Electrode Kinetics: Exchange Current and Overpotential
Equilibrium potentials describe thermodynamic driving forces, not reaction rates. Overpotential is an electrode's departure from its equilibrium potential. Exchange current density describes equilibrium reaction activity and depends on the surface and conditions.
What a learner can do afterwards
- Defines exchange current density and explains what a large value says about an electrode reaction
- Explains overpotential as a kinetic cost and gives the reason platinum and mercury differ for hydrogen evolution
- Interprets current-potential curves or ordered measurements and distinguishes charge-transfer control from a stated mass-transfer limit
- Uses overpotential to explain a product of electrolysis that thermodynamics alone would not predict
1 · Read
At electrode equilibrium, oxidation and reduction continue with equal opposing current densities. Their net current density is zero. Exchange current density, j0, is the magnitude of either opposing current density, measured per unit electrode area. It depends on the reaction, surface, composition and temperature. Under otherwise matched conditions, larger j0 gives more current for the same small overpotential.
Use the Nernst equilibrium potential as a reference. Overpotential is eta = E - Eeq: actual electrode potential minus its equilibrium value. Reduction usually needs negative eta; oxidation needs positive eta. Activation overpotential reflects charge-transfer or surface-reaction barriers. Concentration changes at the surface can also shift the potential. Voltage lost through solution resistance is a separate cell loss.
Compare hydrogen evolution at the same current density under comparable conditions. Platinum usually needs a smaller activation overpotential magnitude than mercury because their surface reaction rates differ. Near equilibrium, a larger j0 means less overpotential magnitude is needed for the same small current density.
When charge transfer controls current, the current-potential slope depends on electrode kinetics and need not be steep. For a dissolved reactant under steady conditions, current can approach a plateau when reactant delivery limits its rate. Further potential change then produces little extra current. A plateau needs interpretation: not every electrode process becomes diffusion limited.
At a suitable anode in concentrated aqueous brine, chloride oxidation can give mainly chlorine. Oxygen evolution competes, but slower oxygen-evolution kinetics and high chloride concentration help favour chlorine. At the cathode, water reduction to hydrogen is thermodynamically easier than sodium-ion reduction to metal. Sodium's absence is not an exception explained only by slow kinetics.
Compare competing reactions at the actual potential, surface and composition. Each path carries a partial current. In the usual convention, oxidation currents are positive and reduction currents negative. Compare magnitudes: for equal electron counts per product molecule, a larger magnitude means faster product formation. If electron counts differ, include them in the comparison. Equilibrium potentials alone do not give rates.
Exchange current density describes equilibrium activity. Overpotential drives net current, reactant supply can limit it, and product rates depend on operating conditions.
Lesson, questions and flashcards written for this dataset.
Readings and flashcards written for this dataset. Videos by Khan Academy and EChem Channel.
More to read
Zero current does not mean zero reaction
Imagine measuring an electrode at its equilibrium potential. The instrument reports zero net current, but oxidation and reduction can still occur. Their equal current densities point in opposite directions and cancel. Exchange current density, j0, names the magnitude of either opposing contribution.
The word density matters. Comparing total currents from electrodes of different areas can confuse a size difference with a kinetic difference. State which electrode area is used when reporting j0, especially for rough or porous surfaces.
The same redox reaction can have different j0 values on different surfaces or at different temperatures and compositions. Near equilibrium, otherwise matched systems with larger j0 respond with more current to a small overpotential. This does not mean that their net equilibrium current is nonzero.
Account for voltage losses separately
Begin with the equilibrium potential for the actual electrode conditions. The Nernst equation provides this reference. Subtract it from the actual electrode potential to obtain overpotential: eta = E - Eeq.
Some overpotential is associated with the rates of charge transfer and surface chemistry. This activation contribution depends on the electrode material. For hydrogen evolution at a matched current density, platinum usually requires less activation overpotential magnitude than mercury under comparable conditions.
Reactant supply can impose another limit. A dissolved reactant may be consumed faster than it reaches the surface. Under steady conditions, the current can then approach a limiting plateau.
Resistance in the electrolyte also causes a voltage drop. This solution-resistance loss is separate from electrode overpotential. During electrolysis at finite current, activation and resistance losses increase the required applied voltage above the reversible requirement for the same reaction and conditions. During galvanic discharge, these losses instead lower the delivered terminal voltage below the reversible cell voltage for the same conditions.
Predict products at each electrode
Aqueous brine presents two different competitions. At the cathode, water reduction to hydrogen is thermodynamically easier than sodium-ion reduction to sodium metal under ordinary brine conditions. Hydrogen production there does not require a claim that sodium reduction loses only because of slow kinetics.
At the anode, chlorine formation competes with oxygen evolution. High chloride concentration and an anode with suitable reaction kinetics can favour chlorine. Changing the surface or solution can change the product mixture.
To compare product rates, examine each reaction's partial current at the operating potential. In the usual sign convention, oxidation currents are positive and reduction currents negative. Compare magnitudes rather than algebraic values: for equal electron requirements per product molecule, the larger magnitude represents faster product formation. Different electron requirements must be included in the comparison. A potential table gives equilibrium information, not the partial currents of a working electrode.
Explore more
2 · Watch
به انگلیسیTake it off screen
Where it sits
Optional connections
8 questions wait behind this lesson, each with its answer explained. Every answer feeds the sky: stars light as they are learned, and dim when it is time to come back.