The document discusses corrosion and identifies oxidation-reduction reaction pairs present in corrosion situations. It lists the basic types of corrosion as uniform attack, galvanic coupling, and localized corrosion. Localized corrosion includes pitting, crevice corrosion, and corrosion fatigue. Corrosion occurs through oxidation and reduction reactions. Oxidation reactions involve the metal going into its ionic state and releasing electrons, while reduction reactions consume those electrons.

3. Corrosion Example: Zn + 2HCl  ZnCl 2 + H 2 Chlorine only peripherally involved Zn + 2H +  Zn 2+ + H 2

4. Example 2 Reactions Oxidation: (Anodic RXN) Zn  Zn 2+ + 2e -

5. Example 2 Reactions Oxidation: (Anodic RXN) Zn  Zn 2+ + 2e - Reduction: (Cathodic RXN) 2H + + 2e -  H 2

6. Example Oxidation: (Anodic RXN) Zn  Zn 2+ + 2e - Reduction: (Cathodic RXN) 2H + + 2e -  H 2 Key Principle - Rate of Reduction = Rate of Oxidation

8. All corrosion falls into Ox-Red pair groups Oxidation RXN (Free Electron): M  M + n +ne - (From metal to its ion)

9. All corrosion falls into Ox-Red pair groups Oxidation RXN (Free electrons): M  M + n +ne - (From metal to its ion) ie: Ag  Ag + + e - Al  Al 3+ + 3e - >>>Produces Electrons

10. Reduction Reactions (Consume electrons) Hydrogen Evolution: 2H + + 2e -  H 2

11. Reduction Reactions (Consume electrons) Hydrogen Evolution: 2H + + 2e -  H 2 Oxygen Reduction (acid): O 2 +4H + +4e -  2H 2 0

12. Reduction Reactions (Consume electrons) Hydrogen Evolution: 2H + + 2e -  H 2 Oxygen Reduction (acid): O 2 +4H + +4e -  2H 2 0 Oxygen Reduction (neutral or basic): O 2 + 2H 2 O + 4e -  4OH -

13. Reduction Reactions (Consume electrons) Hydrogen Evolution: 2H + + 2e -  H 2 Oxygen Reduction (acid): O 2 +4H + +4e -  2H 2 0 Oxygen Reduction (neutral or basic): O 2 + 2H 2 O + 4e -  4OH - Metal Ion Reduction: M 3+ + e -  M 2+

14. 5 Reduction Reactions (Consume electrons) Hydrogen Evolution: 2H + + 2e -  H 2 Oxygen Reduction (acid): O 2 +4H + +4e -  2H 2 0 Oxygen Reduction (neutral or basic): O 2 + 2H 2 O + 4e -  4OH - Metal Ion Reduction: M 3+ + e -  M 2+ Metal Deposition: M + + e -  M

15. Note: Reactions can be controlled from either side (OX/ RED). Example: Add oxygen gas to an acid  Oxygen reduction is available to consume electrons.

16. Note: Reactions can be controlled from either side (OX/ RED). Example: Add oxygen gas to an acid  Oxygen reduction is available to consume electrons.  Higher Rate of Oxidation

17. Note: Reactions can be controlled from either side (OX/ RED). Example: Add oxygen gas to an acid  Oxygen reduction is available to consume electrons.  Higher Rate of Oxidation  Acids with oxygen are worse than acids without.

18. Polarization: What controls rate of RXN Two Types 1. Activation Polarization 2. Concentration Polarization

20. Concentration Diffusion of reducing species controls rate

23. Types 2 . Galvanic Coupling -Dissimilar metals or environments create electrical potential -Will have anode and cathode

25. Terminology Anode Cathode Oxidized Reduced Active Passive

32. Galvanic Example Zn Anode Oxidized Active Pt Cathode Reduced Passive

33. Galvanic Potential Example Dry Cell Battery V cell = 1.5 Volts

34. Calculation of Cell Potential p.568: Table Table Pt 2+ + 2e -  Pt +1.2V Mg 2+ + 2e -  Mg -2.363V

35. Calculation of Cell Potential p.568: Table Table Pt 2+ + 2e -  Pt +1.2V Mg 2+ + 2e -  Mg -2.363V Actual Actual Mg  Mg 2+ + 2e - (oxidation) +2.363V Pt 2+ + 2e -  Pt +1.2V

36. Calculation of Cell Potential p.568: Table Pt 2+ + 2e -  Pt +1.2V Mg 2+ + 2e -  Mg -2.363V Actual Actual Mg  Mg 2+ + 2e - (oxidation) +2.363V Pt 2+ + 2e -  Pt +1.2V Total Total Mg + Pt 2+ + 2e -  Mg 2+ + 2e - + Pt +3.563V