Redox Reactions - Examples, Types, Applications, Balancing

Lavoisier developed the concept of oxidation and is generally called the father of modern Chemistry because he made various significant contributions to understanding chemical reactions. He discovered the role of combustion in the oxygen. The scientist Humphry Davy instrumentally understood the nature of redox reactions from 1778 to 1829. He found that various responses can be done by electricity for a better understanding of oxidation and reduction processes another scientist Michael Faraday worked on electrochemistry which includes the law of electrolysis which is very important for understanding the nature of redox reaction

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This Story also Contains

1. What Is Redox Reaction
2. Some Solved Examples
3. Summary

Michael Faraday developed the oxidation-reduction concept, which involves the transfer of electrons and electrochemical equivalence. The idea of the redox reaction was not Completely discovered at that particular time it evolved as time passed by different Scientists and chemists. The redox reaction discovery was established as a very profound theory as it has various implications in multiple fields, such as from chemistry to biology and environmental science. Redox reaction has multiple benefits in the chemical process as well as in the biological process in such a way that it involves the transfer of electrons between the two substances which results in the change in Oxidation states of the number. In biological systems, the redox reactions are the main center of cellular respiration and photosynthesis, both processes are essential or we can say the main source of energy production. Redox reactions are used in various other processes such as the extraction of metals, in corrosion, and in many manufacturing processes. It has a vital role in the decomposition of pollutants and the recycling of ecosystem components.

Redox reaction includes reducing and oxidizing agents which are the main components of this reaction
Oxidizing Agent- The substance that gains electrons and gets reduced. It causes another substance to be oxidized.
Reducing Agent- The substance that loses electrons and gets oxidized. It causes another substance to be reduced.

What Is Redox Reaction

A redox (reduction-oxidation) reaction involves the transfer of electrons between two substances. This type of reaction is fundamental in chemistry and is central to many biological and industrial processes.

- Oxidation: The process in which a substance loses electrons. This leads to an increase in the oxidation state of the substance.
- *Reduction: The process in which a substance gains electrons. This leads to a decrease in the oxidation state of the substance.

Redox Reaction:

Any reaction, in which the electrons are exchanged between atoms or ions, represents a simultaneous process of oxidation and reduction is called a Redox Reaction.

A⁰ + B⁰ → A⁺ + B^-

Here, A to A⁺, Loss of electron ( Oxidation )

And B to B^-, Gain of electron ( Reduction )

Some examples include:

$\mathrm{Zn}(\mathrm{s})+\mathrm{CuSO}_4(\mathrm{aq}) \rightarrow \mathrm{ZnSO}_4(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$ (Metal displacement)
$\mathrm{Na}(\mathrm{s})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g})($ Non - Metal displacement $)$

Some other examples of redox reactions in which oxidation and reduction are happening simultaneously.

$\mathrm{P}_4(\mathrm{~s})+3 \mathrm{OH}^{-}(\mathrm{aq})+3 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{PH}_3(\mathrm{~g})+3 \mathrm{H}_2 \mathrm{PO}_2^{-}$$\mathrm{S}_8(\mathrm{~s})+12 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow 4 \mathrm{~S}^{2-}(\mathrm{aq})+2 \mathrm{~S}_2 \mathrm{O}_3^{2-}(\mathrm{aq})+6 \mathrm{H}_2 \mathrm{O}(\mathrm{l})$

Types Of Redox Reaction

The different types of redox reactions are:

• Decomposition Reaction
• Combination Reaction
• Displacement Reaction
• Disproportionation Reactions

Decomposition Reaction

This is the reaction that involves the breakdown of a compound into different compounds. Some examples of this type of reaction are:

$2 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \xrightarrow{\Delta} 2 \mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g})$$2 \mathrm{KClO}_3(\mathrm{~s}) \xrightarrow{\Delta} 2 \mathrm{KCl}(\mathrm{s})+3 \mathrm{O}_2(\mathrm{~g})$

This must be noted here that all decomposition reactions are not redox reactions. For example, the decomposition of calcium carbonate is not a redox reaction.$\mathrm{CaCO}_3(\mathrm{~s}) \xrightarrow{\Delta} \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_2(\mathrm{~g})$

Combination Reaction

These types of reactions are the opposite of decomposition reactions and hence involve the combination of two species to form a single compound. Some examples include:
$\mathrm{C}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\Delta} \mathrm{CO}_2(\mathrm{~g})$$3 M g(s)+N_2(g) \rightarrow M g_3 N_2(s)$

Displacement Reaction

Displacement reactions, also known as replacement reactions, involve compounds and the replacement of elements. They occur as single and double replacement reactions. In other words, in these type of reactions, an atom or an ion in a compound is substituted by another element.

