7.1 Classical Idea of Redox Reactions-Oxidation and Reduction Reactions

Originally, the term oxidation was used to describe the addition of oxygen to an element or a compound. Because of the presence of dioxygen in the atmosphere \((\sim 20 \%)\), many elements combine with it and this is the principal reason why they commonly occur on the earth in the form of their oxides. The following reactions represent oxidation processes according to the limited definition of oxidation:

\(
\begin{aligned}
& 2 \mathrm{Mg}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{MgO}(\mathrm{s}) \dots(8.1) \\
& \mathrm{s}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{SO}_2(\mathrm{~g}) \dots(8.2)
\end{aligned}
\)
In reactions (8.1) and (8.2), the elements magnesium and sulphur are oxidised on account of addition of oxygen to them. Similarly, methane is oxidised owing to the addition of oxygen to it.
\(
\mathrm{CH}_4(\mathrm{~g})+2 \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \dots(8.3)
\)
A careful examination of reaction (8.3) in which hydrogen has been replaced by oxygen prompted chemists to reinterpret oxidation in terms of removal of hydrogen from it and, therefore, the scope of term oxidation was broadened to include the removal of hydrogen from a substance. The following illustration is another reaction where removal of hydrogen can also be cited as an oxidation reaction.
\(
2 \mathrm{H}_2 \mathrm{~S}(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{~S}(\mathrm{~s})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \dots(8.4)
\)
As knowledge of chemists grew, it was natural to extend the term oxidation for reactions similar to (8.1 to 8.4 ), which do not involve oxygen but other electronegative elements. The oxidation of magnesium with fluorine, chlorine and sulphur etc. occurs according to the following reactions :
\(
\mathrm{Mg}(\mathrm{s})+\mathrm{F}_2(\mathrm{~g}) \rightarrow \mathrm{MgF}_2(\mathrm{~s}) \dots(8.5)
\)
\(
\mathrm{Mg}(\mathrm{s})+\mathrm{Cl}_2(\mathrm{~g}) \rightarrow \mathrm{MgCl}_2 \text { (s) } \dots(8.6)
\)
\(
\mathrm{Mg}(\mathrm{s})+\mathrm{S}(\mathrm{s}) \rightarrow \operatorname{MgS}(\mathrm{s}) \dots(8.7)
\)
Incorporating the reactions ( 8.5 to 8.7 ) within the fold of oxidation reactions encouraged chemists to consider not only the removal of hydrogen as oxidation, but also the removal of electropositive elements as oxidation. Thus the reaction :
\(
2 \mathrm{~K}_4\left[\mathrm{Fe}(\mathrm{CN})_6\right](\mathrm{aq})+\mathrm{H}_2 \mathrm{O}_2(\mathrm{aq}) \rightarrow 2 \mathrm{~K}_3\left[\mathrm{Fe}(\mathrm{CN})_6\right](\mathrm{aq})+2 \mathrm{KOH}(\mathrm{aq})
\)
is interpreted as oxidation due to the removal of electropositive element potassium from potassium ferrocyanide before it changes to potassium ferricyanide.

Oxidation

The term “oxidation” is defined as the addition of oxygen/electronegative element to a substance or removal of hydrogen/ electropositive element from a substance.

Oxidation is a process which involve (the following are the examples of oxidation processes)

• Addition of \(\mathrm{O} /\) electronegative element
  \(
  \begin{aligned}
  & \text { Ex. } 2 \mathrm{H}_2 \mathrm{~S}(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \longrightarrow 2 \mathrm{~S}(\mathrm{~s})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \\
  & \mathrm{Mg}(\mathrm{s})+\mathrm{F}_2(\mathrm{~g}) \longrightarrow \mathrm{MgF}_2(\mathrm{~s})
  \end{aligned}
  \)
• Removal of \(\mathrm{H} /\) electropositive element
  \(
  \begin{aligned}
  & \mathrm{Ex} \cdot 2 \mathrm{H}_2 \mathrm{~S}(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \longrightarrow 2 \mathrm{~S}(\mathrm{~s})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \\
  & 2 \mathrm{NaI}+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow 2 \mathrm{NaOH}+\mathrm{I}_2
  \end{aligned}
  \)
• Involves loss of electrons by a species
  \(
  \text { Ex. } \mathrm{Fe}^{2+} \longrightarrow \mathrm{Fe}^{3+}+\mathrm{e}^{-}
  \)

