Redox Reactions
'Oxidation is gain of oxygen, loss of hydrogen, loss of electrons. Reduction is the opposite. But redox is about electron transfer.' — Chemistry Concepts
1. Chapter Overview
REDOX reactions are reactions involving ELECTRON TRANSFER. They are among the MOST COMMON and important chemical reactions — from rusting to respiration, from batteries to bleaching. This chapter covers OXIDATION and REDUCTION in classical and electronic terms, OXIDATION NUMBER (a powerful bookkeeping tool), BALANCING redox reactions, and the ELECTROCHEMICAL SERIES.
2. Classical Concept of Oxidation and Reduction
Oxidation (Classical)
- Addition of OXYGEN
- Removal of HYDROGEN
- Gain of ELECTRONEGATIVE element
Reduction (Classical)
- Addition of HYDROGEN
- Removal of OXYGEN
- Gain of ELECTROPOSITIVE element
Mnemonic
- OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)
3. Oxidation Number (Oxidation State)
- Definition: The CHARGE an atom would have if electrons were COMPLETELY transferred (a bookkeeping concept)
- Rules for Assigning ON:
- ON of an element in FREE state = 0 (H₂, O₂, Fe, S₈)
- ON of a MONOATOMIC ion = its charge (Na⁺ = +1, Cl⁻ = -1)
- ON of HYDROGEN = +1 (except metal hydrides where it's -1: NaH, CaH₂)
- ON of OXYGEN = -2 (except peroxides -1: H₂O₂; superoxides -½: KO₂; OF₂ = +2)
- Sum of ON in a NEUTRAL compound = 0
- Sum of ON in a POLYATOMIC ion = charge on the ion
- Halogens: -1 in most compounds (except with O or higher halogens)
Worked Problem
Q: Find ON of Cr in Cr₂O₇²⁻ and S in H₂SO₄. A: Cr₂O₇²⁻: 2x + 7(-2) = -2 → 2x = +12 → x = +6. H₂SO₄: 2(+1) + x + 4(-2) = 0 → 2 + x — 8 = 0 → x = +6.
4. Types of Redox Reactions
| Type | Description | Example |
|---|---|---|
| Combination | Two substances combine | 2Mg + O₂ → 2MgO |
| Decomposition | One substance breaks down | 2H₂O → 2H₂ + O₂ |
| Displacement | One element replaces another | Zn + CuSO₄ → ZnSO₄ + Cu |
| Disproportionation | Same element oxidised AND reduced | 2H₂O₂ → 2H₂O + O₂ |
| Combustion | Rapid oxidation producing heat/light | CH₄ + 2O₂ → CO₂ + 2H₂O |
Disproportionation Reactions
- The ELEMENT must have at least THREE oxidation states
- One atom of the element is OXIDISED, another REDUCED
- Example: 2H₂O₂ → 2H₂O + O₂ (O in H₂O₂: -1 → goes to -2 and 0)
5. Balancing Redox Equations
Method 1: Oxidation Number Method
- Assign ON to ALL atoms
- Identify which atoms change ON
- Calculate CHANGE in ON (increase = oxidation; decrease = reduction)
- Make total increase = total decrease (balance electrons)
- Balance O with H₂O and H with H⁺ (in acidic medium)
- For basic medium: after balancing in acidic, add OH⁻ to NEUTRALISE H⁺
Method 2: Half-Reaction Method (Ion-Electron Method)
- Write OXIDATION and REDUCTION half-reactions separately
- Balance ATOMS other than H and O
- Balance O with H₂O, H with H⁺ (acidic) or OH⁻ (basic)
- Balance CHARGE by adding e⁻
- MULTIPLY half-reactions to equalise electrons lost = gained
- ADD half-reactions
Worked Problem (Acidic Medium)
Q: Balance Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ in acidic medium. A:
- Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (×6)
- Reduction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
- Adding: 6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O
6. Electrochemical Series
| Half-Reaction | E° (V) |
