1. Chapter 23 Corrosion

2. Reading
Sec to 23-3, 23-5 and 23-6 (excluding the 23-6 subsections on “Passivation and Anodic Protection” and “Materials Selection and Treatment”)

3. Homework No. 11
Problems 23-6, 23-12

4. Introduction
Corrosion is the chemical interaction of materials with diverse environments.
Such interactions may impact the integrity of materials and thus their mechanical performance and physical properties.
Corrosion is observed in metals, ceramics and polymers:
Metals and ceramics  corrosion
Polymers  degradation

5. Corrosion
Chemical corrosion
Electrochemical corrosion

6. Chemical corrosion
Direct dissolution, i.e., a material dissolves in a corrosive liquid medium.
E.g., copper dissolving in water to form copper hydroxide.

7. Liquid metal attack – liquid metal attacking a solid first at high-energy locations, such as grain boundaries. If these regions continue to be attacked preferentially, cracks eventually grow.

8. Selective leaching – one particular element in an alloy being selectively dissolved, or leached, from the solid.

9. Electrochemical corrosion
Destructive and unintentional attack of a metal as it is exposed to an environment.
It is an electrochemical process and usually starts at the surface.
Very expensive problem: ~ 5% of an industrialized nation’s income is spent on corrosion prevention and maintenance or replacement.
Typical examples: rusting of car body panels, radiators, exhaust components, etc.

10. ©2003 Brooks/Cole, a division of Thomson Learning, Inc
©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning™ is a trademark used herein under license.
The components in an electrochemical cell: (a) a simple electrochemical cell and (b) a corrosion cell between a steel water pipe and a copper fitting.

11. Electrochemical corrosion involves oxidation
Oxidation reaction:
The metal becomes an n+ positively charged ion
The metal also looses its n valence electrons

12. Examples of oxidation reactions
The site where oxidation occurs is called the anode.
Oxidation is also known as anodic reaction.

13. Reduction reaction
The electrons generated from the oxidation of the metal atom must be transferred to become a part of another chemical species. This process is called reduction reaction.
If the corrosion occurs in an acid solution, this solution has a high concentration of H+ ions and the H+ ions are reduced :

14. The hydrogen electrode

15. Another reduction reaction
If metal is exposed to an acid solution with
dissolved oxygen, reduction occurs as

16. The water electrode

17. Yet another reduction reaction
Exposure to a neutral or basic aqueous solution with oxygen dissolved, the reduction becomes

18. The oxygen electrode

19. Still another reduction reaction
Any metal ions present in the solution in contact with the oxidizing metal, may also be reduced:
Electroplating

20. Metallic materials oxidize to form ions at different rates.
Composition cell
membrane, limits mixing of the 2 solutions.

21. Reduction process
The site where the reduction process occurs is called the cathode.
It is possible that two or more reduction reactions occur simultaneously

22. The complete electrochemical reaction is the sum of the anodic and cathodic reactions.
Each of the oxidation and reduction reactions are termed half-cell reactions.

24. Oxidation of Zn in an acid solution.

25. Rusting of iron
Corrosion (rusting) of iron in H2O with dissolved O2

26. It occurs in two stages: First stage:
Second stage:
Dissolves in corroding solution
Insoluble compound

29. If the Fe and Cu are connected electrically the following reactions will occur:
Anodic reaction (Fe corrodes):
Fe  Fe2+ + 2e-
Cathodic reaction (Cu deposits):
Cu2+ + 2e-  Cu
Reduction of the Cu occurs at the expense of the Fe:
Cu2+ + Fe  Cu + Fe2+
Current flows through the external circuit as the e­s generated in the Fe are used up by the Cu cell to reduce the Cu2+.

30. Composition cell
Negative end of open-circuit voltage is always at the anode.

32. Composition cell

33. Steel pipe
Steel pipe
Composition cell
Components of an electrochemical cell: (a) a simple electrochemical cell and (b) a corrosion cell between a steel water pipe and copper fitting.

34. Composition cells
Microscopic corrosion cells

35. Composition cell

36. In multiple alloys, one phase is more anodic than another.
Corrosion rates are higher in multiphase alloys.
Impurities in metals leads to precipitation of intermetallic phases and hence forms anodic and cathodic regions leading to corrosion.

37. The Standard EMF Series
Metals may be evaluated as to their tendency to oxidize when coupled to other metals immersed in solutions of their respective ions.
The half cell reaction of a metal electrode immersed in a 1M solution of its ions at TRT is termed standard half cell.
The measured cell voltages represent only differences in electrical potential.
Convenient to establish a reference half cell to which the other half-cells may be compared.
The reference cell arbitrarily chosen is the standard hydrogen electrode.