Some examples include:

$\mathrm{Zn}(\mathrm{s})+\mathrm{CuSO}_4(\mathrm{aq}) \rightarrow \mathrm{ZnSO}_4(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$ (Metal displacement)
$\mathrm{Na}(\mathrm{s})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g})(\mathrm{Non}-$ Metal displacement

Disproportionation Reactions

Disproportionation reactions are those reactions in which a single element in one oxidation state is simultaneously oxidized and reduced. Some examples include:

$\mathrm{P}_4(\mathrm{~s})+3 \mathrm{OH}^{-}(\mathrm{aq})+3 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{PH}_3(\mathrm{~g})+3 \mathrm{H}_2 \mathrm{PO}_2^{-}$$\mathrm{S}_8(\mathrm{~s})+12 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow 4 \mathrm{~S}^{2-}(\mathrm{aq})+2 \mathrm{~S}_2 \mathrm{O}_3^{2-}(\mathrm{aq})+6 \mathrm{H}_2 \mathrm{O}(\mathrm{l})$

Recommended topic video on(Redox Reactions)

Some Solved Examples

Example.1

1. Which of the following is a redox reaction?

1)$\mathrm{NaCl}+\mathrm{KNO}_3 \rightarrow \mathrm{NaNO}_3+\mathrm{KCl}$

2)$\mathrm{CaC}_2 \mathrm{O}_4+2 \mathrm{HCl} \rightarrow \mathrm{CaCl}_2+\mathrm{H}_2 \mathrm{C}_2 \mathrm{O}_4$

3)$\mathrm{Mg}(\mathrm{OH})_2+2 \mathrm{NH}_4 \mathrm{Cl} \rightarrow \mathrm{MgCl}_2+2 \mathrm{NH}_4 \mathrm{OH}$

4) (correct)$\mathrm{Zn}+2 \mathrm{AgCN} \rightarrow 2 \mathrm{Ag}+\mathrm{Zn}(\mathrm{CN})_2$

Solution

In a Redox Reaction, the oxidation states of the constituents should change.

Keeping that in mind, let us look at each reaction:

$\stackrel{+1-1}{\mathrm{NaCl}}+\stackrel{+1+5-2}{\mathrm{KNN}} \stackrel{+}{\mathrm{N}}_3 \rightarrow \stackrel{+1+5-2}{\mathrm{Na}} \mathrm{N}_3+\stackrel{+1-1}{\mathrm{KCl}}$

There is no change in the oxidation states.

$\mathrm{CaC}_2 \mathrm{O}_4+2 \mathrm{HCl} \rightarrow \mathrm{CaCl}_2+\mathrm{H}_2 \mathrm{C}_2 \mathrm{O}_4$

There is no change in the oxidation states.

$\mathrm{Mg}(\mathrm{OH})_2+2 \mathrm{NH}_4 \mathrm{Cl} \rightarrow \mathrm{MgCl}_2+2 \mathrm{~N} \mathrm{H}_4 \mathrm{OH}$

There is no change in the oxidation states.

Finally,

$\stackrel{0}{\mathrm{Zn}}+2 \stackrel{+1}{\mathrm{Ag}} \stackrel{+2-3}{\mathrm{~N}} \rightarrow 2 \stackrel{0}{\mathrm{Ag}}+\stackrel{+2}{\mathrm{Zn}}(\stackrel{+2-3}{\mathrm{C}})_2$

There is a change in the oxidation states of the constituents.

The last reaction is redox.

The oxidation states show a change only in reaction (d).

Hence, the answer is the option (4).

Example.2

2. Which of the following reactions is an example of a redox reaction?

1)XeF₆ + H₂O → XeOF₄ + 2HF

2)XeF₆ + 2H₂O → XeO₂F₂ + 4HF

3) (correct)XeF₄ + O₂F₂ → XeF₆ + O₂

4)XeF₂ + PF₅ → [XeF]⁺ PF₆⁻

Solution

In the given reactions:

$\mathrm{XeF}_6+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{XeOF}_4+2 \mathrm{HF}$

The oxidation states are:

$\mathrm{XeF}_6+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{XeOF}_4+2 \mathrm{HF}$

There is no change in the oxidation states of any of the reactants.

