Oxidation-Reduction Reactions

Basics

• Oxidation is loss, reduction is gain (OIL-RIG)
• Oxidizing agent causes oxidation; it gets reduced
• Most oxidizing agents contain O or a strongly electronegative element
• Reducing agents often have metals or hydrides
• More bonds to O = oxidation
• Example: methanol → methanoic acid is oxidation
• Some species can be both (like NAD⁺/H)

Assigning oxidation numbers

• Free element = 0
• Monatomic ion = its charge
  • Ca²⁺ = +2
• Group IA/IIA = +1/+2
• Group VIIA = −1 except when with more electronegative elements
  • HCl → Cl = −1
  • HOCl → Cl = +1
• Hydrogen usually +1
  • In metal hydrides (NaH), H = −1
• Oxygen usually −2
  • In peroxides, O = −1
  • In OF₂, O = +2
• Oxidation number ≠ formal charge
  • Oxidation assumes unequal division (more EN gets e⁻)
  • Formal charge assumes equal division

Balancing redox reactions

• Use half-reaction (ion–electron) method
• Separate into two halves and balance individually
• Balance atoms; add H₂O if needed
• Add e⁻ where necessary
• Multiply to match number of e⁻ transferred
• LAH = strong reducing agent
  • Has metal + hydrides
  • Donates e⁻ via hydride

• Gains and losses of e⁻ drive formation of many ionic compounds
• Splitting reaction gives complete ionic equation
• Ions unchanged on both sides = spectator ions → removed to get net ionic equation
• In net ionic equations, all aqueous compounds split into ions

Combination reactions

• Two or more species come together → product

Decomposition reactions

• One product breaks down into two or more species
• Sometimes no spectator ions

Combustion reactions

• Fuel + O₂ → CO₂ + H₂O
• Can have complex half-reactions
• Example: C in methane goes from −4 to +4 (oxidized)

Double displacement (metathesis) reactions

• Counterions switch
• Not usually redox → oxidation states stay the same
• If all ions appear on both sides → no net ionic equation

Disproportionation (dismutation) reactions

• Same element undergoes both oxidation and reduction
• Catalase catalyzes peroxide breakdown in peroxisomes
  • H₂O goes from −1 → −2 and 0
• Superoxide dismutase disproportionates oxygen free radicals
  • Free radical has −1/2 oxidation state → becomes −1 (in peroxisome) and 0 (O₂)
• Often enzyme-mediated
• Frequently involve metals like Cu or Zn in active sites as reducing agents

Oxidation–reduction titrations

• Follow transfer of charge to equivalence point
• Indicators change color at specific emf (voltage)
• Starch indicator identifies iodine complexes (dark color in presence of iodine)
• Iodimetry can standardize thiosulfate
• Potentiometric titration: no indicator; voltage measured with voltmeter
• Voltage changes as redox titration progresses