Rusting

Question 1

What are the rusting half equations?

Answer 1

Eq 1. O2 (g) + 2H2O (l) + 4e- => 4OH- (aq)
Eq 2. Fe (s) => Fe2+ (aq) + 2e-

Question 2

What happens at position 1 (on the rusting diagram)

Answer 2

The conc of oxygen dissolved in the water is high here so this is where half equation 1 occurs (site of reduction).
The electrons needed come from half equation two, through the metal.

Question 3

What happens at position 2 on the diagram?

Answer 3

The conc of oxygen is low, half eq. 1 cannot occur here. It is the site of half eq. 2, eg. The site of oxidation. This releases e- which travel through the iron to position 1. As iron is converted to Fe2+ a pit forms. Corrosion is greatest at the centre of the water droplet, where conc of oxygen is low.

Question 4

What is the effect of impurities?

Answer 4

Some ionic impurities increase the conductivity of water.

Question 5

Ways to protect steel from rusting

Answer 5

1. Provide a barrier, between the steel and the air/water. Eg use oil or grease, paint or a plastic film.
2. Sacrificial protection; metal with a coating of zinc. Zinc is more negative than iron, so is oxidised first and corrodes protecting the iron.
3. Stainless steel; contains chromium which also sacrificially protects the iron.
4. Impressed current; for iron to corrode it must give e-, so providing e- by connecting the metal to the negative DC power supply provides e- and prevents corrosion. It becomes a protected cathode. Used to protect bridges and underground wires.