Types of Redox Reactions: Redox reactions constantly happen in our surroundings. Humans probably used oxidation-reduction reactions for the first time about 7,500-4,500 years ago during the Copper/Bronze Age, when copper ores were heated in the presence of carbon to produce copper metal. This process involved the reduction of copper to copper metal and the oxidation of carbon to carbon dioxide. A similar process was used to treat iron ores during the Iron Age, which occurred 4,500-3,500 years ago.

At present, Redox reactions govern most of our technology. Oxygen and carbon are redox reactants, and the chemical reactions they undergo can be simple, such as carbon burning with oxygen to produce carbon dioxide, or more complex, such as the human body oxidizing glucose through a variety of electron transfer reactions. To learn more about redox reactions and their types, read the below article.

Redox Reactions

Oxidation can be characterised as a process in which oxygen or another electronegative element is added or a process in which hydrogen or another electropositive element is removed. Reduction is a chemical reaction in which hydrogen or another electropositive element is added, or oxygen or another electronegative element is removed. A redox reaction is defined as a reaction in which both reduction and oxidation occurs together. Redox reactions are classified into four types. They are as follows;

Types of Redox Reactions

1. Combination Reactions

A combination reaction is a reaction in which two atoms or molecules combine to form a single molecule.

For example, \({\text{A}} + {\text{B}} + {\text{C}}\)

Either any or both of \({\text{A}}\) and \({\text{B}}\) must be in the elemental form for the combination reaction to be a redox reaction. Redox reactions include all combustion reactions that utilise elemental oxygen and all other reactions that use elements other than oxygen.

Some important combination reduction reactions are as follows;

\({\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right) + {\overset{0}{\mathop {\text{O}}} _2}\left({\text{g}} \right) \to {\overset{{ + 1}}{\mathop {\text{H}}} _2}{\overset{{ – 1}}{\mathop {\text{O}}} _2}\left({\text{l}} \right)\)

\(3\overset{0}{\mathop {\text{M}}} {\text{g}}\left({\text{s}} \right) + {\overset{0}{\mathop {~{\text{N}}}} _2}\left({\text{g}} \right) \to \overset{{{\text{ + 2}}}}{\mathop {{\text{M}}}} {{\text{g}}_3}{\overset{{ – 3}}{\mathop {~{\text{N}}}} _2}\left({\text{s}} \right)\)

\(\overset{0}{\mathop {\text{C}}} \left({\text{s}} \right) + {\overset{0}{\mathop {\text{O}}} _2}\left({\text{g}} \right) \to \overset{{ + 4}}{\mathop {\text{C}}} {\overset{{ – 2}}{\mathop {\,{\text{C}}}} _2}\left({\text{g}} \right)\)

\(\mathop {{\rm{Fe}}}\limits^0 \left( {\rm{s}} \right) + \mathop {\;{\rm{S}}}\limits^0 \left( {\rm{s}} \right) \to \mathop {\;{\rm{Fe}}}\limits^{ + 2} \mathop {\;{\mkern 1mu} {\rm{S}}}\limits^{ – 2} \left( {\rm{s}} \right)\)

2. Decomposition Reaction

In a decomposition reaction, a molecule breaks down into two or more components. All decomposition reactions are reverse of combination reactions. Furthermore, not all decomposition reactions are redox. One of the decomposition products must be in the elemental state for a decomposition reaction to be a redox reaction.

For example,

1. \(2{\overset{{ + 1}}{\mathop {\text{H}}} _2}\overset{{ – 2}}{\mathop {\text{O}}} \left( {\text{l}} \right)\;\underrightarrow \Delta\; 2{\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right) + {\overset{0}{\mathop {\text{O}}} _2}\left({\text{g}} \right)\)
2. \(2\overset{{ + 1}}{\mathop {{\text{Na}}}} \overset{{ – 1}}{\mathop {\text{H}}} \left({\text{s}} \right)\;\underrightarrow \Delta\; 2\overset{0}{\mathop {\text{N}}} {\text{a}}\left({\text{s}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)
3. \(2~\overset{{ + 1}}{\mathop {\text{K}}} \overset{{ + 5}}{\mathop {{\text{Cl}}}} \overset{{ – 2}}{\mathop {{{\text{O}}_3}}}\; \underrightarrow \Delta\; 2\overset{{ + 1}}{\mathop {\text{K}}} \overset{{ – 1}}{\mathop {{\text{Cl}}}} \left({\text{s}} \right) + 3{\overset{0}{\mathop {\text{O}}} _2}\left({\text{g}} \right)\)

