1. Principles of Corrosion
Electrolytes
Oxidation and Reduction
Corrosion of Iron
Corrosion Protection Methods

2. Electrolytes
Electrolytes are aqueous solutions capable of conducting electricity
The solute is an ionized compound:
an acid such as HCl,
a base such as NaOH
a salt such as NaCl

3. Strong Electrolytes
Completely dissociate in solution, also referred to as 100% ionization
Strong electrolytes are very good conductors of electricity
Strong electrolytes cause metal to corrode
Examples include:
most soluble salts such as NaCl,
“strong” acids such as H2SO4, HCl
“strong” bases such as KOH

4. Oxidation and Reduction

5. Oxidation States of Elements
The oxidation state (or oxidation number) of an atom represents the number of electrons lost, gained, or unequally shared by that atom
We can determine the oxidation state of each element in a compound
We can determine the change in oxidation state of elements that lose or gain electrons in a reaction
This is known as an oxidation/reduction reaction
Redox reaction for short
See slide 12

6. Oxidation/Reduction An element is oxidized if it loses electrons.
Its oxidation number (charge) becomes more positive.
An element is reduced if it gains electrons.
Its oxidation number becomes more negative.

7. Oxidation States of Elements

8. Rules for Writing Oxidation Numbers
Any element has an oxidation number of zero.
The sum of oxidation numbers in a compound is zero.
The sum of oxidation numbers in a polyatomic ion equals the charge on the ion.
H is +1 except in metal hydrides where it is -1
O is -2 except in peroxides where it is -1 and OF2 where it is +1
Metals are always positive
In molecules the negative oxidation number is assigned to the most electronegative element

9. Verify the oxidation numbers of each element in the following compounds:
AlF3
NH3
NH4+
NH2-
Cl2
ClF3
ClF4+
ClF4-
HNO3

10. More examples to verify:
HNO3
NO3-
N3O4
Cu2SO4
Sn3(PO3)4
Co(NO3)2
KSCN

11. Consider the reaction 4Fe + 3O2  2Fe2O3
In a redox reaction when one element is oxidized and one is reduced. A redox reaction can be considered as two half reactions (an oxidation and a reduction)
Consider the reaction 4Fe + 3O2  2Fe2O3
Oxidation half of the reaction
Oxidation of iron: 4 Fe0  4Fe electrons
The iron gives up electrons and is oxidized.

12. Continuing with 4Fe + 3O2  2Fe2O3
Reduction half reaction of the reaction
Reduction of oxygen: 3O electrons  6O-2
The oxygen accepts electrons and is reduced.
The positively charged iron cations and negatively charged oxygen anions attract.
Formation of ionic compound 4Fe+3 + 6O-2  2Fe2O3

13. KMnO4 + KNO2 + H2SO4  MnSO4 + H2O + KNO3 + K2SO4
Write oxidation numbers and identify the element oxidized and the element reduced by writing each half of the reaction:
Mg + HCl  MgCl2 + H2
Fe + V2O3  Fe2O3 + VO
KMnO4 + KNO2 + H2SO4  MnSO4 + H2O + KNO3 + K2SO4

14. Not all reactions involve oxidations and reductions. Consider
NaCl + MgBr2  NaBr + MgCl2
In this reaction no element changes its oxidation state.
This is a double displacement reaction

15. Oxidizing and Reducing Agents
An oxidizing agent is an element that causes the oxidation of the other element.
Good oxidizing agents strongly attract electrons and have high electronegativity.
Oxygen and fluorine are the two strongest oxidizing agents
A reducing agent is an element that causes the reduction of the other element.
Good reducing agents weakly attract electrons and have low electronegativity.
Most metals are good reducing agents
Recall the electronegativity of an element measures its attractive force for electrons.

16. Good oxidizing agents Flourine (4.0)
Oxygen (3.5) Since oxygen is present in air metals exposed to air are usually found in the oxidized state.
Hydrogen (2.1) Hydrogen is found in acids, bases, and water.
Attracts electrons from metals

17. Oxidation of Metals
When a metal forms a compound it becomes positively charged.
It loses electrons and is oxidized.
The element that caused its oxidation is reduced.
The activity series for metals lists metals in order of their reactivity.
Elements likely to be oxidized are located near the top of the series.
They also make good reducing agents.

18. Activity Series for Metals

19. Oxidation of Metals
Metals above zinc are oxidized by the hydrogen in water to release hydrogen gas.
2 Na (s) + 2H2O (l) 2NaOH (aq)+ H2 (g)
Metals above nickel are oxidized by dissolved oxygen in water.
4 Fe (s) + 3 O2 (aq)  2 Fe2O3 (s)
Metals above hydrogen react with acids to release hydrogen gas.
Pb + H2SO4PbSO4 + H2

20. Corrosion of Metals
Iron is the most commonly used in industry due to its abundance and low cost.
It is usually mined in the form of iron(III) oxide Fe2O3.
It must first be refined to metallic iron.
The iron ore is reacted with carbon monoxide to produce elemental iron.
Fe2O3 + 3CO  2Fe + 3 CO2
Iron is reduced by carbon.

21. Corrosion of Iron
Iron will attempt to revert back to its natural oxidized state if an oxidizing agent is present.
Water must be present for corrosion to occur.
There are 3 basic oxidations that occur depending on:
acidity
alkalinity
see note next slide
presence of dissolved oxygen in water.

22. A Note on Alkalinity
Water alkalinity refers to the ability of water to neutralize acid, or H+ ions
The anions: carbonates CO32−, bicarbonates HCO3−, and hydroxide OH− contribute to alkalinity since they neutralize the H+ ions.

