Example 4.6: 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]
\(\ce{ZnCO3(s)}\)\(\ce{->}\)\(\ce{ZnO(s)}\)\(\ce{ + }\)\(\ce{CO2(g)}\)\(\ce{ }\)
[b]
\(\ce{2Ga(l)}\)\(\ce{ + }\)\(\ce{3Br2(l)}\)\(\ce{->}\)\(\ce{2GaBr3(s)}\)\(\ce{ }\)
[c]
\(\ce{2H2O2(aq)}\)\(\ce{->}\)\(\ce{2H2O(l)}\)\(\ce{ + }\)\(\ce{O2(g)}\)\(\ce{ }\)
[d]
\(\ce{BaCl2(aq)}\)\(\ce{ + }\)\(\ce{K2SO4(aq)}\)\(\ce{->}\)\(\ce{BaSO4(s)}\)\(\ce{ + }\)\(\ce{2KCl(aq)}\)\(\ce{ }\)
[e]
\(\ce{C2H4(g)}\)\(\ce{ + }\)\(\ce{3O2(g)}\)\(\ce{->}\)\(\ce{2CO2(g)}\)\(\ce{ + }\)\(\ce{2H2O(l)}\)\(\ce{ }\)
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
\(\ce{Ga(l)}\) to
+3 in
\(\ce{GaBr3(s)}\). The reducing agent is
\(\ce{Ga(l)}\). Bromine is reduced, its oxidation number decreasing from
0 in
\(\ce{Br2(l)}\) to
-1 in
\(\ce{GaBr3(s)}\). The oxidizing agent is
\(\ce{Br2(l)}\).
(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
\(\ce{H2O2(aq)}\) to
0 in
\(\ce{O2(g)}\). Oxygen is also reduced, its oxidation number decreasing from
-1 in
\(\ce{H2O2(aq)}\) to
-2 in
\(\ce{H2O(l)}\). 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
\(\ce{C2H4(g)}\) to
+4 in
\(\ce{CO2(g)}\). The reducing agent (fuel) is
\(\ce{C2H4(g)}\). Oxygen is reduced, its oxidation number decreasing from
0 in
\(\ce{O2(g)}\) to
-2 in
\(\ce{H2O(l)}\). The oxidizing agent is
\(\ce{O2(g)}\).