In the above reactions, the product formed, such as \({{\text{H}}_2}\) and \({{\text{O}}_2}\) in reaction (i), \({\text{Na}}\) and \({{\text{H}}_2}\) in reaction (ii) and \({{\text{O}}_2}\) in reaction (iii) are in the elemental form. Hence, all these reactions are decomposition redox reactions.

However, calcium carbonate undergoes a decomposition reaction and forms calcium oxide and carbon dioxide. Hence, this reaction is not a redox reaction because both the product forms are not in elemental form.

\({\text{CaC}}{{\text{O}}_3}\left({\text{s}} \right)\;\underrightarrow \Delta\; {\text{CaO}}\left({\text{s}} \right) + {\text{C}}{{\text{O}}_2}\left({\text{g}} \right)\)

3. Displacement Reaction

A reaction in which an atom or ion in a compound is replaced by an atom or ion of a different element is a displacement reaction. In general, it is represented by the equation,

\({\text{X}} + {\text{YZ}} \to {\text{XZ}} + {\text{Y}}\)

Types of Displacement Reactions

The types of displacement reactions are explained below:

a) Metal Displacement Reactions

In this reaction, a metal in the compound is displaced by some other metal in the elemental state.

For example,

\(\underset{{{\text{ Copper}}\,{\text{sulphate }}}}{\mathop {\overset{{ + 2}}{\mathop {{\text{Cu}}}} \overset{{ + 6}}{\mathop {\text{S}}} \overset{{ – 2}}{\mathop {{{\text{O}}_4}}} \left({{\text{aq}}} \right)}} + \underset{{{\text{ Zinc }}}}{\mathop {\overset{0}{\mathop {{\text{Zn}}}} \left({\text{s}} \right)}} \to \underset{{{\text{ Copper }}}}{\mathop {\overset{0}{\mathop {{\text{Cu}}}} \left({\text{s}} \right)}} + \underset{{{\text{ Zinc}}\,{\text{sulphate }}}}{\mathop {\overset{{ + 2}}{\mathop {{\text{Zn}}}} \overset{{ + 6}}{\mathop {\text{S}}} \overset{{ – 2}}{\mathop {{{\text{O}}_4}}} }} \)

In the above reaction, copper is less reactive than zinc is displaced by zinc and forms zinc sulphate.

Similarly,

b) Non-metal Displacement Reactions

In these reactions, a metal or a non-metal displaces another non-metal from its compound. The non-metal displaced in most of these reactions is hydrogen. Some reactions involving the displacement of oxygen or halogens, on the other hand, are well-known.

Depending upon the capability of the reducing metal or non-metal, the following cases arise;

I. Dihydrogen is displaced from cold water by all alkali metals and some alkaline earth metals (\({\text{Ca}},{\text{Sr}}\) and \({\text{Ba}}\)), which are excellent reducing agents. For example,

\(2\overset{{\text{0}}}{\mathop {{\text{Na}}}} \left({\text{s}} \right) + 2{\overset{{ + 1}}{\mathop {\text{H}}} _2}\overset{{{\text{-2}}}}{\mathop {\text{O}}} \left({\text{l}} \right) \to 2\overset{{ + 1}}{\mathop {{\text{Na}}}} \overset{{{\text{-2}}}}{\mathop {\text{O}}} \overset{{ + 1}}{\mathop {\text{H}}} \left({{\text{aq}}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)

