Skip to main content
Engineering LibreTexts

1.3: Electrical Double Layer

  • Page ID
    8166
  • An electrical double layer is the name given to any region between two different phases when charge is separated across the interface between them.

    In aqueous corrosion, this is the region between a corroding metal and the bulk of the aqueous environment (“free solution”). In the double layer, the water molecules of the solution align themselves with the electric field generated by applying a potential to the metal. In the Helmholtz model, there is a layer of aligned molecules (or ions), which is one particle thick and then immediately next to that, free solution. In later models (proposed by Louis Georges Gouy, David Leonard Chapman and Otto Stern) the layer is not well defined, and the orientation becomes gradually less noticeable further from the metal surface. However, for the purposes of determining the rate of corrosion, the Helmholtz model will suffice.

    To corrode, an ion in the metallic lattice must pass through the double layer and enter free solution. The double layer presents a potential barrier to the passage of ions and so has an acute effect on corrosion kinetics.

    Like all chemical processes, the kinetics involved in corrosion obey the Arrhenius relationship:

    \[k=k_{0} \exp \left(\frac{-\Delta G}{R T}\right)\]

    where k is the rate of reaction, k0 is a fundamental rate constant, ΔG is the activation energy. R and T have their usual meanings of the ideal gas constant (8.3145 J K−1 mol‑1) and temperature (in Kelvin) respectively.

    The chemical nature of corrosion suggests that it is driven by a change in Gibbs Free Energy, ΔG but the electrical nature of corrosion leads to the conclusion that a voltage drives the reaction. Since both quantities can be considered as the driving force, they must be equivalent and, indeed they are related through the expression ΔG = −z F E , where z is the stoichiometric number of electrons in the reaction, F is Faraday’s constant, 96485 C mol−1 and E is the voltage driving the reaction.

    Note the minus sign, used to correct for the conventions that a chemical reaction only proceeds if ΔG is negative but an electrical reaction only occurs if E is positive.

    Since the absolute driving force of an applied voltage depends on what reaction is occurring, potentials are usually defined as the difference between applied voltage and the equilibrium potential of the reaction. The difference between applied potential and equilibrium potential is defined as the overpotential, η.

    \[\eta=E-E_{\mathrm{e}}\]

    It is worth noting that the “equilibrium potential” is not necessarily the standard electrode potential of the reaction, as this has the added requirement that all reagents are in standard states. The equilibrium referred to is, in fact the “equilibrium electrode potential”, Ee, which is specific to every electrode individually.

    If an electrode is at its equilibrium potential, both forwards and backwards reactions occur at the same rate, so no net reaction will occur. Net reactions only occur when the potential is moved away from equilibrium.

    • Was this article helpful?