Redox Reactions

• Classical Concept

  • Oxidation: Addition of oxygen / electronegative element OR removal of hydrogen / electropositive element
  • Reduction: Addition of hydrogen / electropositive element OR removal of oxygen / electronegative element

• Electron Transfer Concept

  • Oxidation: Loss of electrons (De-electronation)
  • Reduction: Gain of electrons (Electronation)
  • Oxidizing Agent (Oxidant): Accepts electrons, gets reduced
  • Reducing Agent (Reductant): Donates electrons, gets oxidized

• Oxidation Number (O.N.)

  • Definition: Charge an atom appears to have in a chemical species
  • Rules for Calculation
    • Free / Elemental state: O.N. = 0
    • Monoatomic ion: O.N. = Charge on ion
    • Hydrogen: Generally +1 (In metal hydrides: -1)
    • Oxygen: Generally -2 (Peroxides: -1, Superoxides: -1/2, $OF_2$: +2)
    • Halogens: Generally -1 (Fluorine is always -1)
    • Alkali Metals: +1 | Alkaline Earth Metals: +2
    • Neutral compound: Sum of O.N. = 0
    • Polyatomic ion: Sum of O.N. = Charge on ion
  • Stock Notation: O.N. represented by Roman numerals in parentheses (e.g., $Fe(III)_2O_3$)

• Types of Redox Reactions

  • Combination Reaction: $A + B \rightarrow C$ (e.g., $C + O_2 \rightarrow CO_2$)
  • Decomposition Reaction: Compound breaks down (e.g., $2H_2O \rightarrow 2H_2 + O_2$)
  • Displacement Reaction
    • Metal Displacement: Reactive metal displaces less reactive metal
    • Non-Metal Displacement: Displaces hydrogen or halogens
  • Disproportionation Reaction: Same element is simultaneously oxidized and reduced (e.g., $2H_2O_2 \rightarrow 2H_2O + O_2$)

• Balancing Redox Reactions

  • Oxidation Number Method
    • Identify atoms changing O.N.
    • Calculate increase/decrease in O.N. per atom
    • Multiply to equalize change in O.N.
    • Balance remaining atoms (O by adding $H_2O$, H by adding $H^+$ or $OH^-$)
  • Half-Reaction (Ion-Electron) Method
    • Separate into Oxidation and Reduction half-reactions
    • Balance atoms other than O and H
    • Balance O with $H_2O$ and H with $H^+$ (in acidic medium)
    • Add $OH^-$ equal to $H^+$ on both sides (if basic medium)
    • Balance charge by adding electrons ($e^-$)
    • Multiply half-reactions to equalize electrons and combine

• Redox Reactions and Electrode Processes

  • Redox Couple: Conjugate pair of oxidized and reduced forms ($Zn^{2+}/Zn$)
  • Galvanic / Electrochemical Cell
    • Converts chemical energy to electrical energy
    • Anode: Site of Oxidation (Negative terminal)
    • Cathode: Site of Reduction (Positive terminal)
    • Salt Bridge: Maintains electrical neutrality and completes circuit
  • Electrochemical Series
    • Arrangement of elements by standard electrode potential ($E^\circ$)
    • Higher $E^\circ$ (Positive): Stronger oxidizing agent
    • Lower $E^\circ$ (Negative): Stronger reducing agent