\(\overset{{\text{0}}}{\mathop {{\text{Ca}}}} \left({\text{s}} \right) + 2{\overset{{ + 1}}{\mathop {\text{H}}} _2}\overset{{{\text{-2}}}}{\mathop {\text{O}}} \left({\text{l}} \right) \to \overset{{{\text{2}} + }}{\mathop {{\text{Ca}}}} {\left({\overset{{{\text{-2}}}}{\mathop {\text{O}}} \overset{{+ 1}}{\mathop {\text{H}}} } \right)_2}\left({{\text{aq}}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)

II. Less reactive metals such as magnesium and iron react with steam to produce dihydrogen gas. For example,

\(\overset{0}{\mathop {{\text{Mg}}}} \left({\text{s}} \right) + 2{\overset{{ + 1}}{\mathop {\text{H}}} _2}\overset{{ – 2}}{\mathop {\text{O}}} \left({\text{g}} \right) \to \overset{{ + 2}}{\mathop {{\text{Mg}}}} {\left({\overset{{ – 2}}{\mathop {\text{O}}} \overset{{ + 1}}{\mathop {\text{H}}} } \right)_2}\left({{\text{aq}}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)

\(3{\text{Fe}}\left({\text{s}} \right) + 4{\overset{{ + 1}}{\mathop {\text{H}}} _2}\overset{{ – 2}}{\mathop {\text{O}}} \left({\text{g}} \right) \to \overset{{ + 2, + 3}}{\mathop {{\text{F}}{{\text{e}}_3}}} \overset{{ – 2}}{\mathop {{{\text{O}}_4}}} \left({\text{s}} \right) + 4\overset{0}{\mathop {{{\text{H}}_2}}} \left({\text{g}} \right)\)

III. Dihydrogen can be displaced from acids by various metals, including ones that do not react with cold water.

\(\overset{{\text{0}}}{\mathop {{\text{Mg}}}} \left({\text{s}} \right) + 2\overset{{ + 1}}{\mathop {\text{H}}} \overset{{{\text{-1}}}}{\mathop {{\text{Cl}}}} \left({\text{g}} \right) \to \overset{{ + 2}}{\mathop {{\text{Mg}}}} \overset{{ – 1}}{\mathop {{\text{C}}{{\text{l}}_2}}} \left({{\text{aq}}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)

\(\overset{{\text{0}}}{\mathop {{\text{Fe}}}} \left({\text{s}} \right) + 2\overset{{ + 1}}{\mathop {\text{H}}} \overset{{{\text{-1}}}}{\mathop {{\text{Cl}}}} \left({{\text{aq}}} \right) \to \overset{{ + 2}}{\mathop {{\text{Fe}}}} \overset{{ – 1}}{\mathop {{\text{C}}{{\text{l}}_2}}} \left( {{\text{aq}}} \right) + {\overset{0}{\mathop {\text{H}}} _2}\left({\text{g}} \right)\)

IV. Very less reactive metals such as silver \(\left({{\text{Ag}}} \right)\) and gold \(\left({{\text{Au}}}\right)\) which may occur in the native state, do not react even with dilute hydrochloric acid

4. Disproportionation Reactions

A reaction in which the same species is simultaneously oxidised as well as reduced is called a disproportionation reaction. For such redox reactions to occur, the reacting species must contain an element that has at least three oxidation states. The element in the reacting species is in the intermediate oxidation state, whereas the higher and lower oxidation states are available for oxidation and reduction. For example, hydrogen peroxide \(\left({{{\text{H}}_2}{{\text{O}}_2}} \right)\) decomposition is a disproportionation reaction in which the oxygen atom is disproportionated.

Phosphorus, sulphur and chlorine undergo disproportionation reactions in an alkaline medium as shown below:

Applications of Redox Reactions

1. Redox reactions are used in the determination of equivalent masses of the element.
2. It is used in electrochemistry to isolate reactive metals like sodium, potassium, magnesium, calcium, aluminum, etc., by the electrolysis of their fused salts.
3. It is used in the electrorefining of metals such as copper, aluminum, silver, gold, tin, etc.
4. It is used in electroplating metals by applying a coating of superior metal over an interior metal surface.
5. It is also used in the preparation of chemicals and organic compounds.
6. Photosynthesis taking place in plants in the presence of sunlight is a redox reaction.

