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Redox Potential

In any chemical reaction, some substances get oxidized while others get reduced. According to the old concept, a chemical reaction is one in which there is the addition of oxygen or removal of hydrogen.
For example,

S + O2 SO2

In the given reaction, oxygen is added to sulfur to form sulfur dioxide. Hence sulfur gets oxidized to sulfur dioxide.

Zn + 2HCl ZnCl2 + H2

While in this reaction, hydrochloric acid gets oxidized to zinc chloride.

"The addition of electronegative elements or removal of electropositive element is also termed as oxidation". According to the modern concept of oxidation, a reaction in which any element or atom donates a number of electrons and gets converted to a cation is termed as oxidation reaction. Hence the oxidation number increases in an oxidation reaction. Removal of electrons can take place from,

An atom and form cation: Na Na+ + e-
A cation and form another cation with more charge: Sn 2+ Sn4+ + 2e-
An anion and form atom: 2Cl- Cl2
A complex anion: [Fe(CN)6]4- [Fe(CN)6]3- + e-

Hence in oxidation process:

  1. Positive charges increase and negative charges decrease.
  2. Oxidation number increases.
  3. However in reduction reaction, there is an addition of hydrogen or removal of oxygen during reaction.

H2 + Cl2
2HCl (Addition of hydrogen to chlorine)
ZnO + H2 Zn + H2O (Removal of oxygen from zinc oxide)

"A reaction corresponding to the addition of an electropositive element and the removal of an electronegative element is also termed as reduction".

According to the modern concept, a chemical reaction in which an element or ion accepts electrons to form anions is termed as reduction. For example,

Cl2 + 2e- 2Cl-
Na+ + e-
Sn4+ + 2e-

Hence in a reduction reaction,
  1. Positive charge decreases and negative charge increases.
  2. Oxidation number decreases.

The substance which gets reduced is known as the oxidizing agent and the substance which gets oxidized is termed as reducing agent.

If we have pure carbon and a complete combustion, carbon is converted into carbon dioxide. The reaction can be given as below.

C(s) + O2(g) CO2(g)

Here, the oxidation number of carbon changes from 0 in C(s) to +4 in CO2. Simultaneously, oxygen is reduced from 0 in O2 (g) to -2 in CO2 (g).

C (s) + O2 (g) CO2 (g)


Here, the oxidation number of carbon changes from zero in C(s) to +4 in CO2. Simultaneously, oxygen is reduced from zero in O2 (g) to -2 in CO2 (g).

Related Calculators
Redox Reaction Calculator Elastic Potential Energy Calculator
Gravitational Potential Energy Calculator Nernst Potential Calculator

Redox Equations

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In any chemical reaction, the species donating electrons forms the cation and the same electron is accepted by another to form the corresponding anion. In other words, a substance gets oxidized only in the presence of an oxidizing agent which gets reduced at the same time.

Hence oxidation and reduction reactions occur simultaneously. Such reactions are known as redox reactions and the balanced chemical equation of a redox reaction is called redox equation. Reduction and oxidation reactions are complementary to each other. For example,

Fe2+ + MnO4- + 8H+ Fe3+ + Mn2+ + 4H2O

In a given redox reaction, one oxidation and one reduction take place,

Fe2+ Fe3+ + e- (Oxidation half reaction)
MnO4-+ 8H+ + 5e-
Mn2+ + 4H2O (Reduction half reaction)

Here Fe2+ lose one electron and form Fe3+, hence it’s a reducing agent and gets oxidized. However MnO4- is an oxidizing agent and gets reduced to Mn2+. Some substances act as reducing agents as well as oxidizing agents depending upon the reaction conditions. For example, Hydrogen peroxide (H2O2) acts as an oxidant as well as a reductant. In the reaction of hydrogen peroxide with bromide ion, H2O2 acts as the oxidant.

2Br- + H2O2 2OH- + Br2

While in the reaction of H2O2 with silver oxide, H2O2 acts as the reductant.

Ag2O + H2O2 2Ag + H2O + O2

In the same way, nitrous acid and sulfur dioxide also act as both oxidant and reductant.
Some common reducing agents and oxidizing agents are as follows:

Reductant Oxidant
Substance Product
Fe2+ Fe3+
MnO4- Mn2+(acidic)
AsO32- AsO42- Cr2O72- Cr3+
Cu2O42- CO2 MnO4- Mn4+(basic)
Cu Cu2+ ClO3-
Zn2+ OCl-

Let us take a simple example and try writing the redox equations for it. There are two main steps in writing redox equations. They are:

  • Oxidation half reaction
  • Reduction half reaction

Copper metal reacts with aqueous silver nitrate solution to form aqueous Cu (II) ions in solution and solid silver metal.

Copper metal + Aqueous Silver (I) ions aqueous Cu (II) ions + Silver metal

The half reaction for the oxidation process is the loss of two electrons from the more electro positive, copper metal.

That is,

Cu (s) Cu 2+ (aq) + 2e-

In the corresponding reduction process, a silver ion gains a single electron to form silver metal.

That is,

Ag+ (aq) + e- Ag (s)

Now, let us try balancing redox reactions. All the electrons lost in the oxidation half reaction must be accounted for by the redox half reactions. So, in order to balance redox reactions, two silver ions must be accepting the two electrons lost by the Cu to form two atoms of Ag(s).

So, the reduction equation becomes,

2 Ag+ (aq) + 2 e- 2 Ag (s)

We get this by multiplying the reduction half reaction by 2.

