Redox reactions are chemical reactions in which the reactants undergo a change in their oxidation states. The term ‘redox’ is a shorthand for reduction-oxidation. All redox reactions can be divided into two separate processes: oxidation and reduction.

The Oxidation-Reduction (redox) reactions always involve the simultaneous oxidation and reduction of substances. The substance undergoing oxidation is known as the reducing agent, while the substance undergoing reduction is known as the oxidizing agent.

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Table of Contents

Types of Redox Reactions

Examples of Redox Reactions

Oxidation and Reduction Reaction

Oxidizing and Reducing Agents

Identification of Oxidizing and Reducing Agents

Balancing Redox Reactions

Problems on Balancing Redox Reactions

Applications of Redox Reaction

Frequently Asked Questions

What are Redox Reactions?

Redox reactions, or oxidation-reduction reactions, are chemical reactions in which the oxidation states of atoms are changed. These reactions involve the transfer of electrons between reactants.

A redox reaction can be defined as a chemical reaction in which electrons are transferred between two reactants participating in it. This transfer of electrons can be identified by observing the changes in the oxidation states of the reacting species.

An illustration showing the electron transfer between two reactants in a redox reaction is provided below.

It can be seen from the illustration that the reactant A has lost an electron, making it oxidized. On the other hand, reactant B has gained an electron, thus making it reduced.

Oxidation is the loss of electrons and the corresponding increase in the oxidation state of a given reactant. Reduction is the gain of electrons and the corresponding decrease in the oxidation state of a reactant.

Any redox reaction can be broken down into two half-reactions, namely the oxidation half-reaction and the reduction half-reaction. An oxidizing agent is an electron-accepting species which tends to undergo a reduction in redox reactions. On the other hand, an electron-donating species which tends to hand over electrons can be referred to as a reducing agent. These species tend to undergo oxidation.

When writing these half-reactions separately, they must be balanced such that the total number of electrons are equal on both sides.

#Types of Redox Reactions

The different types of redox reactions are:

• Combustion reactions
• Displacement reactions
• Oxidation-reduction reactions
• Acid-base reactions
• Precipitation reactions

Decomposition Reaction

Combination Reaction

Displacement Reaction

Disproportionation Reactions

Decomposition Reaction

This kind of reaction involves the breakdown of a compound into different compounds. Examples of these types of reactions include:

2Na + H2 → 2NaH

2H2O → 2H2 + O2

2Na2CO3 → 4Na2O + 2CO2

AB → A + B

All the above reactions result in the breakdown of smaller chemical compounds.

But, there is a special case that confirms that not all decomposition reactions are redox reactions.

For Example: CaCO3 → CaO + CO2

Check Out: Types of Reactions

Combination Reaction

An example of a combination reaction is: Sodium + Chlorine → Sodium Chloride

2HCl + O2 → CO2 + H2O

2Fe + 3O2 → 4Fe2O3

Displacement Reaction

In this kind of reaction, an atom or an ion of one element is replaced by an atom or an ion of another element. It can be represented in the form of X + YZ → XZ + Y. Further, displacement reactions can be categorized into…

Metal Displacement Reaction

Non-metal Displacement Reaction

Metal Displacement

In metallurgical processes, a metal present in the compound can be displaced by another metal through a reaction. These types of reactions are highly useful in obtaining pure metals from their ores.

For example:

CuSO₄ + Zn → Cu + ZnSO₄

Non-Metal Displacement Reactivity

In this type of reaction, we can find a hydrogen displacement and sometimes rarely occurring reactions involving oxygen displacement.

Disproportionation Reactions

Disproportionation reactions are known as reactions wherein a single reactant is both oxidized and reduced.

P4 + 3NaOH + 3H2O → 3NaH2PO2 + 3H2O + PH3

Disproportionation Reaction Video Lesson

![Disproportionation Reaction]()

Examples of Redox Reactions

1. Combustion of Methane: CH4 + 2O2 → CO2 + 2H2O
2. Rusting of Iron: 4Fe + 3O2 → 2Fe2O3

A few examples of redox reactions, along with their corresponding oxidation and reduction half-reactions, are presented in this section.

Example 1: Reaction Between Hydrogen and Fluorine

The reaction between hydrogen and fluorine can be written as:

$$2H_2 + F_2 \rightarrow 2HF$$

In this reaction, hydrogen is oxidized and fluorine is reduced.

2HF → H2 + F2

2H+ + 2e- → H2

2F^- → F₂ + 2e^-

The hydrogen and fluorine ions go on to combine in order to form hydrogen fluoride.

Example 2: Interaction Between Zinc and Copper

This is an example of a Metal Displacement Reaction, where Zinc displaces the Cu²⁺ ion in the Copper Sulfate Solution to form Copper, as shown in the reaction below.

Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)

Zn2+ + 2e– → Zn

Cu²⁺ + 2e^- → Cu

Thus, zinc displaces copper from the copper sulfate solution in a redox reaction.