Non-Redox Reactions

The reactions in which the oxidation numbers of the reactants and products do not change are called non-redox reactions. In ordinary life, this reaction occurs both spontaneously and artificially. However, the reactant does not change oxidation states before or after the reaction. As a result, it can be classified as a non-redox reaction because there is no change in the oxidation states. Double displacement reactions are non-redox reactions.

Types of Non-redox reactions are as follows:

1. Neutralisation Reaction: The reaction in which acid and base react with each other and forms neutral salt and water is known as neutralisation reaction.
   For example,
   \({\text{HCl}} + {\text{NaOH}} \to {\text{NaCl}} +{{\text{H}}_2}{\text{O}}\)
   \({{\text{H}}^ + } + {\text{O}}{{\text{H}}^ – } \to {{\text{H}}_2}{\text{O}}\)

2. Precipitation Reaction: When two solutions containing soluble salts are mixed, a precipitation reaction occurs, resulting in the formation of an insoluble salt.
For example,
\({\text{AgN}}{{\text{O}}_3} + {\text{HCl}} \to {\text{AgCl}} + {\text{HN}}{{\text{O}}_3}\)
\({\text{A}}{{\text{g}}^ + } + {\text{C}}{{\text{l}}^ – } \to {\text{AgCl}}\)

3. Decomposition Reaction: A decomposition reaction occurs when a compound is broken down into two or more simpler substances.
For example,
\({\text{CaC}}{{\text{O}}_3} \to {\text{CaO}} + {\text{C}}{{\text{O}}_2}\)

4. Combination Reaction: A combination reaction occurs when two reactants combine to form one product.
For example,
\({\text{CaO}} + {{\text{H}}_2}{\text{O}} – {\text{Ca}}{\left({{\text{OH}}} \right)_2}\)

Redox Titration Principle

In redox systems, the titration method can be adapted to determine the strength of a reductant/oxidant using a redox-sensitive indicator. A reducing substance is titrated with a standard solution of an oxidising agent, or an oxidising substance is titrated with the standard solution of the reducing agent in the oxidation-reduction titration method.

The oxidation-reduction titrations are based on the principle that the oxidation process involves the loss of electrons, whereas the reduction process involves the gain of electrons.

Types of Redox Titration

Depending upon the nature of the oxidising agent, these are divided into the following categories:

1. Potassium permanganate titrations
2. Potassium dichromate titrations
3. Ceric sulphate titrations
4. Iodimetric titrations
5. Iodometric titrations

Potassium Permanganate Titrations

In these titrations, reducing agents like \({\text{FeS}}{{\text{O}}_4},\) mohr’s salt \(\left[{{{\left({{\text{N}}{{\text{H}}_4}} \right)}_2}{\text{S}}{{\text{O}}_4}.{\text{FeS}}{{\text{O}}_4}.6{{\text{H}}_2}{\text{O}}} \right],\) oxalic acid \({\left({{\text{COOH}}} \right)_2},\) oxalates \({\left({{\text{COONa}}} \right)_2},{{\text{H}}_2}{{\text{O}}_2},\) etc., are directly titrated against \({\text{KMn}}{{\text{O}}_4}\) in an acidic medium. In this reaction, \({\text{KMn}}{{\text{O}}_4}\) acts as an oxidising agent in the acidic medium, while oxalic acid, oxalate, mohr’s salt, ferrous sulphate acts as a reducing agent.

