Introduction
Redox reactions (reduction-oxidation reactions) are one of the most fundamental concepts in chemistry. These reactions involve the transfer of electrons between chemical species and play a crucial role in both laboratory chemistry and real-life processes such as respiration, corrosion, combustion, and electrochemistry.
Understanding redox reactions is essential for mastering advanced topics like electrochemistry, metallurgy, and organic chemistry.
Concept of Oxidation and Reduction
Classical Concept
Oxidation was initially defined as the addition of oxygen or removal of hydrogen.
Examples:
2Mg + O2 → 2MgO
Here, magnesium is oxidized.
H2 + Cl2 → 2HCl
Hydrogen is oxidized as it combines with chlorine.
Reduction was defined as removal of oxygen or addition of hydrogen.
CuO + H2 → Cu + H2O
Copper oxide is reduced.
Electronic Concept (Modern Definition)
Oxidation = Loss of electrons
Reduction = Gain of electrons
Example:
Zn → Zn2+ + 2e− (Oxidation)
Cu2+ + 2e− → Cu (Reduction)
Oxidation Number (Oxidation State)
Oxidation number is the apparent charge assigned to an atom in a compound.
Rules for Assigning Oxidation Number
- Oxidation number of free element = 0
Example:
O2, H2, Na → 0
- For monatomic ions:
Na+ = +1 Cl− = −1
- Oxygen:
O = −2 (generally)
Exception:
Peroxides: O = −1
- Hydrogen:
H = +1 (with non-metals) H = −1 (with metals)
- Alkali metals:
Group 1 = +1
- Alkaline earth metals:
Group 2 = +2
- Sum rule:
Sum of oxidation numbers = charge on molecule/ion
Calculation of Oxidation Number
Example:
H2SO4
Let oxidation state of S = x
2(+1) + x + 4(−2) = 0
2 + x − 8 = 0
x = +6
Redox Reactions Types
1. Combination Reaction
A + B → AB
Example:
2Mg + O2 → 2MgO
2. Decomposition Reaction
AB → A + B
Example:
2HgO → 2Hg + O2
3. Displacement Reaction
A + BC → AC + B
Example:
Zn + CuSO4 → ZnSO4 + Cu
4. Disproportionation Reaction
Same element undergoes oxidation and reduction
Example:
2H2O2 → 2H2O + O2
Balancing Redox Reactions
Oxidation Number Method
Steps:
- Assign oxidation numbers
- Identify oxidation and reduction
- Equalize electron transfer
- Balance atoms and charges
Half Reaction Method
Steps:
- Split into oxidation and reduction
- Balance atoms
- Balance charges using electrons
- Combine reactions
Reducing and Oxidizing Agents
Oxidizing agent = substance that gets reduced
Reducing agent = substance that gets oxidized
Example:
Zn + Cu2+ → Zn2+ + Cu
Zn = reducing agent Cu2+ = oxidizing agent
Electrochemical Concept
Redox reactions involve electron flow.
Electric current is produced due to electron movement.
Applications of Redox Reactions
1. Respiration
Glucose oxidation:
C6H12O6 + O2 → CO2 + H2O
2. Combustion
CH4 + O2 → CO2 + H2O
3. Corrosion
Rusting of iron:
Fe → Fe2+ + 2e−
4. Photosynthesis
6CO2 + 6H2O → C6H12O6 + 6O2
Equivalent Concept
Equivalent mass = Molar mass / n-factor
n-Factor
n-factor = number of electrons lost or gained
Redox Titrations
Used to determine concentration.
Example:
KMnO4 titration
Important Equations Summary
Oxidation:
Loss of electrons
Reduction:
Gain of electrons
Oxidation number sum:
Sum = total charge
n-factor:
n = electrons transferred
Equivalent mass:
Equivalent mass = molar mass / n
Common Mistakes
- Wrong oxidation number assignment
- Not balancing charge
- Confusing oxidizing and reducing agents
- Skipping electron balance
Exam Tips
- Practice balancing reactions
- Learn oxidation number rules
- Focus on numericals
- Revise examples regularly
Conclusion
Redox reactions are essential for understanding chemical processes involving electron transfer. They are widely applicable in real-life processes and advanced chemistry topics.
FAQs
Q1. What is oxidation?
Loss of electrons
Q2. What is reduction?
Gain of electrons
Q3. What is oxidizing agent?
Substance that gets reduced