1. Corrosion is the unwanted oxidation of a metal.

2. Oxidation of all Metals in general is called corrosion Oxidation of All Metals is called Corrosion

3. Whereas, oxidation of iron metal specifically, is usually called rusting. Oxidation of All Metals is called Corrosion Oxidation of Iron Metal is called Rusting

4. Looking at the standard reduction table, (click) near the bottom, we see that (click) iron, zinc, and aluminum are all near the bottom right,

5. (click after fading)

6. They all have high oxidation potentials Ox. Pot. = +0.45 V Ox. Pot. = +0.76 V Ox. Pot. = +1.66 V

7. Which means they are strong reducing agents, Ox. Pot. = +0.45 V Ox. Pot. = +0.76 V Ox. Pot. = +1.66 V Strong Reducing Agents

8. And are readily oxidized. Ox. Pot. = +0.45 V Ox. Pot. = +0.76 V Ox. Pot. = +1.66 V Strong Reducing Agents Readily Oxidized

9. But notice that aluminum and zinc don’t seem to corrode like the iron in steel does! The iron in steel show obvious corrosion, but aluminum and zinc do not

10. It’s because aluminum and zinc both readily oxidize (click), 4Al (s) + 3O 2(g)  2Al 2 O 3(s) 2Zn (s) + O 2(g)  2ZnO (s)

11. To form hard oxide coatings that are difficult for more oxygen to penetrate. 4Al (s) + 3O 2(g)  2Al 2 O 3(s) 2Zn (s) + O 2(g)  2ZnO (s) Hard oxide coatings that are difficult to penetrate

12. So these metals will not corrode much under normal atmospheric conditions, even when they get rained on. 4Al (s) + 3O 2(g)  2Al 2 O 3(s) 2Zn (s) + O 2(g)  2ZnO (s)

13. But iron, or Fe solid is different than other metals like aluminum and zinc, Iron (Fe (s) ) is different than these other metals

14. It’s oxide does not form an impenetrable coating. Iron (Fe (s) ) is different than these other metals It’s oxide does not form an impenetrable coating. Instead it can flake off, exposing a fresh iron suface to oxidizing agents.

15. Instead it can flake off, exposing a fresh iron surface to oxidizing agents. Iron (Fe (s) ) is different than these other metals It’s oxide does not form an impenetrable coating. Instead it can flake off, exposing a fresh iron surface to oxidizing agents.

16. Steel is mainly iron. Steel is made up primarily of iron. Coating steel with tin will protect its surface and prevent oxidation of the iron in it. An example is a tin can.

17. Coating steel with tin protects it’s surface and prevents oxidation of the iron that’s in it. This is why tin cans, which are tin coated steel, are normally good for keeping food without rusting. Steel is made up primarily of iron. Coating steel with tin will protect its surface and prevent oxidation of the iron in it. An example is a tin can.

18. However, if the cans are dented or pitted, the tin coating is broken and water can collect in the indentations. Now rusting occurs relatively quickly

19. We’ll consider a piece of iron where a dent or a crack in a painted surface makes it easy for (click) a water droplet to come in contact with the iron surface. iron Water droplet

20. Here an iron atom can oxidize to an iron 2+ cation, which will dissolve in the water. iron Fe Fe 2+ e–e– e–e–

21. The equation for this oxidation is Fe (s)  Fe 2+ + 2e –. Further oxidation of iron here causes pits in the surface to increase in size. iron e–e– e–e– Fe 2+ Fe (s)  Fe 2+ + 2e –

22. What is formed here is an electrochemical cell. (click) Because oxidation takes place here, (click) this is called the anode region. iron e–e– e–e– Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region

23. Oxygen from the air can come into contact with the edge of the water droplet and the iron metal surface. iron e–e– e–e– Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region O2O2

24. If we could take this section of our standard reduction table.

25. And split it right here,

26. We could add in this half-reaction. ½O 2(g) + H 2 O + 2e – ⇄ 2OH – … +0.40 ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40

27. It is the reduction of oxygen in the presence of water, to produce hydroxide ions. ½O 2(g) + H 2 O + 2e – ⇄ 2OH – … +0.40 ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Reduction of oxygen to form hydroxide ions

28. And its reduction potential is + 0.40 volts ½O 2(g) + H 2 O + 2e – ⇄ 2OH – … +0.40 ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Reduction of oxygen to form hydroxide ions

29. Here, oxygen combines with water and the electrons formed by the oxidizing iron, to form hydroxide ions. iron e–e– e–e– Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region O2O2 H2OH2O ½O 2(g) + H 2 O + 2e –  2OH – OH –

