ChemistryCode 9701
Standard electrode potentials E⦵, standard cell potentials E⦵cell and the Nernst equation
What this topic asks you to be able to do
- define what is meant by each of these: (a) standard electrode (reduction) potential (b) standard cell potential
- outline the techniques employed to determine standard electrode potentials for: (a) a metal or non-metal placed in contact with its ions in aqueous solution (b) ions of one element that exist in different oxidation states
- work out a standard cell potential by putting together two standard electrode potentials
- use standard cell potentials in order to: (a) deduce each electrode's polarity, and from that explain/deduce the direction electrons take through a simple cell's external circuit (b) predict whether a reaction is feasible
- deduce from the E⦵ values how elements, compounds and ions compare with one another as oxidising agents or reducing agents, in terms of their reactivity
- construct full redox equations from the appropriate half-equations
- predict, in qualitative terms, how E, the value of an electrode potential, responds to changes in the concentrations of aqueous ions
- apply the Nernst equation, e.g. E = E⦵ + (0.059/z) log [oxidised species]/[reduced species], to predict in quantitative terms how an electrode potential depends on aqueous ion concentrations
- understand what the equation ΔG⦵ = –nE⦵cell F means and use it