1. CHAPTER 2:
BASICS OF CORROSION

2. The Basic Principles of Chemistry
Chemistry is a branch of physical science that studies the composition, structure, properties and changes of matter.
Chemistry is chiefly concerned with atoms and molecules and their interactions and transformations.

3. Matter, Elements, Atoms, and Ions
Matter is anything that has mass and occupies space.
All matter is made up of substances called elements, which have specific chemical and physical properties and cannot be broken down into other substances through ordinary chemical reactions.
The structure of the atom
An atom is the smallest unit of matter that retains all of the chemical properties of an element.

4. Electron, Proton, and Neutron
Electrons are present outside the nucleus of an atom and carry negative charges. They also hold atoms together.
The mass of an electron is considered to negligible. They are spread around the nucleus
Proton
Protons are present in the nucleus of an atom and carry positive charges. They are closely bound.
The mass of a proton is approximately 1840 times as the mass of an electron.
Neutron
Neutrons are present in the nucleus of an atom and they are neutral. They are closely bound.
The mass of a neutron is nearly equal to the mass of a proton

5. The Periodic Law
The periodic law states that the properties of elements recur in a repeating pattern when arranged according to increasing atomic number.

6. Hydrogen on the Periodic Table
Hydrogen occupies a special position on the periodic table.
It is a gas with properties similar to nonmetals.
It also reacts by losing one electron, similar to metals.
We will place hydrogen in the middle of the periodic table to recognize its unique behavior.

7. Valence Electrons
When an atom undergoes a chemical reaction, only the outermost electrons are involved.
These electrons are of the highest energy and are furthest away from the nucleus. These are the valence electrons.

8. Ionic Charge
Recall, that metals tend to lose electrons and nonmetals tend to gain electrons.
The charge on an ion is related to the number of valence electrons on the atom.
Group IA/1 metals lose their one valence electron to form 1+ ions.
Na → Na+ + e-
Metals lose their valence electrons to form ions.

9. Ions, Cations, and Anions

10. Chemical Compound Chemical Compound:
Any substance consisting of two or more different types of elements in a fixed proportion of its atoms

11. Ionic Charges

12. The Nature of Corrosion Reactions
The special characteristic of most corrosion processes is that the oxidation and reduction steps occur at separate locations on the metal.
This is possible because metals are conductive, so the electrons can flow through the metal from the anodic to the cathodic regions.
The presence of water is necessary in order to transport ions to and from the metal, but a thin film of adsorbed moisture can be sufficient.

13. The Nature of Corrosion Reactions

14. Galvanic Series
The galvanic series determines the nobility of metals and semi-metals.
When two metals are submerged in an electrolyte, while also electrically connected by some external conductor, the less noble (base) will experience galvanic corrosion.

15. Acidity and Alkalinity (pH)
The atoms of a molecule are held together by a force referred to as chemical bonding. It is this chemical bonding that defines many of the properties of a substance.
When discussing an aqueous medium (including soil), the question often arises as to how acid or alkaline the solution is.
Referring to whether there is an excess of hydrogen (H+) or hydroxyl (OH-) ions present.
When acids dissociate, the cation produced is the hydrogen ion, H+.

16. Acidity and Alkalinity (pH)
A medium is said to be acidic when there is an excess of H+ ions.
The strength of an acid is a measure of the hydrogen ion concentration in an aqueous solution and is classified according to the pH scale.
The definition of pH is the negative logarithm to the base 10 of the hydrogen ion concentration, or: pH = -log [H+]
The neutral pH point is 7.
Acid solutions have a pH below 7 and alkaline, or basic, solutions have a pH above 7.

17. Acidity and Alkalinity (pH)
Since the pH scale is logarithmic, for each unit of pH the environment becomes ten times more acid or alkaline.
This concept is better understood if we look at pure water, H2O.
Pure water will ionize into equal parts of hydrogen ions (H+) and hydroxyl ions (OH-).
An understanding of pH is important in corrosion and cathodic protection work.
For many metals, the rate of corrosion increases appreciably below a pH of about 4.
Between 4 and 8 corrosion rate is fairly independent of pH.
Above 8, the environment becomes passive and corrosion rates tend to decrease.

