Lesson 10: Corrosion and techniques to avoid corrosion
Corrosion: - undesired oxidation of metals. E.g., the corrosion of Iron is called Rust.
The term rust applies to Iron only, while corrosion is a general term that applies to all
other oxidations of metals.
Q: How does Iron rust?
A: Consider the following diagram.
O2(g)
Fe+2
H2O
Fe(s)
When the O2 comes in contact with the outer surface of the water drop, the following
reduction half reaction takes place:
1/2O2(g) + H2O(l) + 2e- ---------
2OH-(aq)
Therefore the outer surface of the water drop is acting as the cathode.
At the surface of the iron and the water drop, the following oxidation half reaction takes
place:
Fe(s) ----------- Fe +2(aq) + 2eTherefore the water closest to the Iron is acting as the anode.
The Fe+2 and the OH- combine the form Fe(OH)2. This substance can undergo further
REDOX to form Fe(OH)3. This molecule under dry conditions will reconfigure itself to a
hydrated version of Fe2O3(s). Each oxidation state of the Iron has a characteristic colour;
Fe+2 is green, while Fe+3 is the classic red/brown or “rust” colour.
Q: So how do I prevent rusting from occurring?
A: There are several answers:
1. Paint the metal or coat it in oil/grease. (something that prevents O2 and H2O from
touching the metal).
2. Keeping the metal cold will slow the reaction rate significantly.
3. Cathodic Protection
Q: What is cathodic protection???
A: Protecting a metal by contacting it with a more reactive one. E.g., metals are often
coated with Zinc because Zinc tends to be more reactive than the metals we are trying to
protect. Metals that are coated this way are said to be galvanized. Nails are often coated
in Zinc to prevent the iron underneath it from rusting.
The zinc that oxidizes forms a hard coating which prevents further corrosion from
occurring because O2 and H2O can no longer come in contact with the surface of the Zinc
or the Iron (unless you make a deep scratch in the nail).
e.g., cathodic protection can also be used to protect boat motors or metal pipes. See the
diagram below.
Magnesium band
Iron pipe
In this case, the Magnesium is a stronger reducing agent than the Iron pipe. When the
Magnesium reacts it oxidizes according to the following half equation:
Mg(s) ------------------ Mg+2(aq) + 2eThe electrons lost by the Mg are deposited on the Iron forcing the Iron to be a cathode
(b/c only reduction happens at the cathode, it is impossible for oxidation of the iron to
occur!)
Standard Reduction Potentials (Eo)
“Standard” conditions for electrochemical cells are listed in the data booklet on pg.8
- 25oC
- 1M solutions are used for each half-cell.
Eo = The voltage, or the work done per electron transferred (which is why we do not have
to multiply voltage when adjusting the number of e- in two half-equations).
The “o” means the voltage occurs in standard conditions.
The Zero voltage for Hydrogen is arbitrary, and the remaining voltages on the REDOX
table simply represent the difference in voltage between the two half-cells. (pg. 218
Hydrogen half-cell).
*Note: That’s why when we set up our own table using the sneaky trick, we can set the
bottom half-cell to zero (the assignment of Zero voltage is arbitrary. The difference in
voltage between the two half-cells will still be the same).
e.g., If the Lithum half-cell is set arbitrarily to Zero volts, what is the new voltage for the
Potassium half-cell? (Note the difference in voltage is the same!)
Things to watch out for!
1. Changing the surface area of an electrode does not change the voltage of a cell.
Since metals are pure, there is no increased tendency to react. (We are not talking
about the rate of the reaction).
2. Decreasing the concentration of a solution will decrease the voltage b/c there will
be a decreased tendency to react. (opposite is true for inc conc)
3. Operating cells are NOT in equilibrium ( a cell in equilibrium has a voltage of
Zero. i.e., it’s dead, it’s not operating!)