redox

• oxidation: loss of electrons
• reduction: gain of electrons
• oxidising agent: accepts electrons to oxidise another species
• reducing agent: donates electrons to reduce another species
• H+ ions:

  • in MnO4-, the oxidation state of Mn is +7, as Mn7+ ions are not stable alone
  • when reduced, the oxidation state of Mn changes to +2, leaving 4O2- ions
  • so, H+ must be added to react with the O2-, forming 4H2O
  • this is why MnO4- is acidified before use as an oxidising agent
• cell potential: voltage between two half-cells
• more negative electrode potential is easier to oxidise
• more positive electrode potential is harder to oxidise
• standard electrode potential: voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode
• standard conditions:

  • concentrations of solutions 1.0 moldm-3
  • 298 K
  • 100 KPa
• electrode potential of a cell equation$E_{cell}=$$E_{reduced}-E_{oxidised}$
• more reactive metal, wants to be oxidised more, more negative electrode potential
• more reactive non-metal, wants to be reduced more, more positive electrode potential
• more positive electrode potential, left hand equation more easily reduced, right hand more stable
• more negative electrode potential, right hand equation more easily oxidised, left hand more stable
• issues with electrode potentials:

  • will not work if conditions are not standard (temperature, pressure and concentrations)
  • reaction kinetics are unfavourable, either a too high activation energy or rate of reaction is so slow reaction may not appear to happen
• altering concentrations: Zn + Cu2+ = Cu + Zn2+

  • increasing concentration of the more negative E cell species shifts equilibrium to the left, decreasing the cell potential (reduces oxidisation)
  • increasing concentration of the more positive E cell species shifts equilibrium to the right, increasing the cell potential (increases reduction)