Redox Reactions – Comprehensive Notes
Introduction to Redox Reactions
Redox reactions are chemical reactions in which oxidation and reduction occur simultaneously. Oxidation is the loss of electrons, while reduction is the gain of electrons.
These reactions are fundamental in various chemical processes, including metabolism, combustion, corrosion, and electrochemical cells.
Key Concepts and Definitions
- Oxidation: Addition of oxygen, loss of hydrogen, or loss of electrons.
- Reduction: Removal of oxygen, gain of hydrogen, or gain of electrons.
- Oxidizing Agent: Substance that gains electrons and gets reduced.
- Reducing Agent: Substance that loses electrons and gets oxidized.
- Oxidation Number: The charge assigned to an atom in a molecule or ion based on electron allocation.
Rules for Assigning Oxidation Numbers
- Oxidation state of an element in its free state is zero.
- Oxygen usually has an oxidation number of -2 (except in peroxides where it is -1).
- Hydrogen has an oxidation number of +1 when bonded to non-metals and -1 when bonded to metals.
- The sum of oxidation numbers in a neutral compound is zero; in ions, it equals the ion charge.
Types of Redox Reactions
- Combination Reactions: Two or more reactants combine to form one product.
Example: 2 Mg + O2 → 2 MgO - Decomposition Reactions: A compound breaks down into two or more products.
Example: 2 KClO3 → 2 KCl + 3 O2 - Displacement Reactions: A more reactive element displaces a less reactive element from its compound.
Example: Zn + 2 HCl → ZnCl2 + H2 - Disproportionation Reactions: An element undergoes both oxidation and reduction simultaneously.
Example: Cl2 + 2 NaOH → NaCl + NaClO + H2O
(Chlorine is oxidized to ClO– and reduced to Cl–) - Comproportionation Reactions: Two different oxidation states of the same element react to form a product with an intermediate oxidation state.
Example: Cu + CuCl2 → 2 CuCl
(Copper in 0 and +2 states forms copper in +1 state)
Rules & Methods for Balancing Chemical Equations
Balancing chemical equations ensures mass and charge conservation during chemical reactions. Different methods apply depending on the type of reaction.
General Rules for Balancing
- Atoms of each element must be equal on both sides of the equation.
- The total charge must be the same on both sides for ionic reactions.
- Start balancing with elements that appear in fewer compounds first.
- Balance oxygen and hydrogen atoms last, especially in redox reactions.
Methods of Balancing
1. Inspection Method (Trial and Error)
Used mainly for simple non-redox equations by adjusting coefficients to balance atoms.
Example: Balance N2 + H2 → NH3
Balanced as: N2 + 3 H2 → 2 NH3
2. Oxidation Number Method (For Redox Reactions)
- Assign oxidation numbers to all elements.
- Identify which elements are oxidized and reduced.
- Calculate the change in oxidation number for each element.
- Balance the increase and decrease in oxidation numbers by adjusting coefficients.
- Balance atoms and charges as needed.
Example 1: Balance Fe + CuSO4 → FeSO4 + Cu
Balanced: Fe + CuSO4 → FeSO4 + Cu
Example 2: Balance H2S + SO2 → S + H2O
Balanced: 2 H2S + SO2 → 3 S + 2 H2O
3. Ion-Electron (Half Reaction) Method
- Separate the reaction into oxidation and reduction half-reactions.
- Balance atoms other than O and H.
- Balance oxygen atoms by adding H2O.
- Balance hydrogen atoms by adding H+ (acidic medium) or OH– (basic medium).
- Balance charges by adding electrons.
- Multiply the half-reactions to equalize electrons lost and gained.
- Add half-reactions and cancel common species.
Example 1: MnO4– + Fe2+ → Mn2+ + Fe3+ (acidic medium)
Balanced: MnO4– + 5 Fe2+ + 8 H+ → Mn2+ + 5 Fe3+ + 4 H2O
Example 2: Cr2O72- + I– → Cr3+ + I2 (acidic medium)
Balanced: Cr2O72- + 14 H+ + 6 I– → 2 Cr3+ + 3 I2 + 7 H2O
Example 3: MnO4– + Cl– → MnO2 + ClO3– (basic medium)
Balanced: MnO4– + Cl– + 2 H2O → MnO2 + ClO3– + 4 OH–
Applications of Redox Reactions
- Used in electrochemical cells and batteries.
- Corrosion prevention and metallurgy.
- Titrations involving redox reactions (e.g. KMnO4 titrations).
- Industrial synthesis such as manufacture of chemicals.
Important Question-Answer
Q1. Define oxidation and reduction with examples.
Oxidation is the gain of oxygen, loss of hydrogen, or loss of electrons (e.g., Zn → Zn2+ + 2e–). Reduction is the loss of oxygen, gain of hydrogen, or gain of electrons (e.g., Cu2+ + 2e– → Cu).
Q2. How are oxidizing and reducing agents identified in a reaction?
The oxidizing agent gains electrons and gets reduced; the reducing agent loses electrons and gets oxidized.
Q3. State the rules to assign oxidation number.
Some rules include elemental state = 0, oxygen = -2 (except in peroxides), hydrogen = +1 with non-metals, and sum of oxidation states equals the overall charge.
Q4. Balance the redox reaction: MnO4– + Fe2+ → Mn2+ + Fe3+ in acidic medium.
Using ion-electron method, balanced equation: MnO4– + 5 Fe2+ + 8 H+ → Mn2+ + 5 Fe3+ + 4 H2O.
Q5. What is a redox couple?
A redox couple consists of the oxidized and reduced form of a species, such as Fe3+/Fe2+ or MnO4–/Mn2+.
Q6. What is disproportionation? Give an example.
Disproportionation is a redox reaction where the same element undergoes both oxidation and reduction simultaneously. For example, chlorine gas reacts with sodium hydroxide: Cl2 + 2 NaOH → NaCl + NaClO + H2O
Q7. What is comproportionation? Give an example.
Comproportionation is a redox reaction where two different oxidation states of the same element combine to form an intermediate oxidation state. For example, Cu + CuCl2 → 2 CuCl
Summary Table of Oxidation and Reduction
| Process | Definition | Electron Transfer | Example |
|---|---|---|---|
| Oxidation | Loss of electrons | Electrons given away | Zn → Zn2+ + 2e– |
| Reduction | Gain of electrons | Electrons accepted | Cu2+ + 2e– → Cu |