Redox

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Redox — Quick Facts

• Oxidation = loss of electrons (OIL: Oxidation Is Loss) → LE0 (Lose Electrons = Oxidation)
• Reduction = gain of electrons (RIG: Reduction Is Gain) → GER (Gain Electrons = Reduction)
• Oxidizing agent gets reduced (gains electrons); Reducing agent gets oxidized (loses electrons)
• Oxidation number rules: Free elements = 0; O = -2 (except peroxides); H = +1 (except metal hydrides); Sum of ONs in neutral compound = 0

⚡ Exam tip: NEET questions on redox often test oxidation number calculation first. Memorize the common oxidation states: Fe shows +2/+3; Mn shows +2/+4/+7; Cr shows +2/+3/+6; S shows -2/+4/+6.

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🹡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Redox — Study Guide

Overview: Redox (oxidation-reduction) reactions are central to both inorganic chemistry and electrochemistry. They form the basis for electrochemical cells, corrosion, and many industrial processes. In NEET, this topic is tested every year — primarily through oxidation number calculations, balancing redox equations, and electrochemical cell potentials. The concepts also connect directly to equivalent weight calculations and the Nernst equation.

Key concepts:

Oxidation Number Rules (Priority Order for NEET):

1. Free element (uncombined atom) = 0
2. Monatomic ion = charge on ion
3. Oxygen = -2 in most compounds; -1 in peroxides (H₂O₂); -½ in superoxides (KO₂)
4. Hydrogen = +1 in most compounds; -1 in metal hydrides (NaH, CaH₂)
5. Sum of all oxidation numbers in a neutral compound = 0
6. Sum of all oxidation numbers in a polyatomic ion = charge on ion

Common Redox Reactions in NEET:

• Combination: A + B → AB (e.g., 2Na + Cl₂ → 2NaCl)
• Decomposition: AB → A + B (e.g., 2H₂O → 2H₂ + O₂)
• Displacement: A + BC → AC + B (e.g., Zn + CuSO₄ → ZnSO₄ + Cu)
• Disproportionation: Same element is both oxidized and reduced (e.g., 2H₂O₂ → 2H₂O + O₂)

Balancing Redox Equations (Ion-Electron Method):

1. Write the unbalanced ionic equation
2. Separate into oxidation and reduction half-reactions
3. Balance atoms other than O and H in each half
4. Balance O by adding H₂O; balance H by adding H⁺
5. Balance charge by adding electrons
6. Multiply half-reactions so electrons are equal
7. Add and cancel common species

Standard Electrode Potentials:

• E°cell = E°cathode - E°anode
• E°cell > 0 → Spontaneous (galvanic cell)
• E°cell < 0 → Non-spontaneous (electrolytic cell)
• For a galvanic cell: ΔG° = -nFE°cell (where n = number of electrons, F = 96500 C/mol)

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Redox — Comprehensive Notes

Full Coverage: Redox chemistry is foundational for understanding electrochemical cells, corrosion, and titration methods (permanganometry, dichrometry). NEET frequently asks about the relationship between E°cell and ΔG°, and about the Nernst equation for non-standard conditions.

Electrochemical Cell Conventions:

• Anode is where oxidation occurs (left side in cell notation)
• Cathode is where reduction occurs (right side in cell notation)
• Salt bridge connects the two half-cells — maintains electrical neutrality
• Cell notation: Anode | Anode solution || Cathode solution | Cathode
• Example: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s) → E°cell = +1.10 V

Nernst Equation:

• E = E° - (RT/nF) ln Q
• At 298 K: E = E° - (0.0591/n) log Q
• This allows calculation of cell potential under non-standard conditions

Equivalent Weight of Oxidizing/Reducing Agents:

• Eq. wt. = Molecular weight / n-factor (change in oxidation number)
• For KMnO₄ in acidic medium: n = 5 (Mn goes from +7 to +2)
• For KMnO₄ in alkaline medium: n = 3 (Mn goes from +7 to +4)
• For K₂Cr₂O₇ in acidic medium: n = 6 (Cr goes from +6 to +3)

Common NEET Mistakes to Avoid:

• Forgetting to multiply half-reaction electron counts when combining
• Confusing oxidation number with ionic charge — oxidation number is a formal concept, not actual charge
• Using the wrong n-factor for permanganate (5 in acid, 3 in base, 1 in neutral)
• Mixing up anode and cathode polarities — remember: oxidation at anode, reduction at cathode

Important Formulas:

• E°cell = E°cathode - E°anode
• ΔG° = -nFE°cell
• log K = nE°cell / 0.0591 (at 298 K)
• E°cell > 0.4 V → strong oxidizing agent

Related Topics: pc-005 (Electrochemistry — cells and potentials), pc-012 (Solutions — molarity and equivalents), pc-006 (Ionic Equilibrium — pH in redox context)

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