39. The Standard EMF Series
Membrane
H2 1 atmosphere
The reference cell consists of an inert Pt electrode immersed in a 1M solution of H+ ions.
The electrolyte is saturated with H2 gas bubbled through 1 atm pressure and 25ºC.
Note that Pt is not involve in the reaction.
Pt acts as a surface where H-atoms may be oxidized or H+ ions may be reduced

40. The Standard EMF Series
The electromotive force (emf) series is generated by coupling standard half-cells for various metals to the standard hydrogen electrode.
The ranking is done according to measured voltages.
The table represents the corrosion tendencies for the different metals.

42. Electrode potential of assumed cathode
– that of assumed anode
> 0
→ Assumption OK.
< 0
→ Assumed cathode is actually the anode.

43. Calculation of overall cell potential
The voltages on the emf Table correspond to half-cell reduction reactions and the voltages are read straight from the table:
This table can also be used for oxidation, but the direction of the reaction is reversed. In addition the sign of the voltage from the table is reversed:

44. The overall cell reaction:
The overall cell potential:
If Vo > 0  the reaction occurs spontaneously.
If Vo < 0  the reaction occurs spontaneously but the cell direction is just reversed.

45. Electrode potential of assumed cathode
– that of assumed anode:
> → Assumption OK.
Electrode potential of assumed anode
– that of assumed cathode:
< → Assumption OK.

47. Assumed anode (Zn) – assumed cathode (Cu)
= (-0.763) – (+0.337) = -1.1 V < 0
So assumption is OK.
Assumed anode is really the anode.

48. Composition cell
Membrane

49. Assumed anode (Zn) – assumed cathode (Fe)
= (-0.76) – (-0.44) = V < 0
So assumption is OK.
Assumed anode is really the anode.

50. Assumed anode (Cu) – assumed cathode (H)
= – 0
= V > 0
So assumption is wrong.
Assumed anode is actually the cathode.

51. Effect of concentration and temperature on cell potential (nonstandard solution)
emf series applies to pure metals at STP conditions in 1M solutions.
Use of alloy electrodes, and/or changing the solution concentration or temperature will alter the cell potential. In some cases the spontaneous cell reaction may be reversed.

52. The Galvanic Series
A more realistic and practical ranking of the metals and commercial alloys in seawater.
emf table was made using idealized conditions and it has limited usefulness.
In the galvanic series table, the alloys near the top are cathodic and unreactive, but those at the bottom are anodic.
No voltages are provided.
Comparison of the two tables shows a high degree of correspondence with respect to the relative positions of the pure base metals.

55. Corrosion rate
n electrons are released for every metal atom that gets corroded. Since each electron has a charge of 1.6 X C, a charge of n(1.6 X 10-19) C must flow from the anode to the cathode for every metal atom that gets corroded. For one mole of metal atoms to get corroded, a charge of Nn(1.6 X 10-19) C, where N = Avogadro’s number. If the mass of a mole of metal atoms is M, that means a charge of Nn(1.6 X 10-19) C is needed for the corrosion of M grams of the metal.

56. 1 Faraday (F) = charge of a mole of electrons, i.e., N(1.6 X 10-19) C

57. Current (in Ampere) = charge (in Coulomb) per unit time (in second)
If the current is set at 1 A, 1 C flows in 1 s.
For a charge of Nn(1.6 X 10-19) C (i.e., n F) to flow in 1 s, the current needs to be Nn(1.6 X 10-19) A.
For a charge of Nn(1.6 X 10-19) C (i.e., n F) to flow in 1 min (60 s), the current needs to be Nn(1.6 X 10-19)/60 A.

58. Electroplating rate
n electrons are received by the cathode for every M (anode) atom that gets electroplated.

59. .

60. Concentration cell

61. Concentration cell

62. Concentration cell

63. Concentration cell

64. Concentration cell

65. Concentration cell

66. Concentration cell

67. Concentration cell

69. Stress cell

70. Stress cell

71. Stress cell

72. Grain
Boundary
Grain boundary
(anode)
Cartridge Brass

73. ©2003 Brooks/Cole, a division of Thomson Learning, Inc
©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning™ is a trademark used herein under license.
Photomicrograph of intergranular corrosion in a zinc die casting. Segregation of impurities to the grain boundaries produces microgalvanic corrosion cells (x50).

74. Corrosion protection

75. Sacrificial anode (Zn)
Galvanized steel
Sacrificial anode (Zn)
Isolation to outside world

76. ©2003 Brooks/Cole, a division of Thomson Learning, Inc
©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning™ is a trademark used herein under license.
Zinc-plated steel and tin-plated steel are protected differently. Zinc protects steel even when the coating is scratched, since zinc is anodic to steel. Tin does not protect steel when the coating is disrupted, since steel is anodic with respect to tin.

78. Sacrificial anode

79. Cathodic protection