Redox Reactions & Oxidation Number

🔍 Why Oxidation Number?

Sometimes, electron transfer isn’t obvious (like in covalent compounds). Oxidation numbers help “book-keep” electron shifts by assuming complete electron transfer from less electronegative to more electronegative atoms. Example:

  • \( 2\ce{H2}(g) + \ce{O2}(g) \rightarrow 2\ce{H2O}(l) \)
    Oxidation numbers: H: 0 → +1, O: 0 → -2

📜 Rules for Assigning Oxidation Numbers

  1. Free elements: Oxidation number = 0
    (e.g., \(\ce{Na}\), \(\ce{O2}\), \(\ce{P4}\)).
  2. Monatomic ions: Oxidation number = charge
    (e.g., \(\ce{Na+}\): +1, \(\ce{Cl-}\): -1).
  3. Oxygen: Usually -2
    Exceptions: Peroxides (\(\ce{H2O2}\): -1), superoxides (\(\ce{KO2}\): -½), bonded to F (\(\ce{OF2}\): +2).
  4. Hydrogen: Usually +1
    Exception: Metal hydrides (\(\ce{NaH}\): -1).
  5. Fluorine: Always -1.
  6. Sum in compound: Zero
    Sum in polyatomic ion: Equals ion charge
    (e.g., \(\ce{(CO3)^{2-}}\): C + 3×O = -2).

⚡ Types of Redox Reactions

  1. Combination: \( \ce{A + B -> C} \) (at least 1 element)
    Example: \(\ce{CH4}(g) + 2\ce{O2}(g) \rightarrow \ce{CO2}(g) + 2\ce{H2O}(l)\)
  2. Decomposition: Compound → Elements + other compounds
    Example: \(\ce{2H2O}(l) \xrightarrow{\Delta} \ce{2H2}(g) + \ce{O2}(g)\)
  3. Displacement:
    • Metal displacement: \(\ce{Zn}(s) + \ce{CuSO4}(aq) \rightarrow \ce{Cu}(s) + \ce{ZnSO4}(aq)\)
    • Non-metal displacement: \(\ce{Cl2}(g) + 2\ce{KI}(aq) \rightarrow 2\ce{KCl}(aq) + \ce{I2}(s)\)
  4. Disproportionation: Same element oxidized & reduced
    Example: \(\ce{2H2O2}(aq) \rightarrow 2\ce{H2O}(l) + \ce{O2}(g)\)
    Oxygen: -1 → 0 (oxidized) and -1 → -2 (reduced).

⚖️ Balancing Redox Reactions

Method 1: Oxidation Number

  1. Write skeletal equation.
  2. Assign oxidation numbers.
  3. Balance atoms with changed oxidation states.
  4. Add \(\ce{H+}\) (acidic) or \(\ce{OH-}\) (basic) to balance charge.
  5. Add \(\ce{H2O}\) to balance H/O atoms.

Example (acidic):
\[ \ce{Cr2O7^{2-} + \ce{SO3^{2-} -> Cr^{3+} + SO4^{2-}} \]
Balanced: \[ \ce{Cr2O7^{2-} + 3SO3^{2-} + 8H+ -> 2Cr^{3+} + 3SO4^{2-} + 4H2O} \]

Method 2: Half-Reaction

  1. Split into oxidation & reduction half-reactions.
  2. Balance atoms (except H/O).
  3. Add \(\ce{H2O}\) and \(\ce{H+}\)/\(\ce{OH-}\) to balance O/H.
  4. Add electrons to balance charge.
  5. Multiply to equalize electrons, then combine.

Example (basic):
\[ \ce{MnO4- + I- -> MnO2 + I2} \]
Balanced: \[ \ce{2MnO4- + H2O + 6I- -> 2MnO2 + 2OH- + 3I2} \]

🧪 Redox Titrations

  • \(\ce{MnO4-}\) (purple) acts as self-indicator (turns colorless → pink at endpoint).
  • \(\ce{Cr2O7^{2-}}\) uses diphenylamine (blue at endpoint).
  • Iodometry: \(\ce{I2}\) with starch (blue) + thiosulphate (\(\ce{S2O3^{2-}}\)) for quantification.

⚠️ Limitations of Oxidation Number

Fractional oxidation states (e.g., \(\ce{Fe3O4}\): avg. +8/3 for Fe) are averages! Actual oxidation states differ:
\(\ce{Fe3O4} = \ce{Fe^{2+} + 2\ce{Fe^{3+}} \). 🧐

🚀 NEET Focus Areas

  1. Oxidation Number Rules: Assign correctly (e.g., O in peroxides, H in hydrides).
  2. Disproportionation: Identify reactions where same element is oxidized & reduced (e.g., \(\ce{Cl2}\) in alkali).
  3. Balancing Equations: Master both methods (oxidation number & half-reaction) for acidic/basic media.
  4. Metal Reactivity: Use activity series (\(\ce{Zn > Cu > Ag}\)) to predict displacement.
  5. Redox Titrations: Indicators for \(\ce{MnO4-}\), \(\ce{Cr2O7^{2-}}\), iodometry.