← Chemistry — Std 11
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Redox Reactions

Ch. 6Std 11

Easy Overview

Every time you charge a phone, start a fire, or watch iron rust — that's redox. OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons. Always happen together. The oxidizing agent gets reduced; the reducing agent gets oxidized. Oxidation numbers track electrons. Changes reveal which atoms are oxidized (↑) and reduced (↓). Balancing redox uses the half-reaction method: split into oxidation and reduction halves, balance atoms (H₂O, H⁺/OH⁻) and charge (e⁻), then combine. Electrochemical cells convert chemical energy to electricity (Daniel cell: Zn/Cu, E° = 1.10 V). The Nernst equation relates potential to concentration. Electrolysis drives non-spontaneous reactions. Faraday's laws: m = (E×I×t)/96,485.

Oxidation and Reduction — The Classic Dance

Oxidation: loss of e⁻ (OIL). Reduction: gain of e⁻ (RIG). Always together. Older: oxidation = gain O/lose H; reduction = lose O/gain H. 2Mg + O₂ → 2MgO: Mg oxidized (0→+2), O₂ reduced (0→-2). Mg = reducing agent (causes reduction by giving e⁻). O₂ = oxidizing agent (causes oxidation by accepting e⁻).

Oxidation Number — The Electron Scorecard

Rules: free element = 0. Monatomic ion = charge. O = -2 (except peroxides -1, superoxides -½, OF₂ +2). H = +1 (except metal hydrides -1). F = -1 always. Sum = 0 for compound, = charge for ion. H₂SO₄: H=+1, O=-2 → S=+6. KMnO₄: K=+1, O=-2 → Mn=+7. Increase = oxidation, decrease = reduction.

Oxidizing and Reducing Agents

Oxidizing agent: accepts e⁻, gets reduced (ON ↓). Reducing agent: donates e⁻, gets oxidized (ON ↑). Zn + CuSO₄ → ZnSO₄ + Cu: Zn reducing agent (0→+2), CuSO₄ oxidizing agent (+2→0). Common oxidants: O₂, Cl₂, HNO₃, KMnO₄, K₂Cr₂O₇, H₂O₂. Common reductants: metals (Na, Zn, Fe), H₂, C, CO, H₂S, KI. Strength: E° value (positive = stronger oxidant).

Balancing Redox — Half-Reaction Method (Acidic)

Steps: (1) Write half-reactions. (2) Balance other atoms. (3) Add H₂O for O. (4) Add H⁺ for H. (5) Add e⁻ for charge. (6) Multiply halves so e⁻ cancel. (7) Add and simplify. Cr₂O₇²⁻ + Fe²⁺ → Cr³⁺ + Fe³⁺ (acidic). Reduction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (×6). Combined: Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 7H₂O + 6Fe³⁺.

Balancing Redox — Basic Medium

Balance in acidic, then add OH⁻ to both sides to neutralize H⁺. For Cr(OH)₃ + IO₃⁻ → CrO₄²⁻ + I⁻ (basic): after acidic balance, add 10OH⁻ → combine H⁺+OH⁻ as H₂O, cancel. Result: 2Cr(OH)₃ + IO₃⁻ + 10OH⁻ → 2CrO₄²⁻ + I⁻ + 6H₂O.

Disproportionation Reactions

Same element simultaneously oxidized and reduced. Need ≥3 oxidation states; intermediate state unstable. Examples: 2H₂O₂ → 2H₂O + O₂ (O: -1→-2 and 0). 3Cl₂ + 6OH⁻ → 5Cl⁻ + ClO₃⁻ + 3H₂O (Cl: 0→-1 and +5). 2Cu⁺ → Cu + Cu²⁺ (Cu⁺: +1→0 and +2).

Electrochemical Cells — Batteries and Beyond

Galvanic cell: chemical → electrical energy. Anode (oxidation, -), cathode (reduction, +). Daniel cell: Zn|Zn²⁺||Cu²⁺|Cu. Zn → Zn²⁺ + 2e⁻ (anode, -0.76 V). Cu²⁺ + 2e⁻ → Cu (cathode, +0.34 V). E°_cell = E°_cathode - E°_anode = 0.34-(-0.76)=+1.10 V. Salt bridge maintains charge balance. Electrons flow anode→cathode.

Standard Electrode Potential and Nernst Equation

E° measured vs SHE (0 V) at 1 M, 1 atm, 298 K. More positive E° = stronger oxidant. Nernst: E = E° - (RT/nF)lnQ. At 298 K: E = E° - (0.0591/n)log₁₀Q. Q = [products]/[reactants] (exclude solids/liquids). Battery voltage drops as Q increases (discharge). At equilibrium: E=0, Q=K.

Electrochemical Series — The Potentials Table

List of E° values. F₂ (+2.87 V) = strongest oxidant. Li (-3.04 V) = strongest reductant. Any species oxidizes what's above it. Metals with E°<0 displace H₂ from acids (Mg, Zn, Fe). Cu, Ag, Au (E°>0) are 'noble' — don't react with non-oxidizing acids. E°_cell = E°_cathode - E°_anode; positive = spontaneous.

Types of Electrodes

Metal-metal ion: Zn|Zn²⁺, Cu|Cu²⁺. Gas: Pt|H₂|H⁺ (SHE, E°=0). Metal-insoluble salt: Ag|AgCl|Cl⁻, Hg|Hg₂Cl₂|Cl⁻ (calomel). Redox: Pt|Fe³⁺,Fe²⁺. Reference electrodes (SHE, calomel, Ag/AgCl) provide stable potentials for measurement.

