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Redox Reaction (Oxidation-Reduction Reaction)

What is a Redox Reaction?

A redox reaction is a chemical reaction in which the atoms change their oxidation numbers. During a redox reaction, some atoms lose electrons and are said to be oxidized – a process known as oxidation. On the other hand, some atoms gain electrons and are said to be reduced – a process known as reduction. Therefore, both REDuction and OXidation take place simultaneously, and hence the term redox. Redox reactions usually occur in one of two environments: acidic or basic [1-3].

Redox Reaction

Identifying Redox Reactions

A way to recognize a redox reaction is that it must involve a change in the oxidation number of the two elements involved in the reaction. The element that is oxidized increases its oxidation number, while the element that is reduced decreases its oxidation number. By checking this change, one can quickly determine and tell whether the reaction is redox or not.

What Happens in Redox Reactions?

In a redox reaction, electrons move from one atom to another, which changes the oxidation number or state of the two atoms. The oxidation state of an element corresponds to the number of electrons that an atom loses, gains, or appears to use when joining with other atoms in compounds.

Redox Half-Reaction

A half-reaction is the part of an overall reaction that represents, separately, either an oxidation or a reduction. Two half-reactions, one oxidation and one reduction, are necessary to describe a redox reaction completely [4].


When a nickel (Ni) strip is placed in an aqueous solution of copper (II) sulfate (CuSO4), an immediate reaction occurs. Copper metal begins to deposit on the strip. The only source for metallic copper in this system is the copper (II) ions (Cu2+) in solution. The reaction is represented as follows.

Cu2+ (aq.) + 2e → Cu(0) (s)

Here, the Cu2+ ion gains two electrons and is reduced to Cu (0). So, the above equation represents the reduction part of a redox reaction, and hence, it is called a redox half-reaction.

On the other hand, nickel releases two electrons to form nickel (II) ions (Ni2+). These are the two electrons that go to copper.

Ni(0} (s) → Ni2+ (aq.) + 2e

Thus, nickel is oxidized, and the above half-reaction represents the oxidation part of the redox reaction.

Balancing Redox Reaction using Half-Reaction [5,6]

The half-reaction can be used to balance the redox reaction. Consider an example of the combustion redox reaction between magnesium (Mg) and oxygen (O2), resulting in magnesium oxide (MgO). The unbalanced equation is given by:

Mg (s) + O2 (g) → MgO (s)

Mg loses two electrons and becomes Mg2+ cation.  Its oxidation number goes from zero to +2. Hence, it is oxidized. The oxidation part of the reaction is written as:

Mg (s) → Mg2+ + 2 e

On the other hand, O gains two electrons from Mg and becomes O2-, thereby changing its oxidation number from zero to -2. Hence, oxygen is reduced, and the reduction part of the reaction is:

½ O2 (g) + 2 e → O2-

Adding the two equation together:

Mg (s) + ½ O2 (g) → MgO (s)

Multiplying both sides by 2:

2 Mg (s) + O2 (g) → 2 MgO (s)

This balanced equation represents the reaction between magnesium and oxygen.

Examples of Redox Reaction

Below are some examples of the redox reaction with explanations [1-9].

Redox Reaction Examples
  • The reaction between lead oxide (PbO) and ammonia (NH3) results in nitrogen (N2), water (H2O), and lead (Pb).

3 PbO (s) + 2 NH3 (g) → N2 (g) + 3 H2O (l) + 3 Pb (s)

The oxidation number of lead decreases from +2 to 0, and Pb is reduced. The oxidation number of nitrogen increases from -3 to zero, and N is oxidized.

  • The reaction between carbon (C) and Iron (III) oxide (Fe2O3) results in iron (Fe) and carbon dioxide (CO2).

3 C (s) + 2 Fe2O3 (s) → 4 Fe (s) + 3 CO2 (g)

The oxidation number of carbon goes from zero to +4, and C is oxidized. The oxidation number of iron decreases from +3 to zero, and Fe is reduced.

  • In a thermite reaction, iron (Fe) atoms in ferric oxide (Fe2O3) lose oxygen (O) atoms to aluminum (Al) atoms, producing aluminum oxide (Al2O3).

Fe2O3 (s) + 2 Al (s) → Al2O3 (s) + 2 Fe (l)

The oxidation number of aluminum goes from zero to +3 (oxidation), and that of iron decreases from +3 to zero (reduction).

  • Chromium (Cr) reacts with hydrochloric acid (HCl) to give chromium chloride (CrCl2) and hydrogen (H2) gas.

Cr (s) + 2 HCl (aq.) → CrCl2 (aq.) + H2 (g)

The oxidation number of chromium increases from zero to +2 (oxidation), and that of hydrogen decreases from +1 to zero (reduction).

  • Sulfur dioxide (SO2) reacts with carbon monoxide (CO) to form sulfur (S) and carbon dioxide (CO2).

SO2 (g) + 2 CO (g) → S (g) + 2 CO2 (s)

The oxidation number of sulfur decreases from +4 to zero (reduction), and that of carbon increases from +2 to +4 (oxidation).

  • In the reaction between manganese (IV) oxide (MnO2) and hydrochloric acid (HCl), the products are manganese (II) chloride (MnCl2), chlorine (Cl2) gas, and water (H2O).

MnO2 (s) + 4 HCl (aq.) → MnCl2 (aq.) + Cl2 (g) + 2 H2O (l)

The oxidation number of manganese decreases from +4 to +2 (reduction), and that of chlorine increases from -1 to zero (oxidation).

Types of Redox Reactions

The following list represents the four types of redox reactions.

  • Combination reactions – two reactants combine to form a product
  • Decomposition reactions – a compound decomposes into two or more products
  • Single-replacement (displacement) reactions – one atom replaces another atom to form two products
  • Disproportionation reactions – a decomposition reaction in which the same atom is simultaneously oxidized and removed

Common Uses of Redox Reaction in Everyday Life

Redox reactions have plenty of applications in industry and daily life with many benefits and effects [10].

  • Conversion of food and oxygen to carbon dioxide and water (cellular respiration, including glycolysis)

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + Energy

  • Functioning of batteries in all electronic devices
  • Photosynthesis – reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen

6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2

  • Burning of organic material and combustion of hydrocarbons in fossil fuels
  • Extraction of ores from minerals
  • Electroplating
  • Developing a photographic film
  • Bleaching and decolorization
  • Corrosion
  • Breathalyzer


Q.1. What is a spontaneous redox reaction?

Ans. In a voltaic cell, if a positive voltage is formed across the electrodes, then the redox reaction is a spontaneous reaction for reduction at the cathode and oxidation at the anode.

Q.2. What is the purpose of redox reactions in a living cell?

Ans. Redox reactions are important because they are the principal sources of energy for a living cell. Oxidation of molecules by removal of hydrogen or combination with oxygen liberates typically large quantities of energy.

Q.3. How do you know if a redox reaction is acidic or basic?

Ans. From the reaction mechanism, if H+ is present while balancing the reaction, then it is an acidic medium. If OH is present, then it is a basic medium.

Q.4. What makes balancing redox reactions different from balancing other reactions?

Ans. Balancing redox reactions is different from balancing other reactions. It is because, in a redox reaction, half-reaction is used where both the number of atoms and the amount of charge must be balanced.

Q.5. Is corrosion a redox reaction?

Ans. Yes, corrosion is a redox reaction.



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