For example

I. Oxidation of Ferrous Salts

\(\mathop {\mathop {{\rm{5F}}{{\rm{e}}^{2 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Ferrous}}\,{\rm{ion}}} \, + \,\mathop {\mathop {{\rm{MnO}}_4^ – \left( {{\rm{aq}}} \right)}\limits_{{\rm{Permaganate}}\,{\rm{ion}}} }\limits_{\rm{ }} \, + \,8{{\rm{H}}^ + }\left( {{\rm{aq}}} \right)\, \to \,\mathop {\mathop {5{\rm{F}}{{\rm{e}}^{3 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Ferric}}\,{\rm{ion}}} \, + \,{\rm{M}}{{\rm{n}}^{2 + }}\left( {{\rm{aq}}} \right)\, + \,4{\mkern 1mu} {{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right)\)

II. Oxidation of Oxalates

Potassium Dichromate Titrations

In these titrations, reducing agents like \({\text{FeS}}{{\text{O}}_4},\) mohr’s salt \(\left[{{{\left({{\text{N}}{{\text{H}}_4}} \right)}_2}{\text{S}}{{\text{O}}_4}.{\text{FeS}}{{\text{O}}_4}.6{{\text{H}}_2}{\text{O}}} \right],\) oxalic acid \({\left({{\text{COOH}}} \right)_2},\) oxalates \({\left({{\text{CONa}}} \right)_2},{{\text{H}}_2}{{\text{O}}_2},\) etc., are directly titrated against \({{\text{K}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_7}\) as the oxidising agent in the presence of dilute sulphuric acid.

For example,

I. Oxidation of Ferrous Salts

\(6{\text{F}}{{\text{e}}^{2 + }}\left({{\text{aq}}} \right) + {\text{C}}{{\text{r}}_2}{{\text{O}}_7}^{2 – }\left({{\text{aq}}} \right) + 14{{\text{H}}^ + }\left({{\text{aq}}} \right) \to 2{\text{C}}{{\text{r}}^{3 + }}\left({{\text{aq}}} \right) + 6{\text{F}}{{\text{e}}^{3 + }}\left({{\text{aq}}} \right) + 7{{\text{H}}_2}{\text{O}}\left({\text{l}} \right)\)

II. Oxidation of Mohr’s Salt

Mohr’s salt is a double salt of \({\left({{\text{N}}{{\text{H}}_4}} \right)_2}{\text{S}}{{\text{O}}_4}\) and \({\text{FeS}}{{\text{O}}_4}.\) Out of these two salts,\({\text{C}}{{\text{r}}_2}{{\text{O}}_7}^{2 – }\) oxidises \({\text{FeS}}{{\text{O}}_4}\) to \({\text{F}}{{\text{e}}_2}{\left({{\text{S}}{{\text{O}}_4}} \right)_3}.\)

Ceric Sulphate Titrations

In these titrations, the reducing agents such as nitrites,\({\text{F}}{{\text{e}}^{2 + }}\) salts, oxalates, \({\text{C}}{{\text{u}}^ + }\) salts, arsenites, etc., are directly titrated against ceric sulphate. In this reaction, \({\text{Ce}}{\left({{\text{S}}{{\text{O}}_4}}\right)_2}\) acts as an oxidising agent.

For example,

I. Oxidation of arsenites \(\left({{\text{As}}{{\text{O}}_3}^{3 – }} \right)\) to arsenates \(\left({{\text{As}}{{\text{O}}_4}^{3 – }} \right)\)

\(\mathop {\mathop {{\rm{ASO}}_3^{3 – }}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Arsenite}}\,{\rm{ion}}} \, + \,\mathop {\mathop {2{\rm{C}}{{\rm{e}}^{4 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Ceric}}\,{\rm{ion}}} \, + \,{{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right)\, \to \mathop {\,\mathop {{\rm{AsO}}_4^{3 – }}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Arsenate}}\,{\rm{ion}}} \, + \,\mathop {\mathop {2{\rm{C}}{{\rm{e}}^{3 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Cerousc}}\,{\rm{ion}}} \, + \,2{{\rm{H}}^ + }\left( {{\rm{aq}}} \right)\)

