Redox Reactions

JKBOSE/CBSE
Redox Reactions Class 11 Chemistry Notes
Redox Reactions - Class 11 Chemistry

Redox Reactions

Overview: Oxidation and Reduction, Oxidation Number, Balancing Redox Reactions, and Electrochemical Cells.

1. Oxidation and Reduction

  • Classical: Ox (Gain O / Lose H). Red (Lose O / Gain H).
  • Electronic: Ox (Loss of e-). Red (Gain of e-). "OIL RIG".
  • Oxidation Number: Ox (Increase in ON). Red (Decrease in ON).

Redox Reaction: Both oxidation and reduction occur simultaneously.

2. Oxidation Number (ON)

Fictitious charge an atom would have if all bonds were ionic.

  • Element in free state: 0.
  • Fluorine: -1 always.
  • Oxygen: -2 (usually). -1 (peroxide). +2 (OF2).
  • Hydrogen: +1 (non-metals). -1 (metals).
Sum of ON in neutral molecule = 0

3. Balancing Redox Reactions

Ion-Electron Method

  1. Split into two half-reactions (Ox and Red).
  2. Balance atoms other than O and H.
  3. Balance O by adding H2O.
  4. Balance H by adding H+ (Acidic) or H2O/OH- (Basic).
  5. Balance charge by adding e-.
  6. Equalize e- and add half-reactions.

4. Electrochemical Cells

  • Galvanic/Voltaic Cell: Chemical Energy -> Electrical Energy. (Spontaneous).
  • Electrolytic Cell: Electrical Energy -> Chemical Energy. (Non-spontaneous).

Daniel Cell (Zn-Cu)

Anode (Ox): Zn -> Zn2+ + 2e- (Negative).
Cathode (Red): Cu2+ + 2e- -> Cu (Positive).

Ecell = Ecathode - Eanode

5. Standard Electrode Potential (E°)

Potential different between electrode and electrolyte at standard conditions (1M, 298K, 1 bar). Measured against Standard Hydrogen Electrode (SHE, E°=0).

  • Negative E°: Strong reducing agent (Li).
  • Positive E°: Strong oxidizing agent (F2).
Numericals - Redox

Numericals

Oxidation Number
Q1. Determine oxidation number of Cr in K2Cr2O7.
2(+1) + 2(x) + 7(-2) = 0
2 + 2x - 14 = 0
2x = 12
x = +6.
Balancing (Acidic)
Q2. Balance: MnO4- + Fe2+ -> Mn2+ + Fe3+ (Acidic).
Ox: Fe2+ -> Fe3+ + e- (x5)
Red: MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O (x1)
Add: MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O.
Disproportionation
Q3. Is 2H2O2 -> 2H2O + O2 a disproportionation reaction?
ON of O in H2O2 is -1.
ON of O in H2O is -2 (Reduction).
ON of O in O2 is 0 (Oxidation).
Yes, Oxygen is both oxidized and reduced.
Paradox of ON
Q4. Calculate ON of S in H2SO5 (Caro's Acid).
Structure has one peroxy linkage (-O-O-).
2 H (+1) + S (x) + 3 O (-2) + 2 O (-1) = 0.
2 + x - 6 - 2 = 0.
x = +6. (Formula method gives +8 which is wrong).
Stock Notation
Q5. Write Stock notation for AuCl3.
Au has ON +3.
Gold(III) chloride.
Equivalent Mass
Q6. Equivalent mass of KMnO4 in acidic medium? (M = 158).
MnO4- -> Mn2+. Change in ON = 7 - 2 = 5.
n-factor = 5.
Eq Mass = M / 5 = 158 / 5 = 31.6.
E Cell
Q7. Calc E0 for Zn/Zn2+//Cu2+/Cu. E0(Zn)=-0.76V, E0(Cu)=+0.34V.
Anode: Zn (Oxidation). Cathode: Cu (Reduction).
E0_cell = E0_cathode - E0_anode.
E0_cell = 0.34 - (-0.76) = 1.10 V.
Oxidizing Strength
Q8. Given E0 values: F2/F- = +2.87, Cl2/Cl- = +1.36. Which is stronger OA?
Higher reduction potential means stronger tendency to get reduced.
F2 (+2.87V) > Cl2 (+1.36V).
F2 is the stronger oxidizing agent.
Fractional ON
Q9. ON of Br in Br3O8 (Tribromooctaoxide).
Average ON = 16/3.
Structure: O3Br-BrO2-BrO3.
Terminal Br: +6 each. Middle Br: +4.
Avg = (6+4+6)/3 = 16/3.
Reaction Spontaneity
Q10. Can we store CuSO4 in Zn pot?
Reaction: Zn + CuSO4 -> ZnSO4 + Cu.
E0_cell = 1.1V (positive).
ΔG = -nFE < 0. Reaction is spontaneous.
Zn pot will dissolve. No, we cannot store.
Formulas & Facts - Redox

Equations & Formulas

ConceptFormula
Standard Cell PotentialE°cell = E°C - E°A
Gibbs EnergyΔG° = -nFE°cell
Average Ox NoTotal Charge / Total Atoms
n-factor (Oxidant)Change in ON per mole
Equivalent MassMolar Mass / n-factor

50 NEET Facts

Key points for Redox Reactions.