18. Acidity and Alkalinity (pH)
The corrosion rate of aluminum and lead, on the other hand, tends to increase in environments above a pH of about 8.
This is because the protective oxide film on the surface of these metals is dissolved in most strong acids and alkalis and the metals corrode.
Metals that corrode under low and high pH levels are termed amphoteric metals.
An understanding of the effect of pH is also important in the application of cathodic protection.
The pH of the environment around the cathode (the protected structure) becomes more alkaline due to the production of hydroxyl ions or removal of hydrogen ions.

19. Electrochemistry
One branch of electrochemistry deals with solid state reactions that take place in semiconductors such as transistors and diodes.
Another branch deals with chemical changes that accompany the passage of an electric current.
Or a process in which a chemical reaction produces an electric current.
Corrosion and cathodic protection pertain to the branch of electrochemistry concerned with chemical changes that accompany the transfer of electric charge, or reactions that produces electric current in aqueous or other liquid environments.

20. Electrochemical Reactions
Metallic corrosion is almost always an electrochemical process.
An electrochemical reaction is a chemical reaction:
involving the transfer of electrons.
involves oxidation and reduction.
Battery or an electrochemical cell is generating electrical energy from chemical reactions.
The principle of a Daniell cell shown here.
Zinc loses electrons and copper gains electrons.
Electrons is transferred through an electrically path.
Zinc anode: Zn(s) → Zn2+ + 2e-
Copper cathode: Cu e- → Cu(s)

21. Electrochemical Reactions
A piece of zinc immersed in hydrochloric acid solution is undergoing corrosion.
Zinc is transformed to zinc ions.
This reaction produces electrons and these pass through the metal to other site where hydrogen ions are reduced to hydrogen gas.

22. Anodic Processes
The individual anodic reactions for iron, nickel, and aluminum are listed as follows:
The corrosion of metal M results in the oxidation of metal M to an ion with a valence charge of n+ and the release of n electrons.

23. Faraday’s Law
By using Michael Faraday relation we convert the current generated by one of the anodic reactions to an equivalent mass loss or corrosion penetration rate.
As shown before:
According to Faraday’s law, the reaction with 1 mol of Metal would require n moles of electrons. (1 mol =Avogadro’s number (6.022 × 1023)).
Charge carried by 1 electron = 1.6 × 10−19 coulomb (C).
The charge carried by 1 mol of electrons = 1 faraday (F).
1 F = 96,485 C/(mol of electrons).
Q=n.F.m/M
I=n.F.K/M

24. Conversion between Current, Mass Loss, and Penetration Rates
The charge of any electrochemical reaction is: Q=nFm/M
Q can be defined in terms of electric current: I=nFK/M
Q=n.F.m/M
I=n.F.K/M
Conversion between Current, Mass Loss, and Penetration Rates for Steel.
The corrosion current itself can be either estimated by using specialized electrochemical methods or by using weight-loss data.

25. Cathodic Processes
The reduction of hydrogen ions at a cathodic surface will make the solution less acidic or more alkaline or basic at the corroding interface.

26. Cathodic Processes
In neutral waters the anodic corrosion of some active metals will split water directly as shown:
Cathodic reactions encountered during the corrosion of metals are:

27. Cathodic Processes
Oxygen reduction is a very common cathodic reaction.
Although less common, metal ion reduction and metal deposition, can cause severe corrosion problems in special situations.

28. Cathodic Processes
All corrosion reactions are simply combinations of one or more of the above cathodic reactions.

29. Cathodic Processes
Consider the corrosion of zinc in a hydrochloric acid solution containing dissolved oxygen.

30. Surface Area Effect
Mild steel is placed in a solution of hydrochloric acid.

31. Surface Area Effect
The fact that there is no net accumulation of charges on a corroding surface is quite important for understanding most corrosion processes.
The absolute equality between the anodic and cathodic currents expressed below does not mean that the current densities for these currents are equal.
Relative anodic (Sa) and cathodic (Sc) surface areas and their associated current densities ia and ic expressed in units of mA cm−2.

32. Surface Area Effect
The larger the cathode compared with the anode, the more oxygen reduction, or other cathodic reaction, can occur and, hence, the greater the galvanic current.