Redox reactions involve the transfer of electrons between substances, leading to changes in oxidation states. Oxidation is electron loss; reduction is electron gain. They are essential in energy production and metabolism.
Redox reactions can be balanced by using various methods like Ion electron method, oxidation number method.
Types of Redox Reactions
Combination Reactions:
Two elements combine to form a compound.
Example: 2 Mg
(s)
+ O
2(g)
→ 2 MgO
(s)
Decomposition Reactions:
A compound breaks down into elements or simpler compounds.
Example: 2H
2
O
(l)
→ 2 H
2(g)
+ O
2(g)
Single Displacement:
Displacement Reactions: An element displaces another element from a compound.
Cu
(s)
+ 2AgNO
3(aq)
→ Cu(NO
3
)
2(aq)
+ 2 Ag
(s)
Double Displacement:
It is a type of chemical reaction in which two compounds exchange ions with each other to form two new compounds.
Pb(NO
3
)
2(aq)
+ 2 KI
(aq)
→ 2 KNO
3(aq)
+ PbI
2(s)
Combustion Reactions:
A substance reacts with oxygen, usually producing energy in the form of heat and light.
Example: CH
4(g)
+ 2O
2(g)
→ CO
2(g)
+ 2H
2
O
(g)
Redox Titrations
Analytical method to determine concentration by adding a known reagent until reaction completion.
Example: Titration of iron (II) solution with potassium permanganate solution.
Oxidation:
Oxidation is the process of losing electrons. It may also involve the increase in oxidation state or the addition of oxygen or removal of hydrogen.
Example: Na → Na
+
+ e
–
Sodium loses an electron, getting oxidized to sodium ion.
Reduction:
Reduction is the process of gaining electrons. It can also involve the decrease in oxidation state or the addition of hydrogen or removal of oxygen.
Example: Cl
2
+2e
−
→ 2Cl
–
Chlorine gains electrons, getting reduced to chloride ions.
Download PDF Redox Reactions Formula
Oxidizing Agent (Oxidant)
A substance that facilitates the oxidation of another substance, and in doing so, gets reduced itself.
Example: Cl
2
- In the reaction
2Na + Cl
2
→ 2NaCl
, chlorine acts as an oxidizing agent, causing sodium to lose electrons, while itself getting reduced.
Also Check –
Malic Acid Formula
Reducing Agent (Reductant)
A substance that facilitates the reduction of another substance, and in doing so, gets oxidized itself. Example: Na In the aforementioned reaction, sodium acts as a reducing agent, causing chlorine to gain electrons, while itself getting oxidized.
Also Read:
Molar Volume Formula
Measurement of Reduction Rotential
The reduction potential (E∘) of a half-cell measures a chemical species' ability to acquire electrons and be reduced in an electrochemical system. It is expressed in volts (V) and is determined by the species' affinity for electrons. A higher reduction potential indicates stronger oxidation, while a lower potential indicates weaker oxidation. The standard hydrogen electrode (SHE) serves as a reference for measuring reduction potentials in other half-reactions.
Example: The reduction potential for the copper half-reaction:
Cu
2+
(aq)
+2e
−
→Cu
(s)
is E
∘
= +0.34V.
This positive value indicates that the copper ion has a tendency to gain electrons and be reduced, especially when compared to the hydrogen electrode.
Also Check –
Periodic Acid formula
Balancing Redox Reactions
Ion-Electron Method (Half-Reaction Method):
Balancing redox reactions by separating and balancing the oxidation and reduction half-reactions.
Example: Balance the following in acidic solution:
MnO
4
−
+ C
2
O
4
2−
→ Mn
2+
+ CO
2
Split into half-reactions:
MnO
4
−
→ Mn
2+
(Reduction)
C
2
O
4
2−
→ CO
2
(Oxidation)
Balance atoms other than O and H:
MnO
4−
→ Mn
2+
is already balanced for Mn
C
2
O
4
2−
→ 2CO
2
to balance the carbon atoms.
Balance the O atoms by adding H
2
O:
MnO
4
−
→ Mn
2+
+ 4H
2
O
Balance the H atoms by adding H
+
:
MnO
4
−
+ 8H
+
→ Mn
2+
+ 4H
2
O
Balance charge by adding electrons:
MnO
4
−
+ 8 H
+
+ 5e
-
→ Mn
2+
+ 4H
2
O
C
2
O
4
2−
→ 2CO
2
+2e
–
Equalize the electron transfer (multiply the oxidation half-reaction by 5):
C
2
O
4
2−
→ 2CO
2
+10e
–
Add the half-reactions together and simplify:
MnO
4
−
+ 5C
2
O
4
2−
+ 8 H
+
→ Mn
2+
+ 10CO
2
+ 4H
2
O
Oxidation Number Method:
Balance redox reactions by adjusting coefficients based on changes in oxidation numbers.
Example: Balance:
H
2
O + Cl
2
→ HClO + HCl
Assign oxidation numbers:
H in H
2
O: +1,
O: -2,
Cl in Cl 2 : 0,
Cl in HClO: +1,
Cl in HCl: -1
Identify oxidized and reduced atoms: Cl
2
is reduced (0 to -1 and +1).
H
2
O is oxidized (+1 to 0).
Balance atoms that change oxidation state:
To balance Cl, add a coefficient of 2 in front of HClO and HCl:
H
2
O + Cl
2
→ 2HClO + 2HCl
Balance remaining atoms: Add a coefficient of 3 in front of H
2
O:
3H
2
O + Cl
2
→ 2HClO + 2HCl