Electrolysis — Driving Non-Spontaneous Reactions

External power drives non-spontaneous redox. Anode (+): oxidation. Cathode (-): reduction. Water: 2H₂O → 2H₂ (cathode) + O₂ (anode). Molten NaCl: 2NaCl → 2Na (cathode) + Cl₂ (anode). Aqueous NaCl (chlor-alkali): H₂ at cathode, Cl₂ at anode, NaOH in solution. Faraday's laws govern product amounts.

Faraday's Laws of Electrolysis

First: m ∝ Q = I×t. Second: m ∝ E (equivalent mass). Combined: m = E×I×t/96,485. One faraday (96,485 C) deposits 1 gram equivalent. Example: 0.5 A for 1 hour deposits Cu: m = (31.75×0.5×3600)/96485 = 0.592 g (E(Cu)=63.5/2=31.75 for Cu²⁺).

Applications of Electrochemical Cells

Primary (non-rechargeable): dry cell (Zn/MnO₂, 1.5 V), mercury battery (Zn/HgO). Secondary (rechargeable): lead-acid (Pb/PbO₂ in H₂SO₄, 2 V/cell, car battery), Li-ion (LiCoO₂/graphite, high energy density). Fuel cells: continuous fuel (H₂+½O₂→H₂O, produces electricity, only byproduct water). Corrosion prevention: galvanizing (Zn coating), cathodic protection (sacrificial Mg/Zn).

Corrosion — Redox Gone Wrong

Electrochemical degradation of metals. Fe anode: Fe → Fe²⁺ + 2e⁻. O₂ cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻. Fe²⁺ + 2OH⁻ → Fe(OH)₂ → Fe₂O₃·xH₂O (rust). Needs O₂, water, electrolyte. Accelerated by: low pH, salts, contact with less active metals. Prevention: paint, galvanizing, cathodic protection, alloying (stainless steel — Cr forms protective Cr₂O₃).

Batteries — Portable Redox Power

Primary (non-rechargeable): dry cell (Zn anode, MnO₂/C cathode, NH₄Cl paste, 1.5 V). Secondary: lead-acid (Pb + PbO₂ in H₂SO₄; discharge → both become PbSO₄; recharge reverses; 2 V/cell, 6 cells=12 V). Ni-Cd (NiOOH/Cd, 1.2 V, memory effect). Li-ion (LiCoO₂/graphite, Li⁺ shuttles, high energy density, no memory effect).

Redox in Biological Systems

Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O (glucose oxidized, O₂ reduced). Electron transport chain: series of redox reactions producing ATP. Photosynthesis: reverse of respiration, powered by light. NAD⁺/NADH and FAD/FADH₂: biological electron carriers. Catalase: 2H₂O₂ → 2H₂O + O₂ (breaks down reactive oxygen species). Antioxidants (vitamin C, E): easily oxidized, protect cells.

Redox Titrations

KMnO₄: purple, self-indicating (→ colorless Mn²⁺ in acid). Equivalent mass = M/5 (MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O). K₂Cr₂O₇: orange, diphenylamine indicator. Primary standard (weighable, stable). Iodometric: I₂ + starch → blue-black; titrate with thiosulfate. 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O.

Key Points

  • •OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons
  • •ON rules: free=0, O=-2, H=+1, sum=0 for compounds
  • •Oxidizing agent: gets reduced (ON↓); Reducing agent: gets oxidized (ON↑)
  • •Half-reaction method: split, balance atoms (Hâ‚‚O, H⁺/OH⁻), add e⁻, combine
  • •Acidic: use H⁺/Hâ‚‚O; Basic: balance acidic then neutralize H⁺ with OH⁻
  • •Disproportionation: same element oxidized and reduced simultaneously
  • •Daniel cell: Zn (anode, -0.76 V) / Cu (cathode, +0.34 V), E°cell=1.10 V
  • •Nernst: E = E° - (0.0591/n)logQ at 298 K
  • •Electrochemical series: positive E° = strong oxidant (Fâ‚‚), negative = strong reductant (Li)
  • •Galvanic: spontaneous → electricity. Electrolytic: electricity → non-spontaneous
  • •Faraday: m = E×I×t/96,485; one F = 96,485 C deposits 1 g equivalent
  • •Primary cells: non-rechargeable (dry cell). Secondary: rechargeable (Li-ion, lead-acid)
  • •Fuel cell: Hâ‚‚+½O₂→Hâ‚‚O produces electricity directly, only byproduct water
  • •Rusting: Fe oxidized by Oâ‚‚ + Hâ‚‚O; prevention: paint, galvanizing, cathodic protection
  • •KMnOâ‚„ titrations: self-indicating; equivalent mass = M/5 in acid
  • •Biological redox: respiration, photosynthesis, electron transport chain

Practice Questions

  • Calculate ON: Cr in Kâ‚‚Crâ‚‚O₇, Mn in KMnOâ‚„, S in Hâ‚‚SOâ‚„, Cl in Ca(OCl)â‚‚.
  • Balance in acidic: Crâ‚‚O₇²⁻ + Fe²⁺ → Cr³⁺ + Fe³⁺. Also in basic medium.
  • Identify OA and RA: Zn+CuSO₄→ZnSOâ‚„+Cu; 2KMnOâ‚„+10FeSOâ‚„+8Hâ‚‚SO₄→...
  • Explain Daniel cell construction, working, calculate E°cell.
  • 0.5 A passed through CuSOâ‚„ for 1 hour. Mass of Cu deposited? (F=96,485 C/mol)
  • Define standard electrode potential. Significance of electrochemical series?
  • State and explain Faraday's laws with examples.
  • Distinguish between electrolytic and galvanic cells.
  • Explain mechanism of rusting and three prevention methods.