23. Corrosion of iron in acidic or neutral water.
Hydrogen evolution corrosion occurs in water of pH < 9. The oxidizing agent is H+1
In carbonic acid solution the reaction is
Fe (s) + H2CO3 (aq) FeCO3 (s) + H2 (g)
 In hydrochloric acid solution the reaction is
Fe (s)+ 2HCl(aq) FeCl2 (s) + H2 (g)

24. Corrosion of iron in acidic or neutral water.
The acid etches the metal.
It appears free of corrosion products but becomes thinner.
In these reactions hydrogen gas evolves.
The amount of corrosion can be estimated by the amount of hydrogen gas evolved.
To prevent this in boilers make sure alkaline solution is added to neutralize acids. Also in condensate.

25. Corrosion of iron in water containing dissolved oxygen.
Oxygen absorption corrosion. The oxidizing agent is dissolved O2
2Fe (s) + 2H2O (l) + O2 (aq)  2Fe(OH)2 (s)
If additional dissolved oxygen is readily available the further oxidation of iron(II) hydroxide occurs.
4Fe(OH)2 + O2 (aq)  2Fe2O3 (s)+ 4H2O (l)
This produces red rust that offers no protection of underlying metal.

26. Magnetite
If there is little dissolved oxygen present the iron(II) hydroxide, Fe(OH)2, is oxidized only partially to magnetite (black rust).
6Fe(OH)2 + O2  2FeO.Fe2O3 (magnetite) + 6H2O
The magnetite forms a durable coat that protects the underlying metal from further corrosion.
This reaction is beneficial.

27. Oxygen Pitting
Corrosion of iron by dissolved oxygen is particularly bad since it is highly localized.
Oxygen pitting can occur under surface contaminants.
It is important to keep the metal surface free from build up of any substance.
Bacteria

28. Boiler tube oxygen pit

29. Alkaline Corrosion
In strongly basic solution (pH>11) the iron is oxidized by the hydrogen in the hydroxide ion causing caustic corrosion.
Fe + 2NaOH  Na2FeO2 + H2
The corrosion product (sodium ferrate) diffuses into underlying iron making it brittle.
This corrosion causes cracking of metals at pipe elbows where sodium hydroxide can accumulate.
To prevent this in boiler water make sure the alkalinity is not too high.

30. Corrosion Mechanisms
oxy redox summary of 3 types of corrosion

31. In all corrosion mechanisms:
The iron is oxidized into the Fe+2 ion which then dissolves into the electrolyte.
This results in the degradation of the metal.
The oxidized electrons in the metal
This constitutes a corrosion current in the metal.
The electrons enter the solution where they reduce an available element that is present.
This is hydrogen (in acids and caustic) or dissolved oxygen.
No corrosion occurs here at this site.
The corrosion product forms in the electrolyte
insoluble and forms sludge.

32. Corrosion Protection Methods
Keep the pH of the water between 9 and 11.
This reduces acid or alkaline corrosion to a minimum.
This is usually done by adding sodium hydroxide to neutral water.
Alkaline amines such as cyclohexylamine are added since they are volatile and carry over with the steam to prevent acid corrosion in the condensate return system.

33. Corrosion Protection Methods
a) Deaeration (removal of dissolved oxygen)
Remove dissolved oxygen from water.
This is done by heating the water to close to °C to drive off the dissolved gases in a deaerator.
Oxygen solubility decreases with temperature
This water must not be re-exposed to air or more oxygen will dissolve.

34. Corrosion Protection Methods
b) Oxygen Scavengers
If an oxygen scavenger chemical is dissolved in the water it will react with the dissolved oxygen to reduce its concentration.
Sodium sulfite protects boilers.
2Na2SO3 + O2  Na2SO4
Carbohydrazide is volatile so it carries over with steam to protect condensate.
(N2H3)2CO + 2O2  2N2 + 3H2O + CO2

35. Corrosion Protection Methods
Sacrificial Metals
If a metal higher than iron on the activity series such as zinc is bonded in good electrical contact to the iron, the zinc will corrode first, protecting the iron from corrosion.
This method of corrosion protection is called galvanic protection.
See the video referenced in the next slide

36. Video on Galvanic Cell Galvanic Cell Video
Both ZnSO4 and CuSO4 are 100% ionized
Zn, more anodic (reactive) than Cu, gives up its electrons to Cu
referring to the Activity Series for metals, Zn is more reactive than Cu
The zinc electrode corrodes over time

37. Corrosion Protection Methods
Protective Coatings
If the metal is coated so it is not exposed to water it will resist corrosion.
Aluminum oxide forms a durable oxide coat that protects aluminum from corrosion.
The chromium in stainless steel forms a durable oxide coat offering corrosion protection.
Magnetite (a form of iron oxide) protects iron.
Rust paints work by isolating the metal from water.

38. Corrosion Protect Methods
Impressed current
If the iron is connected to the negative terminal of a DC power supply, this will supply electrons to the iron preventing its loss of electrons (oxidation).
A current of about 50 to 100 mA is required for each square meter of iron to be protected.

39. Corrosion of Copper
Copper is low on the activity series meaning it is resistant to corrosion.
Over long periods of time the following reaction occurs:
4Cu + O2  2Cu2O
The green copper(I) oxide protects the underlying metal.
If carbonic acid is present this oxide coat reacts with it and dissolves and further corrosion can occur.
Cu2O (s) + 2H2CO3 (aq)  2Cu(HCO3)2 (aq)