What are redox reactions?
Redox reactions are oxidation-reduction reactions where reacting substances are changed their oxidation number or state. Oxidation may be defined as a process where the oxidation number of reacting substances increases. Similarly, reduction may be defined as a process where the oxidation number of reacting substances decreases. In another word, redox reactions are chemical reactions where one or more electrons are transferred between two reacting substances like atoms or ions participating in it. Redox reactions have wide industrial applications like electricity generation, purification or extraction of metal like zinc, aluminum, chromium, copper, etc by electrolysis, combustion of fuel in the fuel cell. Redox reaction based on oxidation number and electronic concept in chemistry given below the picture.
Oxidation and reduction are always found to go hand in hand during a redox reaction. In a redox reaction, an oxidant is reduced and a reductant is simultaneously oxidized. Indirect redox reactions take place in the electrochemical cell for conversion of chemical energy into electrical energy by transfer of electron. The electrolytic cell is another example where electrical energy is converted to chemical energy by the transfer of electrons.
Types of redox reactions
Redox reactions are generally following types
- Decomposition reaction
- Combination reaction
- Displacement reaction
- Disproportionation reaction
Examples of decomposition reaction
The chemical reaction where one reactant can be breakdown into two or more products is called decomposition reaction. Decomposition can be represented by, AB → A + B. Breakdown of hydrogen peroxide or water into hydrogen and oxygen are examples of redox reactions where decomposition can occur.
Examples of combination reaction
In a combination reaction, a metal can be reacted with non-metal to form an ionic crystalline solid. For example, when magnesium metal burns in oxygen to produce magnesium oxide. Here two electrons can be transferred between the two reacting substances like magnesium and oxygen.
Examples of displacement reaction
A displacement reaction is a type of chemical reaction where one or more atoms can be displaced by another atom. For example, the electron charge is transferred from the zinc atom to the copper atom in this electrochemical cell.
Examples of disproportionation reaction
A disproportionation reaction is a type of redox reaction where one element of a reactant in a particular oxidation state can be simultaneously oxidized and reduced to form two products. The chemical reaction of phosphorus with sodium hydroxide where phosphorus can simultaneously be oxidized and reduced.
Examples of redox reactions
Oxidation of iron by potassium permanganate
Oxidation of ferrous ion to ferric ion by acidic potassium permanganate is an example of a redox reaction where oxidation and reduction occur simultaneously. The half-reaction is given below the picture,
Reaction of potassium dichromate with potassium iodide
Oxidation of potassium iodide by potassium dichromate in dilute sulfuric acid medium is another example of redox reaction. Here chromium(+6) is reduced to chromium(+3) and iodine oxidized to form an elementary iodine molecule.
|Cr2O7-2 + 14H+ + 6I– → 2Cr+3 + 3I2 + 7H2O|
|2I– → I2 + 2e|
|Cr2O7-2 + 14H+ + 6e → 2Cr+3 + 7H2O|
Redox electrode examples
- When oxidation (loss of electrons) reactions take place on the redox electrode, the chemical potential of the electrode is called oxidation potential.
- When reduction (gain of electrons) takes place on a redox electrode is called reduction potential. If the oxidation potential of an electrode = x volt, then its reduction potential = – x volt.
Two half cells chemical reactions of ferric ion and ceric ion in the presence of dilute sulfuric acid are examples of redox electrode. Two electrode potentials in balancing chemical equations can be given below the table.
|Redox electrode examples|
|Ce+4 + Fe+2 ⇄ Ce+3 + Fe+3|
|Half cell reaction||Electrode potential|
|Ce+4 + e ⇄ Ce+3||1.44 volt|
|Fe+2 ⇄ Fe+3 + e||-077 volt|
Applications of redox reaction
Redox reactions have wide industrial application, it is used for constructing electrochemical and electrolytic cells in our daily life.
Redox reaction in electrochemistry
Redox reactions take place in the electrochemical cell converted chemical energy into electrical energy by electron transfer reaction. Therefore, the electrochemical or simply chemical cell is the device in which energy is produced due to the indirect redox reaction and converted electricity for mankind. We used Galvanic and Voltaic cell according to the name of the scientist, Luigi Galvani, and Alessandro Volta.
Application of redox reaction in metallurgy
An electrolytic cell is a device where a non-spontaneous redox reaction takes place by an external source of electricity. The electrolysis process occurring in the electrolytic cells. The very useful application of electrolysis is found in commercial electroplating and purification of metals. For copper, an electroplating reaction is carried out by dipped copper into the suspension of an aqueous copper sulfate solution. The electrolysis or redox reaction at the two electrodes can be represented as follows,
|Copper electroplating reaction|
|At cathode||Cu+2 + 2e → Cu|
|At anode||Cu → Cu+2 + 2e|
Application of fuel cell
A fuel cell is electrochemical cells where chemical energy available from the fuel like methane, hydrogen, solid oxide, alkaline, etc converted into electrical energy by redox reactions. At the cathodes of the cell, oxygen is dissolved to form OH– ions. The cell has low voltage but capable of supplying current for long periods.
Such cells provide power at comparatively low cost and low maintenance. The efficiency depends particularly on the surface of the electrodes. It serves as the chemical catalyst for the attainment of chemical equilibrium. The electric polarization due to the adoption and poisoning of the redox electrode. It reduces cell performance considerably. The density of the electrolyte and the temperature have a great influence on the potential of the redox reactions.