30. Which then dissolve in the water droplet. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH –

31. Because reduction has taken place near the point where the water, iron, and air meet, iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region

32. This is called the cathode region. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region

33. Looking at the solubility table, we see that Fe 2+ and OH – form a low solubility compound. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region Including Fe 2+ & Fe 3+

34. The Fe2+ ion formed at the anode can combine with the OH– ions formed at the cathode, iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region Including Fe 2+ & Fe 3+

35. To produce the precipitate Fe(OH) 2. iron Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – Cathode Region Fe 2+ OH – Fe(OH) 2(s) Including Fe 2+ & Fe 3+

36. Which can build up on the iron surface. iron Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – Cathode Region Fe(OH) 2(s) Including Fe 2+ & Fe 3+

37. What can happen is Oxygen from the air and water in the droplet can act to (click) further oxidize the iron in iron(II) hydroxide, iron Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – Cathode Region Fe(OH) 2(s) O2O2 H2OH2O

38. And through a fairly complex series of reactions, form what is called hydrated iron(III) oxide, which has the formula Fe2O3 dot x H2O. iron Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – Cathode Region Fe(OH) 2(s) O2O2 H2OH2O Fe 2 O 3 xH 2 O (s)

39. The “x” here means that variable numbers of water molecules can appear in this formula, depending on conditions. Fe2O3 dot x H2O, or hydrated iron(III) oxide is one of the major components of rust. iron Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – Cathode Region Fe 2 O 3 xH 2 O (s) Variable numbers of water molecules can occur in this formula

40. The process that caused the oxidation of iron that we just discussed takes place when neutral or basic water contacts iron. The anode half-rx (click) is the reverse of this one found lower on the reduction table. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Neutral or Basic Water

41. And the cathode half-reaction is the one we added to the table proceeding in a forward direction ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Cathode: ½O 2(g) + H 2 O + 2e –  2OH – Neutral or Basic Water

42. Adding the cathode half-reaction, with it’s E naught value of +0.40 Volts ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Cathode: ½O 2(g) + H 2 O + 2e –  2OH – Neutral or Basic Water E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V

43. To the anode half-reaction, which because it is the reverse of the one on the table, has an E naught value of (click) + 0.45 volts. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Cathode: ½O 2(g) + H 2 O + 2e –  2OH – Neutral or Basic Water E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V

44. Gives the overall redox equation: ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH –, which has an E naught value of 0.40 V plus 0.45 V, which is (click) positive 0.85 volts. The + value for the E° of the overall reaction, means this process is spontaneous. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Cathode: ½O 2(g) + H 2 O + 2e –  2OH – Neutral or Basic Water E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V

45. Now if the water contacting the iron is acidic, the cathode becomes this half-reaction up here. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Cathode: ½O 2(g) + 2H + + 2e –  H 2 O Acidic Water E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water

46. Which is ½O 2(g) + 2H + + 2e –  H 2 O. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water

47. The H+ here means acid is present ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O Acid is Present E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water

48. To determine the equation for the overall reaction and find it’s cell potential, we add (click) the half- reaction at the cathode, ½O 2(g) + 2H + + 2e –  H 2 O, with its E naught value of +1.23 volts, ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water

49. To the half-reaction at the anode, which is Fe (s)  Fe 2+ + 2e –, with its E naught value of +0.45 Volts ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water

50. And we get the overall reaction, ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O, and its E naught value is (click) 1.23 V + 0.45 V, which is equal to (click) positive 1.68 volts. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water

51. We see that the overall cell potential for the reaction in acidic water, 1.68 volts. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water

52. Is significantly higher than the overall cell potential in neutral or basic water, which is only 0.85 Volts. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water

53. This means there is a greater tendency for rusting of iron to occur when an acid is present, than when it is not. In acidic conditions, iron tends to rust more readily. ½O 2(g) + H 2 O + 2e – ⇄ 2OH –........……..+0.40 Anode: Fe (s)  Fe 2+ + 2e – Acidic Water Cathode: ½O 2(g) + 2H + + 2e –  H 2 O E° ½O 2(g) + H 2 O + 2e –  2OH – +0.40 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + H 2 O + Fe (s)  Fe 2+ + 2OH – +0.85 V In neutral or basic water E° ½O 2(g) + 2H + + 2e –  H 2 O +1.23 V Fe (s)  Fe 2+ + 2e – +0.45 V. ½O 2(g) + 2H + + Fe (s)  Fe 2+ + H 2 O +1.68 V In acidic water There is a greater tendency for rusting to occur when an acid is present.