The above is not a redox reaction.

Similarly, $\mathrm{XeF}_6+2 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{XeO}_2 \mathrm{~F}_2+4 \mathrm{HF}$

$\therefore$ Above is not a redox reaction

In the reaction,

(+2) (-1) (+5) (-1) (+2) (-1) (+5) (-1)

$X e F_2+P F_5 \rightarrow[X e F]^{+}\left[P F_6\right]$

$\therefore$ Above is not a redox reaction

Now,

+4 -1 +1 -1 +6 -1 0

$X e F_4+O_2 F_2 \rightarrow X e F_6+O_1$

The oxidation state of Xe changed from +4 to +6 and that of oxygen changed from (+1) to 0

This is an example of a redox reaction.

Hence, the answer is an option (3).

Example.3

3. In the following reaction which species undergoes reduction?

$\left.2 \mathrm{Na}_{(} \mathrm{s}\right)+\mathrm{H}_{2(g)} \rightarrow 2 \mathrm{NaH}_{(\mathrm{s}}$

1)Na

2) (correct)H

3)NaH

4)Not a redox reaction.

Solution

Redox Reaction -

Redox reaction is a class of reactions in which oxidation and reduction reactions occur simultaneously.

Since hydrogen is more electronegative than sodium, so hydrogen undergoes reduction.

Hence, the answer is the option (2).

Example.4

4. In the following reaction, which species undergoes oxidation?

$\mathrm{Cl}_{2(g)}+3 \mathrm{~F}_{2(g)} \rightarrow 2 \mathrm{ClF}_{3(g)}$

1) (correct)Cl

2)F

3)Neither

4)The reaction is not redox

Solution

Redox Reaction -

Redox reaction is a class of reactions in which oxidation and reduction reactions occur simultaneously.

Since fluorine is more electronegative than chlorine, so chlorine undergoes oxidation.

Hence, the answer is the option (1).

Example.5

5. On heating $\mathrm{KClO}_2$ it gives KCl and O₂ This reaction is known as

1)Oxidation

2)Reduction

3)Disproportionation

4) (correct)Redox

Solution

Redox reaction is a class of reactions in which oxidation and reduction reactions occur simultaneously.

In the reaction :

$2 \mathrm{KClO}_3 \rightarrow 2 \mathrm{KCl}+3 \mathrm{O}_2$

Cl is reduced from +5 to -1 oxidation state and O is oxidized from -2 to 0 oxidation state.

So, it is an example of a redox reaction.

Hence, the answer is the option (4).

Summary

Redox reactions are essential in mining for extracting metals, processing ores, managing environmental impacts, and optimizing energy use. They are fundamental to both the technical and environmental aspects of mining operations. Redox reactions are often exothermic or endothermic, impacting the energy requirements of mining processes. Efficient use of energy in reduction and other processes can significantly affect the overall cost and environmental footprint of mining operations. Redox reaction has various impacts on the environment such as Redox reactions are involved in the formation of acid mine drainage, a byproduct of mining. Sulfide minerals in ore can oxidize to form sulfuric acid when exposed to air and water, leading to environmental contamination. Understanding these reactions helps in developing methods to mitigate their impact. Knowledge of redox reactions aids in the remediation of contaminated sites. For instance, chemical treatments can be used to reduce harmful metals or neutralize acidic conditions resulting from mining activities.

Redox reactions are also involved in metal extraction and mineral processing in such a way that Most metals are found in nature as oxides or sulfides. To obtain the pure metal, these ores must undergo reduction reactions. For example, in iron ore extraction, iron oxide is reduced to produce iron metal. This is achieved by reacting the ore with a reducing agent like carbon monoxide in a blast furnace. Redox reactions are also used in the leaching process, where oxidizing agents like oxygen or chlorine are used to dissolve valuable metals from ores. For instance, gold is often extracted from ore using cyanide leaching, where gold is oxidized to form a soluble complex. Redox reactions help in concentrating ores. For example, in the flotation process, reagents that promote or inhibit oxidation/reduction reactions are used to separate valuable minerals from gangue (waste material). pH Control as Managing the pH of the ore slurry is essential for efficient processing. Redox reactions help control the acidity or alkalinity of the solution, optimizing the extraction and concentration of desired metals.