Electrode Potential

  • The concept of electrode potential depends on the understanding of redox reactions where electrons are transferred from one species to another, resulting in changes in oxidation states.
  • Electrode potential, also referred to as redox potential, is a measure of a system’s ability to transfer electrons from one reactant (the oxidant) to another (the reductant).
  • An electrochemical cell consists of two half-cells: each includes an electrode (metal or graphite) and an electrolyte. A salt bridge connects the two half-cells.
  • The electrode potential of half-cell is determined by the equilibrium set up between the metal atom and its ion present in the solution.
  • There are two types of electrodes in terms of potential: the anode (where oxidation occurs) and the cathode (where reduction takes place).
  • Each electrode pair has a standard electrode potential (E°), which is measured under standard conditions: temperature at 298 K, concentration of ions at 1 mol dm^-3, and pressure at 100 kPa.
  • Positive electrode potential indicates a greater tendency of the species to be reduced, while negative potential indicates a weaker tendency to be reduced.
  • The standard hydrogen electrode (SHE) serves as the universal reference for measuring electrode potentials, with an assigned potential of 0 V.
  • The electrochemical series is a list of elements organised according to their standard electrode potentials. This series helps predict the direction of spontaneous redox reactions.
  • In a cell, the overall potential, often known as the electromotive force (emf) or cell potential (Ecell), can be calculated using the equation: Ecell = Ecathode - Eanode.
  • Ecell is positive for galvanic/voltaic cells (spontaneous reactions) and negative for electrolytic cells (non-spontaneous reactions requiring external electric supply).
  • Remember that both temperature and ion concentration changes can shift the electrode potential from the standard state.
  • Standard electrode potential also aids in predicting the feasibility of a reaction and comparing the strengths of different oxidants or reductants.