How to Identify a Redox Reaction

How to identify a redox reaction? -A redox reaction can be identified by looking for changes in the oxidation states of the atoms involved in the reaction.

In the realm of chemistry, certain reactions involve a shift in the way atoms share electrons. These reactions are known as redox reactions. A redox reaction is a type of chemical reaction where there is an exchange of electrons between different substances. In these reactions, one substance loses electrons (oxidation) while another gains electrons (reduction). This article focuses on how to identify redox reactions by understanding the changes in electron distribution, providing a straightforward approach to recognizing them in various chemical processes.

Key Areas Covered

1. What is a Redox Reaction
      – Definition, Features
2. How to Identify a Redox Reaction
      – Finding the Oxidation Number
3. Practice Identifying Redox Reactions
      – Examples of Reactions
4. Types of Redox Reaction
      – Various Types 

Key Terms

Redox Reaction, Oxidation, Half-Reaction

What is a Redox Reaction

Oxidation-reduction or redox reactions involve an electron transfer process. A redox reaction has two half-reactions, namely, an oxidation reaction and a reduction reaction. Oxidation reaction involves the loss of electrons, while reduction reaction involves the acceptance of electrons. Therefore, a redox reaction contains two species; the oxidizing agent undergoes the oxidization half-reaction, and the reducing agent undergoes the reducing half-reaction. The extent of reduction in a redox reaction is equal to the extent of oxidization; that means the number of electrons lost from the oxidizing agent equals the number of electrons accepted by the reducing agent. It is a balanced process in terms of electron exchange.

How to Identify a Redox Reaction

How to Identify a Redox Reaction

Find the Oxidation Number

To identify a redox reaction, first, we need to know the oxidation status of each element in the reaction. We use the following rules to assign oxidation numbers.

  • The free elements, which are not combined with others, have the oxidation number zero. Thus, atoms in H2, Br2, Na, Be, Ca, K, O2 and P4 have the same oxidation number zero. 
  • For ions that are composed of only one atom (monoatomic ions), the oxidation number equals the charge on the ion. For example:

Na+, Li+ and K+ have the oxidation number +1.
F, I, Cl and Br have the oxidation number -1.
Ba2+, Ca2+, Fe2+ and Ni2+ have the oxidation number +2.
O2- and S2- have the oxidation number -2.
Al3+ and Fe3+ have the oxidation number +3.

  • The most common oxidation number of oxygen is -2 (O2-: MgO, H2O), but in hydrogen peroxide, it is -1 (O22-: H2O2).
  • The most common oxidation number of hydrogen is +1. However, when it is bonded to metals in group I and group II, the oxidation number is -1 (LiH, NaH, CaH2).
    Fluorine (F) shows only -1 oxidation status in all its compounds; other halogens (Cl, Br and I) have both negative and positive oxidation numbers.
  • In a neutral molecule, the sum of all oxidation numbers equals zero.
  • In a polyatomic ion, the sum of all oxidation numbers equals the charge on the ion.
  • Oxidation numbers need not be only integers.

Example: Superoxide ion (O22-) – Oxygen has the -1/2 oxidation status.

Identify the Oxidation Reaction and Reduction Reaction

Consider the following reaction.

2Ca + O2(g) —> 2CaO(s)

Step 1: Determine the oxidizing agent and the reducing agent. For this, we need to identify their oxidation numbers.

2Ca + O2(g) —> 2CaO(s)
0           0                (+2) (-2)

Both of the reactants have the oxidation number zero. Calcium increases its oxidation state from (0) —> (+2). Therefore, it is the oxidizing agent. Conversely, in oxygen, the oxidation state decreases from (0) —> (-2). Therefore, Oxygen is the reducing agent.

Step 2: Write half-reactions for the oxidation and the reduction. We use electrons to balance the charges on both sides.

Oxidation: Ca(s)     —>   Ca2+ + 2e ——(1)
Reduction: O2 + 4e —>  2O2-         ——(2)

Step 3: Obtaining the redox reaction. By adding (1) and (2), we can obtain the redox reaction. Electrons in the half-reactions should not appear in the balanced redox reaction. For this, we need to multiply reaction (1) by 2 and then add it with reaction (2).

(1)*2 + (2):
2Ca(s)                              —> 2Ca2+ + 4e ——(1)
O2 + 4e                            —> 2O2-             ——(2)
2Ca + O2(g)                   —> 2CaO(s)

Practice Identifying Redox Reactions

Example: Consider the following reactions. Which one resembles a redox reaction?