1. Oxidation Addition of O, removal of H, loss of electrons, increase in ON.
2. Reduction Addition of H, removal of O, gain of electrons, decrease in ON.
3. Oxidizing Agent (Oxidant) Gets reduced. Accepts electrons.
4. Reducing Agent (Reductant) Gets oxidized. Donates electrons.
5. Half Reaction Only oxidation or only reduction part.
6. Spectator Ions Ions that do not participate in reaction (ON remains same).
7. Electrochemistry Study of production of electricity from chemical energy.
8. SHE Standard Hydrogen Electrode. Pt,H2(g)|H+(aq). E=0V.
9. Fluorine ON Always -1. Most electronegative.
10. Oxygen ON Mostly -2. Peroxides -1. Superoxides -0.5. With F +2 or +1.
11. Alkali Metals ON Always +1.
12. Alkaline Earth Metals ON Always +2.
13. Max ON Group number (for p-block). Mn shows +7. Os/Ru shows +8.
14. Min ON Group number - 8. (N is 5-8 = -3).
15. Disproportionation Same element undergoes ox and red. (e.g., P4 in NaOH).
16. Comproportionation Reverse of disproportionation. Two oxidation states combine to form intermittent one.
17. Balancing (Acidic) Add H2O for O. Add H+ for H.
18. Balancing (Basic) Add H2O for O. Add H+ then neutralize with OH-.
19. KMnO4 Acidic n-factor 5 (Mn+7 to Mn+2). Eq wt = M/5.
20. KMnO4 Basic n-factor 1 (Mn+7 to Mn+6, Manganate). Eq wt = M/1.
21. KMnO4 Neutral n-factor 3 (Mn+7 to Mn+4, MnO2). Eq wt = M/3.
22. K2Cr2O7 Acidic n-factor 6 (Cr+6 to Cr+3, per Cr is 3, for 2 Cr is 6).
23. Bleaching Powder CaOCl2. One Cl is -1, other is +1. Avg 0.
24. Fe3O4 Mixed oxide (FeO.Fe2O3). Siderite.
25. C3O2 Carbon suboxide. O=C=C=C=O. Avg ON of C = 4/3. (+2 for terminal, 0 for middle).
26. Electrochemical Series Arrangement in increasing reduction potential. Li is top (lowest SRP, best reductant). F is bottom (highest SRP, best oxidant).
27. Metal Displacement Metal lower in series (higher neg potential) displaces metal higher in series. Zn displaces Cu.
28. Non-metal Displacement F2 displaces Cl-. Cl2 displaces Br-.
29. Salt Bridge Maintains electrical neutrality. Agar-agar + KCl/KNO3.
30. Hydrogen Economy Use of H2 as fuel due to high calorific value.
31. Redox Indicator Determines end point in redox titrations (Diphenylamine).
32. Self Indicator KMnO4 (Pink to Colorless).
33. Iodometry Liberated I2 titrated with hypo (Na2S2O3). Starch is indicator (Blue).
34. Corrosion Rusting of iron (Redox process).
35. Sacrificial Protection Coating with more reactive metal (Zn on Fe - Galvanization).
36. Anode of Galvanic Cell Negative pole. Oxidation occurs.
37. Cathode of Galvanic Cell Positive pole. Reduction occurs.
38. Electron Flow Anode to Cathode (external circuit).
39. Current Flow Cathode to Anode (Conventional).
40. H2O2 as Oxidant H2O2 + 2H+ + 2e- -> 2H2O.
41. H2O2 as Reductant H2O2 -> O2 + 2H+ + 2e-.
42. Nitric Acid HNO3. Oxidizing agent. Becomes NO2 or NO.
43. Conc H2SO4 Oxidizing agent. Becomes SO2.
44. Structure of CrO5 Butterfly structure. Blue peroxide. Cr is +6. Two peroxy rings.
45. Fractional Oxidation No Average of different ON of atoms of same element in a molecule.
46. Zero Oxidation No CH2Cl2 (C is 0). Glucose (C is 0).
47. Non-Redox reaction Double displacement (BaCl2 + Na2SO4). No change in ON.
48. Strongest Oxidant in Solution Mn3+ (unstable, tends to become Mn2+).
49. Reference Electrode SHE, Calomel Electrode (Hg/Hg2Cl2), Ag/AgCl.
50. Absolute Electrode Potential Cannot be measured directly. Only difference can be measured.
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