Activity 1M (Textbook) – Chemistry Form 5 Chapter 1 (Redox Equilibrium)

Activity 1M:
Conduct this activity in groups.
1. Discuss and explain the following situation:
(a) What is the difference between rusting of iron and metal corrosion?

(b) Describe the mechanisms for rusting from the aspects of oxidation and reduction.

(c) How can metals that are more electropositive than iron prevent rusting of iron? State examples of application in daily life for each situation.

2. Present your discussion with the group members in a Gallery Walk activity.


Answer:
1.
(a) Rusting of iron is a metal corrosion that only occurs to iron, while metal corrosion occurs to all metals.

(b)
– Iron undergo oxidation reaction when iron atoms lose electrons to form Fe²⁺.
$$\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+2 \mathrm{e}^{-}$$

– Oxygen undergo reduction reaction when oxygen gains electrons to form OH⁻ ions.
$$\mathrm{O}_2+2 \mathrm{H}_2 \mathrm{O}+4 \mathrm{e}^{-} \rightarrow 4 \mathrm{OH}^{-}$$

Fe²⁺ ions react with OH⁻ ions to form iron(II) hydroxide, Fe(OH)₂.
$$\mathrm{Fe}^{2+}+2 \mathrm{OH}^{-} \rightarrow \mathrm{Fe}(\mathrm{OH})_2$$

– Fe(OH)₂ undergoes continuous oxidation with oxygen to form hydrated iron(III) oxide, Fe₂O₃. xH₂O, rust.


(c)
– A metal that is more electropositive than iron has a higher tendency to lose electrons. 

– A metal that is more electropositive than iron corrodes and rusting of iron is prevented.

– A metal that is more electropositive than iron becomes a sacrificial metal.

– Examples: Bridge pillars, ship hulls and underground pipes.