Now, the final step is to add the two half equations:

Cu (s) Cu 2+ (aq) + 2 e- oxidation half reaction

2 Ag + (aq) + 2 e- 2 Ag (s) reduction half reaction

Cu (s) + 2 Ag + (aq)
Cu 2+ (aq) + 2 Ag (s) net balanced redox reaction

This is the process for balancing redox equations.

There are a number of steps involved in the redox reactions. They are as follows.

Electron Transfer

In oxidation reduction reactions or redox reactions the electrons are transferred from one species to another, the species which loses the electrons is oxidized while the one which gains electrons is reduced. There are numerous examples of redox reactions in our every day life. Some of them are corrosion reactions, reactions in batteries, and some are metabolic reactions in the body.


Oxidation is the loss of electrons by a species, leading to an increase in the oxidation number of one or more atoms.


Reduction is the gain of electrons by a species, leading to a reduction in the oxidation number of one or more atoms.

Oxidizing Agent

Oxidizing agent is the chemical species causing the oxidation. This species is reduced and can also be called the electron acceptor.

Reducing Agent

Reducing agent is the species causing the reduction. This species is oxidized and can be called the electron donor.

A Redox Reaction: Silver Coating Copper

Silver nitrate + copper silver + copper nitrate

redox reactions

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Fuel Cell

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Fuel cell is a device which converts the energy produced during the combustion of fuels like hydrogen, methane etc., directly into electrical energy through oxygen and other oxidizing agents.

In general hydrogen is used as a common fuel for these types of cells. Fuel cells require a constant source of fuel and oxygen and they can produce electricity as long as these inputs are supplied.

A hydrogen-oxygen fuel cell is one of the best examples of fuel cells. This cell was used as the primary source of electrical energy on the Apollo moon flight. This cell consists of porous carbon electrodes containing suitable catalyst like finally divided platinum and palladium which is incorporated in them. Concentrated KOH or NaOH solution is placed between the electrodes to act as the electrolyte. Hydrogen and oxygen gas is bubbled through the porous electrodes into the electrolyte solution.

The electrode reactions are as follows,

At anode : 2H2(g) + 4OH-(aq) 4H2O (l) + 4e- (Oxidation
At cathode: O2(g) +2H2O(l) +4e-
4OH-(aq) (Reduction reaction)
Redox equation:2H2(g) + O2(g)

Thus, the reactant has to be fed continually in to the cell to get a continuous supply of electricity.

Advantages of A Fuel cell
  • The fuel cell has an unlimited lifetime as the supply of electricity depends upon the sources of fuel. They operate at 343-413 k temperature and give a potential of about 0.9V.
  • The fuel cells are expected to have efficiency of about 100% but practically they show only 60-70% efficiency. Still they are much better than thermal power plants which show only 40% efficiency.
  • The by product of fuel cells is water with electrons which does not cause any pollution.
Limitation of fuel cell
  • It’s difficult to provide contact among three phases- solid, liquid and gas.
  • The catalysts (Pt & Pd) are expensive metals and it’s difficult to handle gaseous fuels at low temperature or high pressure.

Solid Oxide Fuel Cell

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A solid oxide fuel cell (SOFC) is a type of fuel cell which produces electricity from oxidizing a fuel. Like in the hydrogen-oxygen fuel cell, the electrolyte is not a solution of potassium hydroxide or sodium hydroxide but a solid oxide or ceramic acts as the electrolyte.

The main advantages of SOFC are high efficiency, stability with fuel flexibility and low emissions as well as relatively low cost. However, the disadvantage associated with SOFC is the high operating temperature which results in longer start-up times and mechanical and chemical compatibility issues. They work at a very high temperature (500-1000°c) so no expensive catalyst is required for the reaction.

The solid electrolyte like ceramics are not electrically active but because of high temperature ranging from 500 to 1,000°c it becomes active. At cathode, reduction of oxygen takes place and forms oxygen ions. These oxygen ions are diffused toward the anode through the solid oxide electrolyte. At the anode, oxygen ions can electrochemically oxidize the fuel.

Anode Reaction: 2H2 + 2O2- 2H2O + 4e
Cathode Reaction: O2 + 4e
Overall Cell Reaction: 2H2 + O2

The water produced as by-product with two electrons. These electrons are responsible for the flow of electricity in the external circuit. The main disadvantage of SOFC systems is its high operating temperatures. It's difficult to build up the potential for carbon dust on the anode which slows down the internal reforming process. Hence in place of carbon dust the copper-based cermet is more efficient in performance.

Solid oxide fuel cells take a long time to start which makes it less useful for mobile applications. Despite these disadvantages it provides an advantage by removing the need of expensive metals as catalyst and thereby reducing the cost.

Redox Reactions Examples

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Given below are some of the redox reaction examples along with their solutions.

Solved Examples

Question 1: Identify the oxidation and reduction in the reaction of zinc metal with hydrochloric acid.

The reaction between zinc and hydrochloric acid can be represented as below:

Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g)

The net ionic equation can be written as,

Zn (s) + 2H+ (aq) Zn 2+ (aq) +H2 (g)

Now, let us determine the oxidation numbers.

Zn (s) and H 2 (g) are elements. So, their oxidation numbers are zero each. H+ has an oxidation number of +1 and Zn 2+ that of 2+.

Zn loses electrons to form Zn 2+. This is the oxidation reaction. H+ accepts electrons to form H2. This is the reduction reaction.

redox reaction

Question 2: When natural gas burns in a furnace, carbon dioxide and water form. Identify the oxidation and reduction in this reaction.

CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (g)

oxidation reduction

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