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Describing Redox Reactions

Example on Describing Redox Reactions

Identify which equations represent redox reactions, providing a name for the reaction if appropriate. For those reactions identified as redox, name the oxidant and reductant.

(a) $$\mathrm{ZnCO_3}(s) \longrightarrow \mathrm{ZnO}(s) + \mathrm{CO_2}(g)$$

(b) $$\mathrm{2Ga}(l) + \mathrm{3Br_2}(l) \longrightarrow \mathrm{2GaBr_3}(s)$$

(c) $$\mathrm{2H_2O_2}(aq) \longrightarrow \mathrm{2H_2O}(l) + \mathrm{O_2}(g)$$

(d) $$\mathrm{BaCl_2}(aq) + \mathrm{K_2SO_4}(aq) \longrightarrow \mathrm{BaSO_4}(s) + \mathrm{2KCl}(aq)$$

(e) $$\mathrm{C_2H_4}(g) + \mathrm{3O_2}(g) \longrightarrow \mathrm{2CO_2}(g) + \mathrm{2H_2O}(l)$$

Solution

Redox reactions are identified per definition if one or more elements undergo a change in oxidation number.

(a) This is not a redox reaction, since oxidation numbers remain unchanged for all elements.

(b) This is a redox reaction. Gallium is oxidized, its oxidation number increasing from 0 in Ga(l) to +3 in GaBr3(s). The reducing agent is Ga(l). Bromine is reduced, its oxidation number decreasing from 0 in Br2(l) to −1 in GaBr3(s). The oxidizing agent is Br2(l).

(c) This is a redox reaction. It is a particularly interesting process, as it involves the same element, oxygen, undergoing both oxidation and reduction (a so-called disproportionation reaction). Oxygen is oxidized, its oxidation number increasing from −1 in H2O2(aq) to 0 in O2(g). Oxygen is also reduced, its oxidation number decreasing from −1 in H2O2(aq) to −2 in H2O(l). For disproportionation reactions, the same substance functions as an oxidant and a reductant.

(d) This is not a redox reaction, since oxidation numbers remain unchanged for all elements.

(e) This is a redox reaction (combustion). Carbon is oxidized, its oxidation number increasing from −2 in C2H4(g) to +4 in CO2(g). The reducing agent (fuel) is C2H4(g). Oxygen is reduced, its oxidation number decreasing from 0 in O2(g) to −2 in H2O(l). The oxidizing agent is O2(g).