54. Recall that any oxidizing agent above Fe(s) on the left, will oxidize solid iron. So if iron is in contact with any of these, this will increase iron’s tendency to corrode. Will Oxidize Fe (s)

55. This includes some common things like (clk) all acids, which contain H+, or hydrogen ions, (clk) sulphuric acid, with hydrogen and sulphate ions, (clk) nitric acid with hydrogen and nitrate ions, (clk) bromine and chlorine sometimes found in hot tubs and pools, (clk) acidified oxygen as we discussed before, and (clk) acidified hydrogen peroxide. All acids Sulphuric acid Nitric acid Acidified oxygen Chlorine Bromine Acidified hydrogen peroxide

56. Rusting occurs more quickly when salt or other electrolytes are present. (click) When salt or other electrolytes are added to water, iron Fe 2+ Fe (s)  Fe 2+ + 2e – ½O 2(g) + H 2 O + 2e –  2OH – OH – Salt or other electrolytes dissolved in the water increase conductivity and speed up corrosion of iron Anode Region Cathode Region

57. The conductivity is increased, so charged particles can move faster. iron Fe 2+ Fe (s)  Fe 2+ + 2e – ½O 2(g) + H 2 O + 2e –  2OH – OH – Salt or other electrolytes dissolved in the water increases conductivity and speed up corrosion of iron Anode Region Cathode Region

58. This speeds up the corrosion of iron. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Salt or other electrolytes dissolved in the water increases conductivity and speeds up corrosion of iron Cathode Region

59. It is also known that chloride ions present in salt solutions tend to eat through any protective films that may exist on the surface of iron, thus speeding up corrosion. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region Chloride (Cl – ) ions present in salt solutions tend to “eat” through any protective films on the iron surface

60. In wintertime, salt is spread on roads to melt ice and snow. Salty water that is splashed on vehicles can greatly increase the rate of rusting if they are unprotected. Thank you to Jacques from Beloeil, Canada for sharing this photo on Pixabay

61. We should also point out that as (click) temperature increases, the rate of corrosion also increases. All reactions increase in rate as the temperature increases. iron Fe 2+ Fe (s)  Fe 2+ + 2e – Anode Region ½O 2(g) + H 2 O + 2e –  2OH – OH – Cathode Region Increase temperature Rate of Corrosion

62. A type of corrosion we should be aware of is something called galvanic corrosion. Galvanic Corrosion

63. Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte

64. For example, let’s say we have two pieces of steel (which is mainly iron) (click) bolted together with bolts made of copper, or an alloy of copper. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Steel (iron) Cu Cu

65. We’ll add some saltwater, which is an electrolyte. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Steel (iron) Saltwater Cu

66. This establishes a type of electrochemical or galvanic cell, in which the metal lower on the reduction table, (click) the iron in this case, Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Steel (iron) Saltwater Cu

67. Becomes the anode, and undergoes (click) oxidation, which slowly causes the steel to corrode where the bolts contact it. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Steel (iron) Saltwater ANODE Fe (s)  Fe 2+ + 2e – Cu

68. Galvanic corrosion can also happen when a brass boat propeller is used in saltwater. If we take a closer look at brass (click) we see it is an alloy of mainly copper and zinc. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Saltwater ZnCu ZnCu Zn Cu ZnCu Brass Propeller

69. Because zinc is lower on the right side of the reduction table, (click) it is a stronger reducing agent and is more readily oxidized than copper. You may recall that zinc normally forms a protective oxide layer. However, saltwater tends to penetrate through this layer, making zinc more susceptible to oxidation. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Saltwater ZnCu ZnCu Zn Cu ZnCu Brass Propeller

70. Zinc atoms on the surface (click) oxidize to Zn 2+ ions Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Saltwater ZnCu ZnCu Zn Cu ZnCu Brass Propeller Zn 2+ Zn  Zn 2+ + 2e –

71. These ions will leave the zinc and dissolve in the surrounding water, Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Saltwater Cu ZnCu Brass Propeller Zn  Zn 2+ + 2e – Zn 2+ Cu

72. Gradually corroding the brass propeller. Galvanic Corrosion Galvanic corrosion results when two dissimilar metals are attached in the presence of an electrolyte Saltwater Cu ZnCu Brass Propeller Zn  Zn 2+ + 2e – Cu Zn 2+

73. To summarize,.