Question

Identify the redox reactions out of the following reactions and identify the oxidizing and reducing agents in them.

A) $3HCl\left(aq\right)+HN{O}_3\left(aq\right)\to C{l}_2\left(g\right)+NOCl\left(g\right)+2H_{2}O\left(l\right)$

B) $HgCl2\left(aq\right)+2KI\left(aq\right)\to HgI2\left(s\right)+2KCl\left(aq\right)$

C) $F{e}_2{O}_3\left(s\right)+3CO\left(g\right)\Delta \to 2Fe\left(s\right)+3C{O}_2\left(g\right)$

D) $PCl3\left(l\right)+3H2O\left(l\right)\to 3HCl\left(aq\right)+H3PO3\left(aq\right)$

E) $4NH3+3O2\left(g\right)\to 2N2\left(g\right)+6H2O\left(g\right)$

Answer 1

A)
Oxidation states of elements
$3HCl+HN{O}_3\to C{l}_2+NOCl+2H_{2}O$

$H=+1Cl=-1$ $H=+1N=+5O=-2$ $Cl=0$ $N=+1O=-2Cl=+1$ $H=+1O=-2$

Identifying oxidizing and reducing agents


$H=+1Cl=-1$ $H=+1N=+5O=-2$ $Cl=0$​ ​​​​​​$N=+1O=-2Cl=+1$ $H=+1O=-2$

In this reaction, $Cl$ (in $HCl$) getting oxidized from $-1$ to $0$ (in $C{l}_2$) and to $+1$ (in $NOCl$), so $HCl$ acts as a reducing agent. Nitrogen getting reduced from $+5$ (in $HN{O}_3$) to $+1$ (in $NOCl$), so $HN{O}_3$ in this reactions acts as an oxidizing agent. Because $HCl$ and $HN{O}_3$ are getting oxidized and reduced simultaneously in this reaction, it can be said that this reaction is a redox reaction.

B)
Oxidation states of elements
$HgC{l}_2+2KI\to Hg{I}_2+2KCl$
$Hg=+2Cl=-1$ $K=+1I=-1$ $Hg=+2I=-1$ $K=+1Cl=-1$

Identifying whether reaction is redox or not
This reaction is an example of displacement reaction as none of the element is changing its oxidation state. So, none of the reactants is undergoing oxidation or reduction. Thus, this reaction is not an example of a redox reaction.

C)
Oxidation states of elements
$F{e}_2{O}_3+3CO\to 2Fe+3C{O}_2$
$Fe=+3O=-2$ $C=+2O=-2$ $Fe=0$ $C=+4O=-2$

Identifying oxidizing and reducing agents


$Fe=+3O=-2$ $C=+2O=-2$ $Fe=0$ $C=+4O=-2$

In this reaction, $CO$ is getting oxidized as carbon is changing its oxidation state from $+2$ (in $CO$) to $+4$ (in $C{O}_2$). $F{e}_2{O}_3$ is getting reduced as iron is changing its oxidation state from $+3$ (in $F{e}_2{O}_3$) to $+0$ (in $Fe$). So, $F{e}_2{O}_3$ in this reaction acts as an oxidizing agent and $CO$ acts as a reducing agent. Moreover, because $F{e}_2{O}_3$ and $CO$ are getting reduced and oxidized simultaneously in this reaction, it can be said that this reaction is a redox reaction.

D)
Oxidation states of elements
$PC{l}_3+3H_{2}O\to 3HCl+H_{3}P{O}_3$
$P=+3Cl=-1$ $H=+1O=-2$ $H=+1Cl=-1$ $P=+3H=+1O=-2$

Identifying whether reaction is redox or not
In this reaction, no change in oxidation state is taking place which signifies that reactant is not undergoing either reduction or oxidation. Hence, this is not a redox reaction.

E)
Oxidation states of elements
$4N{H}_3+3O_2\to 2N_2+6H_{2}O$
$N=-3H=+1$ $O=0$ $N=0$ $H=+1O=-2$

Identifying oxidizing and reducing agents


$N=-3H=+1$ $O=0$ $N=0$ $H=+1O=-2$

In this reaction, $N{H}_3$ is getting oxidized as nitrogen is changing its oxidation state from $-3$ (in $N{H}_3$) to $0$ (in $N_2$). $O_2$ is getting reduced as oxygen is changing its oxidation state from $0$ (in $O_2$) to $-2$ (in $H_{2}O$). So, $O_2$ in this reaction acts as an oxidizing agent and $N{H}_3$ acts as a reducing agent. Moreover, because $O_2$ and $N{H}_3$ are getting reduced and oxidized simultaneously in this reaction, it can be said that this reaction is a redox reaction.