|---|---|
| Li⁺ + e⁻ → Li | -3.04 |
| K⁺ + e⁻ → K | -2.93 |
| Ca²⁺ + 2e⁻ → Ca | -2.87 |
| Na⁺ + e⁻ → Na | -2.71 |
| Mg²⁺ + 2e⁻ → Mg | -2.36 |
| Al³⁺ + 3e⁻ → Al | -1.66 |
| Zn²⁺ + 2e⁻ → Zn | -0.76 |
| Fe²⁺ + 2e⁻ → Fe | -0.44 |
| Ni²⁺ + 2e⁻ → Ni | -0.25 |
| 2H⁺ + 2e⁻ → H₂ | 0.00 (Reference) |
| Cu²⁺ + 2e⁻ → Cu | +0.34 |
| Ag⁺ + e⁻ → Ag | +0.80 |
| Au³⁺ + 3e⁻ → Au | +1.40 |
| F₂ + 2e⁻ → 2F⁻ | +2.87 |
Uses of Electrochemical Series
- Predicting spontaneity: More NEGATIVE E° → better reducing agent (tendency to get oxidised)
- Predicting displacement reactions: A metal with MORE negative E° reduces the ion of a metal with LESS negative E°
- Calculating cell EMF: E°_cell = E°_cathode — E°_anode
Worked Problem
Q: Zn + Cu²⁺ → Zn²⁺ + Cu. Calculate E°_cell. (E°_Zn = -0.76 V, E°_Cu = +0.34 V) A: Zn is oxidised (anode), Cu²⁺ reduced (cathode). E°_cell = E°_Cu — E°_Zn = 0.34 — (-0.76) = 1.10 V. Positive → SPONTANEOUS.
7. Common Mistakes
- ON is NOT the actual charge (except for monatomic ions): It is a BOOKKEEPING convention
- H is +1 EXCEPT in metal hydrides: Many students forget NaH has H = -1
- O is -2 EXCEPT in peroxides and superoxides: H₂O₂ has O = -1, not -2
- Disproportionation requires the element to have an INTERMEDIATE oxidation state: Not all reactions of an element are disproportionation
- E°_cell positive means spontaneous, but real rate may be slow: Thermodynamic favourability ≠ fast kinetics
8. CBSE Exam Focus
- Oxidation number calculations (3-mark)
- Identifying oxidising and reducing agents (1-mark)
- Balancing redox equations — both methods (5-mark)
- Electrochemical series — predicting reactions (3-mark)
- Disproportionation reactions — examples (3-mark)
9. Key Formulas
- E°_cell = E°_cathode — E°_anode (standard cell potential)
- ΔG° = -nFE°_cell (relation between free energy and cell potential)
- F = 96500 C/mol (Faraday constant)
10. Self-Test (5+ Q&A)
Q1: Assign ON to Mn in KMnO₄ and Cr in Cr₂O₃. A: KMnO₄: K = +1, each O = -2 → +1 + x + 4(-2) = 0 → x = +7. Cr₂O₃: 2x + 3(-2) = 0 → 2x = +6 → x = +3.
Q2: Balance in basic medium: MnO₄⁻ + I⁻ → MnO₂ + I₂. A: After balancing in acidic then converting to basic: 2MnO₄⁻ + 6I⁻ + 4H₂O → 2MnO₂ + 3I₂ + 8OH⁻
Q3: Which is a stronger reducing agent — Zn or Cu? Why? A: Zn (E° = -0.76 V) is stronger than Cu (E° = +0.34 V). MORE negative E° means greater tendency to LOSE electrons (be oxidised).
Q4: What is a disproportionation reaction? Give an example. A: Reaction where the SAME element is BOTH oxidised and reduced. Example: 2H₂O₂ → 2H₂O + O₂ (O goes from -1 to -2 and 0).
Q5: Calculate E°_cell for Fe²⁺ + Zn → Fe + Zn²⁺. (E°_Fe = -0.44 V, E°_Zn = -0.76 V) A: Zn is oxidised (anode), Fe²⁺ is reduced (cathode). E°_cell = E°_Fe — E°_Zn = -0.44 — (-0.76) = +0.32 V → SPONTANEOUS.
11. Conclusion
Redox reactions are FUNDAMENTAL to chemistry and life. The oxidation number system provides a POWERFUL bookkeeping tool for tracking electron transfer. Balancing redox equations is an ESSENTIAL skill. The electrochemical series is a ROADMAP of reduction potentials — predicting spontaneity, displacement reactions, and cell voltages. These concepts prepare you for the in-depth study of electrochemistry in Class 12 and are VITAL for understanding biological redox processes (respiration, photosynthesis).
