2.3 Oxidizing and Reducing Agents | Electrochemistry | Chemistry 12

Key Definitions

In every redox reaction, one species gains electrons and another loses electrons. The two participants are called the oxidizing agent and the reducing agent.

Term	Electron Transfer	Change in Oxidation State	What Happens to It
Oxidizing agent (oxidant)	Gains electrons	Oxidation state decreases	Gets reduced
Reducing agent (reductant)	Loses electrons	Oxidation state increases	Gets oxidized

Memory aid: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).

Identifying Oxidizing and Reducing Agents

To identify the oxidizing and reducing agents in a reaction, assign oxidation numbers to every element and track the changes.

Example: The reaction of $M n O_{4}^{-}$ with $F e^{2 +}$ in acidic solution:

$M n O_{4}^{-} + 8 H^{+} + 5 F e^{2 +} \to M n^{2 +} + 5 F e^{3 +} + 4 H_{2} O$

Mn: $+ 7 \to + 2$ (decrease of 5) → oxidizing agent (gets reduced)
Fe: $+ 2 \to + 3$ (increase of 1) → reducing agent (gets oxidized)

Role of Oxidizing Agents

An oxidizing agent accepts electrons from the reducing agent. In doing so:

Its own oxidation state falls (it is reduced).
It enables the oxidation of the other species.

Common oxidizing agents and their colour changes:

Oxidizing Agent	Reduced Form	Colour Change
$K M n O_{4}$ (acidified)	$M n^{2 +}$	Deep purple → colourless
$K_{2} C r_{2} O_{7}$ (acidified)	$C r^{3 +}$	Orange → green
$C l_{2}$	$C l^{-}$	Yellow-green → colourless

Role of Reducing Agents

A reducing agent donates electrons to the oxidizing agent. In doing so:

Its own oxidation state rises (it is oxidized).
It enables the reduction of the other species.

Common reducing agents: $F e^{2 +}$ , $I^{-}$ , $S O_{2}$ , $H_{2}$ , metals such as Zn and Mg.

Relative Strength of Oxidizing and Reducing Agents

The standard electrode potential ( $E^{\circ}$ ) is used to compare the strengths of oxidizing and reducing agents:

Stronger oxidizing agent → more positive $E^{\circ}$ (greater tendency to be reduced).
Stronger reducing agent → more negative $E^{\circ}$ (greater tendency to be oxidized).

Trend in Halide Ions as Reducing Agents

Among the halide ions $F^{-}, C l^{-}, B r^{-}, I^{-}$ :

$Reducing power: F^{-} < C l^{-} < B r^{-} < I^{-}$

As ionic radius increases down the group, the outermost electrons are held less tightly by the nucleus, making them easier to lose. Therefore, $I^{-}$ is the strongest reducing agent among the halides.

Conversely, the halogens themselves follow the opposite trend as oxidizing agents:

$Oxidizing power: F_{2} > C l_{2} > B r_{2} > I_{2}$

$F_{2}$ has the highest $E^{\circ}$ value ( $+ 2.87 V$ ) and is the strongest oxidizing agent.

Worked Example

Q: In the reaction $2 K M n O_{4} + 10 F e S O_{4} + 8 H_{2} S O_{4} \to 2 M n S O_{4} + 5 F e_{2} (S O_{4})_{3} + K_{2} S O_{4} + 8 H_{2} O$ , identify the oxidizing and reducing agents.

Solution:

Mn: $+ 7 \to + 2$ (reduced) → $K M n O_{4}$ is the oxidizing agent
Fe: $+ 2 \to + 3$ (oxidized) → $F e S O_{4}$ is the reducing agent

Key Definitions

In every redox reaction, one species gains electrons and another loses electrons. The two participants are called the oxidizing agent and the reducing agent.

Term	Electron Transfer	Change in Oxidation State	What Happens to It
Oxidizing agent (oxidant)	Gains electrons	Oxidation state decreases	Gets reduced
Reducing agent (reductant)	Loses electrons	Oxidation state increases	Gets oxidized

Memory aid: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).

Identifying Oxidizing and Reducing Agents

To identify the oxidizing and reducing agents in a reaction, assign oxidation numbers to every element and track the changes.

Example: The reaction of $M n O_{4}^{-}$ with $F e^{2 +}$ in acidic solution:

$M n O_{4}^{-} + 8 H^{+} + 5 F e^{2 +} \to M n^{2 +} + 5 F e^{3 +} + 4 H_{2} O$

Mn: $+ 7 \to + 2$ (decrease of 5) → oxidizing agent (gets reduced)
Fe: $+ 2 \to + 3$ (increase of 1) → reducing agent (gets oxidized)

Role of Oxidizing Agents

An oxidizing agent accepts electrons from the reducing agent. In doing so:

Its own oxidation state falls (it is reduced).
It enables the oxidation of the other species.

Common oxidizing agents and their colour changes:

Oxidizing Agent	Reduced Form	Colour Change
$K M n O_{4}$ (acidified)	$M n^{2 +}$	Deep purple → colourless
$K_{2} C r_{2} O_{7}$ (acidified)	$C r^{3 +}$	Orange → green
$C l_{2}$	$C l^{-}$	Yellow-green → colourless

Role of Reducing Agents

A reducing agent donates electrons to the oxidizing agent. In doing so:

Its own oxidation state rises (it is oxidized).
It enables the reduction of the other species.

Common reducing agents: $F e^{2 +}$ , $I^{-}$ , $S O_{2}$ , $H_{2}$ , metals such as Zn and Mg.

Relative Strength of Oxidizing and Reducing Agents

The standard electrode potential ( $E^{\circ}$ ) is used to compare the strengths of oxidizing and reducing agents:

Stronger oxidizing agent → more positive $E^{\circ}$ (greater tendency to be reduced).
Stronger reducing agent → more negative $E^{\circ}$ (greater tendency to be oxidized).

Trend in Halide Ions as Reducing Agents

Among the halide ions $F^{-}, C l^{-}, B r^{-}, I^{-}$ :

$Reducing power: F^{-} < C l^{-} < B r^{-} < I^{-}$

Conversely, the halogens themselves follow the opposite trend as oxidizing agents:

$Oxidizing power: F_{2} > C l_{2} > B r_{2} > I_{2}$

$F_{2}$ has the highest $E^{\circ}$ value ( $+ 2.87 V$ ) and is the strongest oxidizing agent.

Worked Example

Q: In the reaction $2 K M n O_{4} + 10 F e S O_{4} + 8 H_{2} S O_{4} \to 2 M n S O_{4} + 5 F e_{2} (S O_{4})_{3} + K_{2} S O_{4} + 8 H_{2} O$ , identify the oxidizing and reducing agents.

Solution:

Mn: $+ 7 \to + 2$ (reduced) → $K M n O_{4}$ is the oxidizing agent
Fe: $+ 2 \to + 3$ (oxidized) → $F e S O_{4}$ is the reducing agent