II. Oxidation of ferrous salts:

\(\mathop {\mathop {{\rm{F}}{{\rm{e}}^{2 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Ferrous}}\,{\rm{ion}}} \, + \,\mathop {\mathop {{\rm{C}}{{\rm{e}}^{4 + }}}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Ceric}}\,{\rm{ion}}} \, \to \,\mathop {\mathop {{\rm{F}}{{\rm{e}}^{3 + }}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Ferric}}\,{\rm{ion}}} }\limits_{\rm{ }} \, + \,\mathop {\mathop {{\rm{C}}{{\rm{e}}^{3 + }}}\limits_{\rm{ }} \left( {\rm{l}} \right)}\limits_{{\rm{Cemus}}\,{\rm{ion}}} \)

Iodimetric Titrations

These titrations involve the direct use of iodine as the oxidising agent (in neutral or slightly acidic medium) using starch as an indicator. Thiosulphates, sulphites, arsenites, and antimonites are examples of reducing agents utilised in these titrations.

For example,

I. Oxidation of Thiosulphates

\({{\rm{I}}_2}\left( {{\rm{aq}}} \right)\, + \,\mathop {\mathop {2\;{{\rm{S}}_2}{\rm{O}}_3^{2 – }}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Thiosulphate}}\,{\rm{ion}}} \, \to \,2{{\rm{I}}^ – }\left( {{\rm{aq}}} \right)\, + \,\mathop {\mathop {{{\rm{S}}_4}{\rm{O}}_6^{2 – }}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Tetrathionate}}\,{\rm{ion}}} \)

II. Oxidation of Arsenite

\({{\rm{I}}_2}\left( {{\rm{aq}}} \right)\, + \,\mathop {\mathop {{\rm{AsO}}_3^{3 – }}\limits_{\rm{ }} \left( {{\rm{aq}}} \right)}\limits_{{\rm{Arsenite}}\,{\rm{ion}}} \, + \,{{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right)\, \to \,2{{\rm{I}}^ – }\left( {{\rm{aq}}} \right)\, + \,\mathop {\mathop {{\rm{AsO}}_4^{3 – }\left( {{\rm{aq}}} \right)}\limits_{{\rm{Arsenate}}\,{\rm{ion}}} }\limits_{\rm{ }} \, + \,2{{\rm{H}}^ + }\left( {{\rm{aq}}} \right)\)

Iodometric Titrations

All titrations in which iodine is liberated from potassium iodide with the help of an oxidising agent when treated against a standard solution of sodium thiosulphate are called iodometric titrations.

Iodometric titrations are carried out in two steps:

In the first step, oxidising agents such as \({\text{KMn}}{{\text{O}}_4},{{\text{K}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_{7,}}{\text{CuS}}{{\text{O}}_4},\) peroxides, and other oxidising agents are treated with an excess of \({\text{KI}}\) to liberate \({{\rm{I}}_2}\) rapidly and quantitatively.

For example,

\(2{\text{MnO}}_4^ – \left({{\text{aq}}} \right) + 16{{\text{H}}^ + }\left({{\text{aq}}} \right) + 10{{\text{I}}^ – }\left({{\text{aq}}} \right) \to 2{\text{M}}{{\text{n}}^{2 + }}\left({{\text{aq}}} \right) + 5{{\text{I}}_2}\left({\text{s}} \right) + 8{{\text{H}}_2}{\text{O}}\left({\text{l}} \right)\)

\({\text{C}}{{\text{r}}_2}{\text{O}}_7^{2 – }\left({{\text{aq}}} \right) + 14{{\text{H}}^ + }\left({{\text{aq}}} \right) + 6{{\text{I}}^ – }\left({{\text{aq}}} \right) \to 2{\text{C}}{{\text{r}}^{3 + }} + 3{{\text{I}}_2}\left({\text{s}} \right) + 7{{\text{H}}_2}{\text{O}}\left({\text{l}} \right)\)

\(2{\text{C}}{{\text{u}}^{2 + }}\left({{\text{aq}}} \right) + 4{{\text{I}}^ – }\left({{\text{aq}}} \right) \to {\text{C}}{{\text{u}}_2}{{\text{I}}_2}\left({\text{s}} \right) + {{\text{I}}_2}\left({\text{s}} \right)\)

