Oxidation, Reduction and Redox Reactions

• Oxidation

  • the loss of hydrogens and electrons
  • the gain of oxygen
  • oxidation state will become more negative
• Reductions

  • loss of oxygen
  • gain of electrons and hydrogen
  • oxidation state will become more positive
• Oxidation States

  • fluorine -> -1
  • Group 7 (not in presence of F or O) -> -1
  • in presence of F or O -> +1
  • hydrogen -> +1
  • metal hydrides -> -1
  • group 1 -> +1
  • group 2 -> +2
  • oxygen (not peroxides) -> -2
• Peroxides

  • due to O-O bond, oxygens have shared electron density so do not reach their full “2-” oxidation state
  • e.g. in a hydrogen peroxide, O is usually -1
• Oxidising agents

  • oxidises another species by gaining the electrons the other species loses, reducing iself
  • e.g. Cr2O7+ -> Cr2O7
• Reducing Agents

  reduces another species by donating the electrons the other species gains, oxidising itself
• Disproportionation

  when atoms of the same element are oxidised and reduced simultaneously
• Constructing half equations

  • in order:
  • balance atoms
  • balance O atoms by adding H2O
  • balance H atoms by adding H+
  • balance charges by adding e-
• Full redox equations

  • two half equations can be combined to make an overall redox equations
  • the number of electrons in each equation must match before combining - scale up half equations as appropriate
  • cancel out the electrons after combining