Predictions from electrode potentials

• predictions can be made about the feasibility of any potential redox reactions using std electrode potentials
• the most negative system has the greatest tendency to be oxidised and lose electrons
• the most positive system has the greatest tendency to be reduced and gain electrons
• an oxidising agent takes electrons away from the species being oxidised - so oxidising agents are reduced and are on the left
• a reducing agent adds electrons to the species being reduced and so reducing agents are oxidised and are on the right
• a reaction should take place between an oxidising agent on the left and a reducing agent on the right, provided that the redox system of the oxidising agent has a more positive E value than the redox system of the reducing agent
• • the strongest reducing agent is at the top on the right
  • the strongest oxidising agent is at the bottom on the left
• to write overall equations, the half-equations must be combined
• • the reduction half-equation is the same way round as the equilibrium
  • the oxidation half equation is obtained by reversing the equilibrium
• • the redox system with the more positive (less negative) E value will react from left to right, and gain electrons
  • the redox system with the less positive (more negative) E value will react from right to left and lose electrons
• one limitation of predictions for feasibility based on delta G lies with reactions that have very large activation energy, resulting in a very slow rate - this same limitation applied to predictions based on E values

  electrode potentials may indicate the thermodynamic feasibility of a reaction but they give no indication of the rate of a reaction
• standard electrode potentials are measured using concentrations of 1 moldm-3 - many reactions take place using concentrated or dilute solutions

  • if the concentration of a solution is not 1moldm-3, then the value of the electrode potential will be different from the standard value
• example of the effect of varying concentration:

  • Zn2+ + 2e- (reversible reaction) Zn (s) - E = -0.76
  • if the concentration of Zn2+ is greater than 1moldm-3, the equilibrium will shift to the right, removing electrons from the system and making the electrode potential less negative
  • if concentration Zn2+ is less than imoldm-3, the equilibrium will shift left, increasing the electrons in the system and making the electrode potential more negative
  • any change to the electrode potential will affect the value of the overall cell potential
• other factors:

  • the actual conditions used for the reaction may be different from the standard conditions used to record E values - will affect the value of the electrode potential
  • standard electrode potentials apply to aqueous equilibria - many reactions take place that are not aqueous