Applications of Redox Titration

Following are the applications of redox titration:

1. Redox titration is used to determine the medium and concentration of elements.
2. It is used for the analysis of organic analytes. For example, chemical oxygen demand (COD) of natural water and wastewater is determined by redox titration.
3. It is employed in the extraction of metals from ores and the combustion of fuel in metallurgical processes.
4. Titration is a frequently used analytical technique in the food industry.
5. In natural waters, such as lakes and rivers, the level of dissolved oxygen is determined by redox titration.
6. The chlorination of public water sources is one of the most important industrial applications of redox titrations.

Summary

A redox reaction is defined as a reaction in which both reduction and oxidation occur simultaneously. Redox reactions are classified into four types. Redox titration is used to determine the medium and concentration of elements. It is also used to extract metals from ores and the combustion of fuel in metallurgical processes. Redox titration is a frequently used analytical technique in the food industry.

FAQs About Types of Redox Reactions

Let’s look at some of the commonly asked questions about types of redox reactions:

Q.1. Are all reactions redox reactions?
Ans: Not all chemical reactions are redox reactions.

Q.2. What are the types of redox titration?
Ans: There are five types of redox titration reactions, these are;
1. Potassium permanganate titrations
2. Potassium dichromate titrations
3. Ceric sulphate titrations
4. Iodimetric titrations
5. Iodometric titrations

Q.3. What are the applications of Redox reactions?
Ans: Following are the applications of Redox reactions:

1. Redox reactions are used in the determination of equivalent masses of the element.
2. Electrochemistry uses it to isolate reactive metals like sodium, potassium, magnesium, calcium, aluminum, etc., by the electrolysis of their fused salts.
3. It is used for electrorefining metals such as copper, aluminum, silver, gold, tin, etc.
4. It is used in electroplating metals by applying superior metal coating over an interior metal surface.
5. It is also used in the preparation of chemicals and organic compounds.
6. Photosynthesis taking place in plants in the presence of sunlight is a redox reaction.

Q.4. What is a displacement redox reaction? Give an example?
Ans: A displacement reaction is defined as a reaction in which an atom or ion in a compound is replaced by an atom or ion of a different element.
For example, displacement of copper metal from copper sulphate solution by zinc is a displacement redox reaction.
\(\mathop {\mathop {{\rm{Cu}}}\limits^{ + 2} \mathop {\rm{S}}\limits^{ + 6} \mathop {{{\rm{O}}_4}}\limits^{ – 2} \left( {{\rm{aq}}} \right)}\limits_{{\rm{ Copper}}{\kern 1pt} {\rm{sulphate }}}  + \mathop {\mathop {{\rm{Zn}}}\limits^0 \left( {\rm{s}} \right)}\limits_{{\rm{ Zinc }}}  \to \mathop {\mathop {{\rm{Cu}}}\limits^0 \left( {\rm{s}} \right)}\limits_{{\rm{ Copper }}}  + \mathop {\mathop {{\rm{Zn}}}\limits^{ + 2} \mathop {\rm{S}}\limits^{ + 6} \mathop {{{\rm{O}}_4}}\limits^{ – 2} }\limits_{{\rm{ Zinc}}{\kern 1pt} {\rm{sulphate }}}\)

Q.5. What is a combination redox reaction?
Ans: In a combination reaction, two atoms or molecules combine to form a single molecule. In a combination reaction, one or both reactants must be an element. Such a reaction is called a combination reduction reaction.

Q.6. What is the most common redox reaction?
Ans: The most common redox reactions are;
1. Combination reaction
2. Decomposition reaction
3. Displacement reactions
4. Disproportionation reactions

Q.7. What are \(4\) types of redox reactions?
Ans: The four types of redox reactions are :
1) Combination reaction
2) Decomposition reaction
3) Displacement reactions
4) Disproportionation reactions

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