passive film has formed, the current has a steady, low value - the passive current density
Rapid rate of cathodic reaction leads to passivation, and low rate of corrosion
log |current density|
Electrode Potential
Lower rate of cathodic reaction leads to activity, and high rate of corrosion
Current falls as the passive film starts to form - the active-passive transition
But it may also lead to low rate of corrosion?
Very slow cathodic reaction leads to low rate of corrosion
Active corrosion gives normal activation polarization

31. Polarization Curves Iron in hydrochloric acid Electrode Potential
log |current density|
Electrode Potential
Anodic iron dissolution
Cathodic hydrogen evolution

32. Polarization Curves Iron in sulphuric acid
log |current density|
Electrode Potential
Anodic iron dissolution (with active-passive transition)
Oxygen evolution on passive film (or transpassive corrosion as metal is oxidised to a higher oxidation state)
Cathodic hydrogen evolution

33. Polarization Curves Iron in aerated neutral NaCl solution
log |current density|
Electrode Potential
Cathodic oxygen reduction
Anodic iron dissolution
Cathodic hydrogen evolution

34. Effect of pH on reaction rate
Consider hydrogen evolution reaction
2H e-  H2
The concentration of hydrogen ions will influence the rate of the reaction
As the hydrogen ion concentration is increased (i.e. the solution made more acid), so the rate of the reaction increases
Similarly the potential will influence the reaction - the more negative the potential the faster the reaction

35. Effect of pH and potential on rate of hydrogen evolution
Slower
Potential
Faster pH

36. Effect of pH on reaction rate
On platinum no metal dissolution will occur, but to balance the charge a reaction which creates electrons must occur
If the solution contains dissolved hydrogen, the reverse of the hydrogen evolution reaction can occur: H2 2H e-

37. Effect of pH on reaction rate
H2 2H e-
This reaction will go faster in alkaline solution (since H+ will be removed by H+ + OH-  H2O)
This reaction will go faster at more positive potentials (because electrons will be removed from metal)

38. Effect of pH and potential on rate of hydrogen oxidation
Oxidation Faster
Reduction Faster
Reduction Slower
Rates equal
Potential
Oxidation Slower
Electrochemical Equilibrium pH

39. Thermodynamic Equilibrium
2H e-  H2
The potential at which it occurs for a given solution composition is known as the equilibrium potential.
The concentrations of reactants controls the rates of the forward and reverse reactions and hence the equilibrium potential

40. Thermodynamic Equilibrium
∆Gº= -nFEº ∆G= -nFE
Nernst Equation:

41. The Nernst equation 2H+ + 2e- = H2 The Nernst equation gives
For 1 atm. hydrogen gas

42. The Pourbaix (E-pH) Diagram
2H2O = O2 + 4H+ + 4e-
Equilibrium potential
falls as pH increases
2.0
1.6
O2 is stable
1.2
2H+ + 2e- = H2 Equilibrium potential falls as pH increases
0.8
Potential
0.4
H2O is stable
H2 is stable
0.0
-0.4
-0.8
pH = - log [H+]
-1.2
-1.6 7 14

43. Pourbaix Diagram for Zinc
Equilibrium for
Zn(OH)2 + 2OH-  ZnO H2O
2.0
1.6
Equilibrium for
Zn2+ + 2OH-  Zn(OH)2
1.2
0.8
Zn(OH)2
stable
solid
Potential
0.4
Equilibrium for
Zn + 2OH-  Zn(OH)2 + 2e-
ZnO22-
stable in
solution
0.0
Zn metal stable
Zn2+ stable
in solution
Equilibrium for
Zn  Zn2+ + 2e-
-0.4
Equilibrium for
Zn + 4OH-  ZnO H2O + 2e-
-0.8
-1.2
-1.6 7 14

44. Pourbaix Diagram for Zinc
Corrosion possible with oxygen reduction
Potential 7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
Corrosion
Corrosion requires strong oxidising agent
Passivity
Corrosion
Corrosion is possible, but likely to be stifled by solid corrosion product
Corrosion possible with hydrogen evolution
Zn(OH)2
stable
solid
ZnO22-
stable in
solution
Corrosion is thermodynamically impossible
Zn metal stable
Zn2+ stable
in solution
Immunity

45. Pourbaix Diagram for Gold
Potential 7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2 C
Passivity
Gold can’t corrode with oxygen reduction or hydrogen evolution
Immunity
Gold metal stable

46. Pourbaix diagram for Copper
Will copper corrode in acid?
Potential 7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
Will copper corrode in neutral waters?
No - hydrogen evolution only occurs below the potential for copper corrosion
Cu oxides
stable
CuO22- stable in soln.
Cu2+ stable
in solution
Usually it will just passivate, but corrosion can occur in slightly acid solutions
Cu metal stable

47. Pourbaix Diagram for Iron
Will iron corrode in acid?
Will iron corrode in neutral waters?
Potential 7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
Will iron corrode in alkaline solution?
Yes - there is a reasonably wide range of potentials where hydrogen can be evolved and iron dissolved
Fe3+
Fe oxides
stable
Yes - although iron can form an oxide in neutral solution, it tends not to form directly on the metal, as the potential is too low, therefore it is not protective.
Fe2+ stable
No - iron forms a solid oxide at all potentials, and will passivate
Fe metal stable

48. Pourbaix diagram for Aluminium
Potential 7 14
1.2
0.8
0.0
0.4
-1.2
-0.4
-0.8
-2.4
-1.6
-2.0
Al3+
Al2O3
AlO2- Al

49. Limitations of Pourbaix Diagrams
Tell us what can happen, not necessarily what will happen
No information on rate of reaction
Can only be plotted for pure metals and simple solutions, not for alloys

50. Thanks for your attention