Zn(s) + CuSO4(aq)     —> ZnSO4(aq) + Cu(s)

HCl(aq) + NaOH(aq) —> NaCl(aq) + H2O(l)

In a redox reaction, oxidation numbers change in reactants and products. There should be an oxidizing species and a reducing species. If the oxidation numbers of elements in the products do not change, it cannot be considered as a redox reaction.

Zn(s) + CuSO4(aq) —>  ZnSO4(aq)  +  Cu(s)
Zn(0)     Cu (+2)                      Zn (+2)     Cu (0)
                  S (+6)                       S (+6)
                  O (-2)                       O (-2)

This is a redox reaction because zinc is the oxidizing agent (0 —> (+2) and Copper is the reducing agent (+2) —> (0).

HCl(aq)         +   NaOH(aq)                  —>   NaCl(aq)         +   H2O(l)
H(+1),Cl (-1)      Na(+1), O(-2), H(+1)           Na(+1) , Cl (-1)      H(+1), O(-2)

This is not a redox reaction because the reactants and the products have the same oxidation numbers. H (+1), Cl (-1), Na(+1) and O(-2)

Types of Redox Reactions

There are four different types of redox reactions: combination reactions, decomposition reactions, displacement reactions and disproportionation reactions.

Combination Reactions

Combination reactions are the reactions in which two or more substances combine to form a single product.
A          + B             —> C
S(s)     + O2(g)     —> SO2(g)
S(0)        O(0)              S(+4),O(-2)

3 Mg(s) + N2(g) —> Mg3 N2(s)
Mg(0)       N(0)           Mg(+2), N(-3)

Decomposition Reactions

In decomposition reactions, a compound breaks down into two or more components. It is the opposite of combination reactions.

C                         —>   A              +           B
2HgO (s)           —>    2Hg(l)     +       O2(g)
Hg(+2), O(-2)              Hg(0)              O(0)


2 NaH (s)         —->  2 Na (s) + H2 (g)
Na(+1), H(-1)               Na(0)      H(0)

2 KClO3(s)      —>  2KCl(s) + 3O2(g)

Displacement Reactions

In a displacement reaction, an ion or atom in a compound is replaced by an ion or an atom of another compound. Displacement reactions have a wide range of applications in industry.

A + BC —> AC + B

Hydrogen displacement:

All alkali metals and some alkaline metals (Ca, Sr and Ba) are replaced by hydrogen from cold water.

2Na( s) + 2H2O (l) —> 2NaOH (aq) + H2(g)
Ca( s) + 2H2O (l)   —> Ca(OH)2 (aq) + H2(g)

Metal displacement:

Some metals in the elemental state can displace a metal in a compound. For example, Zinc replaces Copper ions, and Copper can replace Silver ions. The displacement reaction depends on the place activity series (or electrochemical series).

Zn(s) + CuSO4(aq) —> Cu(s) + ZnSO4(aq)

Halogen displacement:

Activity series for halogen displacement reactions: F2 > Cl2 > Br2 > I2. As we go down the halogen series, the power of oxidizing ability decreases.

Cl2(g) +  2KBr (aq)  —> 2KCl (aq)  + Br2(l)
Cl2(g) +  2KI (aq)    —> 2KCl (aq)  +  I2(s)
Br2(l) +  2I (aq)      —> 2Br(aq)    +  I2(s)

Disproportionation Reactions

This is a special type of the redox reaction. An element in one oxidation state is simultaneously oxidized and reduced. In a disproportionation reaction, one reactant should always contain an element that can have at least three oxidation states.

2H2O2(aq) —>  2H2O (l) + O2(g)

Here the oxidation number in the reactant is (-1); it increases to zero in O2 and decreases to (-2) in H2O. The oxidation number in Hydrogen does not change in the reaction.


Redox reactions are considered electron transfer reactions. In a redox reaction, one element is oxidizing and releasing electrons, and one element is reducing by gaining the released electrons. The extent of oxidation is equal to the extent of reduction in terms of electrons exchanging in the reaction. There are two half-reactions in a redox reaction; they are called oxidation half-reactions and reduction half-reactions. There is an increase in oxidation number in oxidation; similarly, the oxidation number decreases in the reduction. Therefore, the best way to identify a redox reaction is by looking at its oxidation state. 

Image Courtesy:

1. “Redox orga” By Stéphane Mons – Own work (CC BY 2.5) via Commosn Wikimedia

About the Author: Hasa

Hasanthi is a seasoned content writer and editor with over 8 years of experience. Armed with a BA degree in English and a knack for digital marketing, she explores her passions for literature, history, culture, and food through her engaging and informative writing.

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