Example 3: Reaction between Iron and Hydrogen Peroxide

Hydrogen peroxide oxidizes Fe2+ to Fe3+ when an acid is present. The reaction is provided below.

2Fe²⁺ + H₂O₂ + 2H⁺ → 2Fe³⁺ + 2H₂O

Fe2+ → Fe3+ + e⁻

2 H2O2 → O2 + 4 OH- + 2e-

Thus, the hydroxide ion formed from the reduction of hydrogen peroxide combines with the proton donated by the acidic medium to form water.

Oxidation and Reduction Reactions

In order to gain an understanding of redox reactions, let us first discuss oxidation and reduction reactions separately.

What is an Oxidation Reaction?

Oxidation may be defined as the loss of electrons from a substance; the other definition of oxidation reactions states that the addition of oxygen or the more electronegative element or removal of hydrogen or the more electropositive element from a substance is called an oxidation reaction.

Examples of Oxidation Reactions

2S(s) + O2 (g) → SO2 (g)

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)

What is a Reduction Reaction?

A reduction reaction is a type of chemical reaction in which electrons are added to an atom or molecule, resulting in the formation of a new compound.

Reduction reactions are defined as the gain of electrons, just like oxidation reactions. Any substance that gains electrons during a chemical reaction is said to be reduced.

The reduction reaction can also be expressed as the addition of hydrogen or a more electropositive element or the removal of a more electronegative element or oxygen from a substance.

Examples of Reduction Reactions:

2CH2CH2 (g) + 2H2 (g) → CH3CH3 (g)

2FeCl3 (aq) + 2H2 (g) → 2FeCl2 (aq) + 4HCl (aq)

Now, if we closely examine the above reaction, we would find that all the reactions above contain both reduction and oxidation reactions.

The reaction in which FeCl3 is getting reduced is due to the removal of the electronegative element chlorine. Meanwhile, hydrogen is getting oxidized as a result of the addition of chlorine, an electronegative element, in the same reaction.

Oxidizing Agents and Reducing Agents

• The substance (atom, ion, and molecule) that loses electrons and is thereby oxidised to a high valency state is called Oxidising agent.

The substance that gains electrons and is thereby reduced to a lower valency state is called an oxidizing agent.

Important Oxidizing Agents

Molecules are composed of electronegative elements, such as O2, O3, and X2 (halogens).

Compounds containing an element in an oxidized state Eg: KMnO4, K2Cv2O7, HNO3, KClo3*

Oxides of Metals and Non-Metals:

• Examples: MgO, CuO, CrO3, P4O10

Fluorine is the strongest oxidizing agent.

Important Reducing Agents

All metals, such as Na, Zn, Fe, and Al

A few non-metals such as Carbon (C), Hydrogen (H), Sulphur (S), and Phosphorus (P)

Hydracids, such as HCl, HBr, HI, and H2S

Examples of compounds containing an element in the lower oxidation state include:

• FeCl2
• FeSo4
• SnCl2
• Hg2Cl2

Metallic hydrides such as NaH, LiH, CaH2, etc.

Organic compounds such as HCOOH

Lithium is the strongest reducing agent in solution, whereas Cesium is the strongest reducing agent in the absence of water. The substances which act as both oxidizing and reducing agents are: H2O2, SO2, H2SO3, HNO2, and NaNO2.

Reduction Potential of a Half-Reaction

Each half-reaction that makes up a redox reaction has a standard electrode potential. This potential is equal to the voltage produced by an electrochemical cell, in which the cathode reaction is the half-reaction being considered, while the anode is a standard hydrogen electrode.

The reduction potentials (denoted by $E_0_{red}$) of the half-reactions are determined by the voltage produced. The reduction potential of a half-reaction is positive for oxidizing agents that are stronger than H$^+$, and negative for the weaker ones.

Examples of the reduction potentials of some species are +2.866 V for F₂ and -0.763 V for Zn²⁺.

Identification of Oxidizing and Reducing Agents

If an element is in its highest possible oxidation state in a compound, it can act as an oxidizing agent, such as KMnO₄, K₂Cr₂O₇, HNO₃, H₂SO₄, and HClO₄.

If an element is in its possible lower oxidation state within a compound, it can act as a reducing agent. For example:

• H2S
• H2C2O4
• FeSO4
• SnCl2

If a highly electronegative element is in its highest oxidation state, the compound will act as an oxidising agent.

The compound acts as a reducing agent when a highly electronegative element is in its lowest oxidation state.

Identify the oxidizing agent and reducing agent in the reactions.

* $2Na_2S_2O_3 + I_2 \rightarrow Na_2S_4O_6 + 2NaI$

* \begin{array}{l}2FeCl_{3} + H_{2}S \rightarrow 2FeCl_{2} + S + 2HCl \end{array}

* $3Mg + 2N_2 \rightarrow Mg_3N_2$

* \begin{array}{l}AgCN + C{{N}^{-}} \rightarrow {{\left[ Ag{{\left( CN \right)}_{2}} \right]}^{-}}\end{array}

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Solution

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| Oxidizing Agent |