Reduction

In the beginning, reduction was considered as removal of oxygen from a compound. However, the term reduction has been broadened these days to include removal of oxygen/electronegative element from a substance or addition of hydrogen/ electropositive element to a substance.

Reduction is a process which involves (the following are the examples of reduction processes)

• Removal of oxygen from mercuric oxide
  \(
  2 \mathrm{HgO}(\mathrm{s}) \stackrel{\Delta}{\longrightarrow} 2 \mathrm{Hg}(\mathrm{l})+\mathrm{O}_2(\mathrm{~g}) \dots(8.8)
  \)
• Removal of electronegative element, chlorine from ferric chloride
  \(
  2 \mathrm{FeCl}_3(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g}) \rightarrow 2 \mathrm{FeCl}_2(\mathrm{aq})+2 \mathrm{HCl}(\mathrm{aq}) \dots(8.9)
  \)
• Addition of hydrogen
  \(
  \mathrm{CH}_2=\mathrm{CH}_2(\mathrm{~g})+\mathrm{H}_2(\mathrm{~g}) \rightarrow \mathrm{H}_3 \mathrm{C}-\mathrm{CH}_3(\mathrm{~g}) \dots(8.10)
  \)
• Addition of mercury to mercuric chloride
  \(
  2 \mathrm{HgCl}_2(\mathrm{aq})+\mathrm{SnCl}_2(\mathrm{aq}) \rightarrow \mathrm{Hg}_2 \mathrm{Cl}_2(\mathrm{~s})+\mathrm{SnCl}_4(\mathrm{aq}) \dots(8.11)
  \)
• Involves gain of electrons by a species
  \(
  \mathrm{Sn}^{4+}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Sn}^{2+}
  \)

In reaction (8.11) simultaneous oxidation of stannous chloride to stannic chloride is also occurring because of the addition of electronegative element chlorine to it. It was soon realised that oxidation and reduction always occur simultaneously (as will be apparent by re-examining all the equations given above), hence, the word “redox” was coined for this class of chemical reactions.

Example 8.1: In the reactions given below, identify the species undergoing oxidation and reduction:
(i) \(\mathrm{H}_2 \mathrm{~S}(\mathrm{~g})+\mathrm{Cl}_2\) (g) \(\rightarrow 2 \mathrm{HCl}\) (g) \(+\mathrm{S}\) (s)
(ii) \(3 \mathrm{Fe}_3 \mathrm{O}_4\) (s) \(+8 \mathrm{Al}\) (s) \(\rightarrow 9 \mathrm{Fe} (s)+4 \mathrm{Al}_2 \mathrm{O}_3(\mathrm{~s})\)
(iii) \(2 \mathrm{Na}\) (s) \(+\mathrm{H}_2\) (g) \(\rightarrow 2 \mathrm{NaH}\) (s)

Answer:(i) \(\mathrm{H}_2 \mathrm{S}\) is oxidised because a more electronegative element, chlorine is added to hydrogen (or a more electropositive element, hydrogen has been removed from \(\mathrm{S}\) ). Chlorine is reduced due to addition of hydrogen to it.
(ii) Aluminium is oxidised because oxygen is added to it. Ferrous ferric oxide \(\left(\mathrm{Fe}_3 \mathrm{O}_4\right)\) is reduced because oxygen has been removed from it.
(iii) With the careful application of the concept of electronegativity only we may infer that sodium is oxidised and hydrogen is reduced.

Reaction (iii) chosen here prompts us to think in terms of another way to define redox reactions.