Self Assess 1.2 (Textbook) – Chemistry Form 5 Chapter 1 (Redox Equilibrium)

Question 1:
Table 1.7 shows the standard electrode potential values of half-cells for some metals.


(a) Arrange the atoms or ions in Table 1.7 in an ascending order of the strength of oxidising agents and reducing agents.

(b) Based on your answer in (a), explain if the reaction will occur for the following reactants:
(i) Mg(s) + Cu²⁺(aq).
(ii) Mg(s) + Zn²⁺(aq).
(iii) Cu(s) + Zn²⁺(aq).


Answer:
(a)
Arrange the standard electrode potential value, E⁰ from the most negative to the most positive.
$$\begin{array}{ll} \mathrm{Mg}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{Mg}(\mathrm{~s}) & \mathrm{E}^0=-2.38 \\ \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{Zn}(\mathrm{~s}) & \mathrm{E}^0=-0.76 \end{array}$$
$$\mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{Cu}(\mathrm{~s}) \quad \mathrm{E}^0=+0.34$$
$$\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{e}^{-} \rightleftharpoons \mathrm{Ag}(\mathrm{~s}) \quad \mathrm{E}^0=+0.80$$
$$\text { Oxidising agent : } \mathrm{Mg}^{2+}, \mathrm{Zn}^{2+}, \mathrm{Cu}^{2+}, \mathrm{Ag}^{+}$$
$$\text { Reducing agent : } \mathrm{Ag}, \mathrm{Cu}, \mathrm{Zn}, \mathrm{Mg}$$
(b)(i)
Reaction occur
– Mg is a stronger reducing agent. Mg easily releases electrons and oxidation reaction occurs.

– Magnesium can displace copper from its salt solution because magnesium is a stronger reducing agent compared to copper.


(ii)
Reaction occur
– Mg is a stronger reducing agent. Mg easily releases electrons and oxidation reaction occurs.

– Magnesium can displace zinc from its salt solution because magnesium is a stronger reducing agent compared to zinc.

(iii)
Reaction does not occur
– Cu is a weaker reducing agent. Cu atom is difficult to lose electrons and oxidation reaction does not occur.

– Copper cannot displace zinc from its salt solution because copper is a weaker reducing agent compared to zinc.